EQUILIBRIUM
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Chemistry and Society LEARNING OUTCOMES
|Section 1 Getting the most from costly reactants |
|Subsection (a) Factors influencing the design of an industrial process. |
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|1. |I can state what industrial processes are designed to maximise and minimise. |
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| |Industrial processes are designed to maximise profit and minimise the impact on the environment. |
|2. |I can give the 7 factors that influence the design of an industrial process. |
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| |Factors influencing process design include: availability, sustainability and cost of feedstock(s); opportunities for |
| |recycling; energy requirements; marketability of by-products; product yield. |
|3. |I know what environmental issues need to be considered when designing a chemical process. |
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| |Environmental considerations include: minimising waste; avoiding the use or production of toxic substances; designing |
| |products which will biodegrade if appropriate. |
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|Subsection (b) calculation of the mass. |
|4. |I am able to use balanced equations to work out mole ratios of reactants and products. |
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|5. |I can use balanced equations and formula mass to work out the mass of product from reactant and vice versa. |
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|6. |I can calculate the volume of a gas from the number of moles and vice versa. Give the units for molar volume. |
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| |litres mol -1 |
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|7. |I know that molar volume is the same for ALL gases at the same temperature and pressure. |
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| |The volume taken up by 1 mole of a gas is called the molar volume. |
| |The value of the molar gas volume is approx 22.4 litres mol -1 depending on the temperature and pressure of the environment.|
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|8. |I can calculate the volumes of reactant and product gases from the number of moles of each reactant and product. |
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|Subsection (c) Calculations in reactions that involve solutions. |
|9. |I can give the units for concentration |
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| |mol l-1 |
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|10. |I can work out quantities of reactants and/or products using one or more of the following: |
| |Balanced equations. |
| |Concentrations and volumes of solutions. |
| |Masses of solutes. |
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|Subsection (d) Reversible reactions. |
|11. |I understand what is meant by “dynamic equilibrium”. |
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| |Many reactions are reversible. A reversible reaction can reach equilibrium in a closed system. |
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|12. |I know what happens to reaction rates at equilibrium. |
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| |A reaction reaches equilibrium when the rate of the forward reaction equals the rate of the reverse reaction. |
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|13. |I know what happens to concentrations of products and reactants at equilibrium. |
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| |At equilibrium, the concentration of the products and the reactants will remain constant. |
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|14. |I can state Le Chatelier’s Principle |
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| |An equilibrium will move to undo any change imposed upon it by temporarily favouring either the forward or backward reaction|
| |until a new equilibrium position is reached. |
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|15. |I can explain using Le Chatelier’s Principle, the likely effect on the equilibrium position of changing:- |
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| |Pressure |
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| |An increase in pressure favours the side with less gas molecules |
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| |Concentration of reactants or product |
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| |Addition of a reagent or removal of product equilibrium shifts to the right. |
| |Addition of product or removal of reactant equilibrium shifts to the left. |
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| |Temperature |
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| |Increase in temperature favours the endothermic reaction |
| |Decrease in temperature favours the exothermic reaction |
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| |Catalyst |
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| |No effect on equilibrium position; equilibrium more rapidly attained. |
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|16. |I can relate the effects of temperature, pressure, concentration of reactants and products to industrial processes (Haber |
| |Process). |
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| |Haber uses high pressure (balanced by cost of equipment), moderately high temperature (lower temperature improves yield but |
| |reaction too slow). Unreacted products are recycled, catalyst (iron) used to speed up the reaction |
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|17. |I understand what is meant by “percentage yield” and can give the formula. |
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| |The percentage yield is a measure of how much of a product is obtained compared to the amount expected if there was complete|
| |conversion. |
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| |The percentage yield = actual yield x 100 |
| |theoretical yield |
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|18. |I can calculate percentage yield from balanced equations and masses of reactants and products. |
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|19. |I can use given costs, and percentage yields to work out the cost of the feedstock(s) to produce a given mass of product. |
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|20. |I understand what is meant by “atom economy”. |
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| |The atom economy measures the proportion of the total mass of all starting materials successfully converted into the desired|
| |product. |
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|21. |I can calculate the atom economy of a reaction using the correct formula. |
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| |Atom economy = mass of desired products x 100 |
| |Total mass of reactants |
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|22. |I understand why some reactions that have a high percentage yield may have a low atom economy. |
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| |High percentage yield and a low atom economy means a lot of waste products a re being produced. |
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|23. |I can relate percentage yield and atom economy to different routes taken in manufacturing products. |
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| |Industrial process want efficiency i.e. high percentage yield an high atom economy. |
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|Subsection (g) Excess. |
|24. |I can explain and identify the excess reactant(s). |
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| |The reaction will be over once one of the reactants has been used up – the other one is said to be ‘in excess’. |
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|25. |I can relate cost to excess reactant(s) used in industrial processes. |
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| |In order to minimise costs, industry aims to have the cheaper reactant in excess where possible. |
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|26. |I can use a balanced equation to work out the reactant in excess and therefore the limiting reactant, for a chemical |
| |reaction. |
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|27. |I can relate excess reactant(s), percentage yields and atom economy to the idea of an economic/environmental balance in |
| |industrial processes. |
|SECTION 2 Chemical Energy |
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|Subsections (a) and (b) Enthalpy. |
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|28. |I can understand what is meant by “enthalpy”. |
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| |Enthalpy (H) is a measure of the energy stored in a chemical. |
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|29. |I can give a definition of enthalpy of combustion. |
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| |The enthalpy of combustion of a substance is the amount of energy given out when one mole of a substance burns in excess |
| |oxygen. |
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|30. |I can use specific heat capacity, mass and temperature to calculate the enthalpy change for a reaction. |
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|31. |I can work out the enthalpy of combustion from practical experiments. |
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|Subsection (c) Hess’s Law |
|32. |State Hess’s Law. |
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| |The overall enthalpy change for a reaction is the same whichever route is taken. |
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|33. |Applying Hess’s Law:- |
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| |I can use enthalpies of combustion, given in the data book, to calculate enthalpies of formation. |
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| |I can calculate the overall enthalpy change for a reaction using given enthalpy data. |
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|Subsection(d) Bond Enthalpies. |
|34. |I understand what is meant by:- |
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| |molar bond enthalpy |
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| |the molar bond enthalpy is the energy required to break one mole of bonds. |
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| |mean molar bond enthalpy |
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| |Where bonds are present in a molecule with more than 2 atoms, the bond enthalpy will be affected by the environment the bond|
| |is in, so a C-H bond in methane may have a slightly different bond enthalpy from one in propene. Therefore the average or |
| |mean bond enthalpy is worked out. |
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|35. |I can use bond enthalpies to estimate the enthalpy change taking place for a gas phase reaction. |
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|Section 3 Oxidising and Reducing Agents |
|Subsection (a) Elements as Oxidising and Reducing Agents |
|36. |I can define oxidising agent and reducing agent in terms of electrons. |
| |Oxidising agents accept electrons. Reducing agents give away electrons. |
|37. |I can identify oxidising agents and reducing agents in redox reactions. |
| |Oxidising agents are themselves reduced so tend to become more negative. |
| |Reducing agents are themselves oxidised so tend to become more positive. |
|38. |I can use electronegativity to predict which elements tend to lose electrons and which elements tend to gain electrons when |
| |they form ions. |
| |Metals have low electronegativities so tend to lose electrons. |
| |Non-metals have high electronegativities so tend to gain electrons. |
|39. |I can use electronegativity to predict which elements can act as reducing agents and which can act as oxidising agents. |
| |Metals tend to lose electrons when they form ions so metals are reducing agents. |
| |Non -metals tend to gain electrons when they form ions so non-metals are oxidising agents. |
|40. |I can state which group in the Periodic Table contains the strongest reducing agents and which group contains the strongest |
| |oxidising agents. |
| |The alkali metals (group 1) are the strongest reducing agents. |
| |The halogens (group 7) are the strongest oxidising agents. |
|41. |I can use the Electrochemical Series to determine the effectiveness of oxidising and reducing agents. |
| |The strongest oxidising agents are at the bottom left hand corner of the Electrochemical Series. |
| |The strongest reducing agents are at the top right hand corner of the Electrochemical Series. |
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|Subsection (b) Molecules and group ions can act as Oxidising and Reducing |
|Agents |
|42. |I can use the Electrochemical Series to identify oxidising agents and reducing agents. |
| |Oxidising agents are at the left hand side of the electrochemical series. |
| |Reducing agents are on the right. |
|43. |I can predict how acidified dichromate and acidified permanganate react. |
| |H+/Cr2O72- and H+/MnO4- are oxidising agents. |
|44. |I can predict how hydrogen peroxide and carbon monoxide react. |
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|Subsection (c) Everyday uses for Strong Oxidising Agents |
|45. |I can state examples of uses for oxidising agents |
| |Used as bleach for clothes and hair. |
| |Can kill fungi and bacteria and can inactivate viruses. |
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|Subsection (d) Ion-electron and Redox Equations |
|46. |I can complete an ion electron half equation when the reactant and product is known. |
| |Balance the unusual atoms first. |
| |Balance the oxygens by adding water. |
| |Balance the hydrogens by adding H+ ions. |
| |Balance the charge by adding electrons to the more positive side. |
|47. |I can combine ion-electron equations to produce a redox equation. |
| |Make sure there is an oxidation and a reduction. |
| |Multiply one or both of the ion-electron equations to make sure they both have the same number of electrons. |
| |When the equations are combined the electrons should cancel. |
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|SECTION 4 Chemical Analysis as part of quality control. |
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|Subsection (a) checking composition and purity of reactants and products. |
|48. |I can explain the basic principle of how chromatography works, defining the mobile and stationary phases |
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| |Chromatography separates compounds according to their relative affinity for the ‘mobile phase’ and the ‘stationary phase’. |
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| |The mobile phase is a liquid or a gas. The size of molecules and their polarity may affect how soluble they are in the |
| |mobile phase |
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| |The stationary phase may be paper, silica gel, or an inert packing material. The size and polarity of the compounds may |
| |affect their affinity for the stationary phase. |
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|49. |I can read and interpret retention/time graphs from results of chromatography experiments. |
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|50. |I can interpret chromatograms using Rf values |
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| |Rf = distance travelled by sample |
| |distance travelled by solvent |
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|Subsection (b) Volumetric titrations. |
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|51. |I can state the principle of volumetric analysis using titration. |
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| |Volumetric analysis involves using a solution of known concentration to determine the amount of another substance present. |
| |The volume of the reactant needed to complete the reaction is determined by titration. |
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|52. |I can give a definition of an indicator and name some examples. |
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| |Indicator: A substance that changes color in response to a chemical change. |
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| |Phenolphalein (pink/colourless) |
| |Permaganate self indicating (purple/colurless) |
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|53. |I know what the “end point” of a reaction is. |
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| |The term "end point" is where the indicator changes colour and means that the solutions have been mixed in exactly the right|
| |proportions according to the equation. |
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|54. |Use the results of titrations and balanced redox equations to calculate the concentration of a reactant, given the |
| |concentration of the other. |
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|55. |Explain what a standard solution is and how to make up a standard solution. |
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| |A solution of accurately known concentration is known as a standard solution. |
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| |Dissolve the calculated mass of solid in the minimum volume of water in a beaker, transfer to standard flask rinsing the |
| |beaker, dilute up to the graduated mark on the standard flask |
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|56. |Explain the relative accuracy of a range of volumetric measuring equipment. |
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| |Burette and pipette allow the volume to be measured accurately |
| |Errors may occur in judging the endpoint/colour change |
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CHEMISTRY IN SOCIETY
The products we use every day, from shampoo to concrete, are produced by industry. Industrial processes are designed to maximise profit and to minimise the impact on the environment.
Reactants are costly, and it is essential that the processes used provide the best possible return on investment. This unit looks at the underlying principles used to develop the design of these industrial processes.
CHEMICAL INDUSTRY CASE STUDY
Using the resources, research the following learning outcomes
➢ Industrial processes are designed to maximise profit and minimise the impact on the environment.
➢ Factors influencing process design include: availability, sustainability and cost of feedstock(s); opportunities for recycling; energy requirements; marketability of by-products; product yield.
➢ Environmental considerations include: minimising waste; avoiding the use or production of toxic substances; designing products which will biodegrade if appropriate.
Designing an Industrial Process
1. Aluminium is extracted from its purified oxide by molten electrolysis. Suggest two advantages and two disadvantages of siting aluminium smelters in the Scottish Highlands.
2. Write down 5 factors which will affect the design of a chemical process.
3. A reaction in a chemical plant is exothermic.
(a) Explain what is meant by an exothermic reaction.
(b) How does the exothermic reaction help the chemical plant make a profit?
4. By-products are usually created in industrial chemical processes.
(a) Explain what is meant by a by-product.
(b) How can by-products be
(i) profit making?
(ii) profit losing?
5. As well as transportation, what other reasons could there be for siting chemical works close to rivers?
6. What are the four stages in the manufacture of a new product? Write a sentence or two to explain each stage.
7. Imagine you were researching a method for converting benzene to a long chain alkylbenzene for detergent manufacture. What mass of reagents would you use?
A milligrams
B grams
C kilograms
D metric tonnes (1000 kg)
8. In the bulk manufacturing process for converting benzene to alkylbenzene for detergents, what mass of reagents would be used?
A milligrams
B grams
C kilograms
D metric tonnes (1000 kg)
9. Give three ways that the operators of chemical plants can minimize the effect of the processes on the environment.
10. The Haber process is used to make ammonia. For this process list
(a) the feedstocks
(b) the raw materials
Using the mole ratio to predict the amount of solid formed in a reaction
Carry out the practical to confirm the mass of magnesium oxide formed
When magnesium ribbon reacts with oxygen, magnesium oxide is formed. If we know the mass of magnesium we start with, we can predict the mass of product expected.
| |Mass (g) |
|Mass of crucible | |
|Mass of crucible + Mg ribbon | |
|Calculated mass of Mg ribbon | |
|Mass of crucible + MgO | |
|Calculated mass of MgO | |
2Mg + O2 → 2MgO
➢ Using the balanced equation calculate the mass of MgO you would expect to be formed?
➢ How does that compare to your result?
➢ Account for any differences.
Reduction of copper(II)oxide by hydrogen
If 3.9755g of copper(II)oxide was reacted with hydrogen, what mass of copper should be formed?
CuO + H2 → Cu + H2 O
Calculations from Equations
number of moles = mass concentration = number of moles
GFM volume
1. What is the mass of
a) 1 mole of H2O
b) 5 moles of CO2
c) 20 moles of NH3
d) 0.1 moles of C2H5OH
e) 0.05 moles of C2H4?
2. How many moles are in
a) 1.8g of H2O
b) 8.8g of CO2
c) 1.755kg of NaCl
d) 40kg of MgO
e) 0.12g of NaOH
3. What is the concentration of a solution containing:
a) 5 moles of NaCl dissolved in 5 l of water
b) 0.3 moles of NaCl dissolved in 0.5 l of water
c) 2 moles of NaCl dissolved in 250 ml of water
d) 175.5g of NaCl dissolved in 500 ml of water
e) 11.7g of NaCl dissolved in 10 l of water
4. Calculate the number of moles of solute required to prepare 250cm3 of a 0.100 mol l-1 solution of oxalic acid.
5. What mass of copper(II) sulphate crystals (CuSO4.5H2O) would be contained in 100cm3 of a 2.50 mol l-1 solution ?
6. How many grams of magnesium oxide would be produced by reacting completely 4.0 g of magnesium with oxygen?
2Mg + O2 ( 2MgO
7. Oxygen can be converted into ozone (O3) by passing an elecrical discharge through it.
Calculate the number of ozone molecules that would be formed if 16g of oxygen were completely converted into ozone.
3O2 ( 2O3
8. Ammonia reduces copper(II) oxide to copper. The other products of the reaction are water and nitrogen.
2NH3 + 3CuO ( 3Cu + H2O + N2
Calculate the mass of copper produced and the mass of ammonia consumed when 56.4g of copper(II) oxide are reduced in this way.
9. What mass of aluminium will be needed to react with 10 g of CuO, and what mass of Al2O3 will be produced?
3CuO + 2Al ( Al2O3 + 3Cu
10. In a reaction magnesium carbonate powder is used to neutralise 250 cm3 of 2 mol l-1 dilute hydrochloric acid. Calculate the mass of magnesium carbonate required to neutralise the dilute hydrochloric acid
MgCO3 + 2HCl ( MgCl2 + H2O + CO2
11. In a reaction sodium carbonate powder is used to neutralise 200 cm3 of 2 mol l-1 dilute sulphuric acid. Calculate the mass of sodium carbonate required to neutralise the dilute sulphuric acid.
Na2CO3 + H2SO4 ( Na2SO4 + H2O +CO2
12. 20 cm3 of a solution of NaOH is exactly neutralised by 25 cm3 of a solution of HCl of concentration 0.5 mol l–1.
HCl + NaOH → NaCl + H2O
Calculate the concentration of the NaOH solution in mol l –1.
13. 100 cm3 of a solution of KOH is exactly neutralised by 150 cm3 of a solution of H2SO4 of concentration 0.25 mol l–1.
H2SO4 + 2KOH → K2SO4 + 2H2O
Calculate the concentration of the KOH solution in mol l –1.
14. 50 cm3 of a solution of HCl is exactly neutralised by 20 cm3 of a solution of Ca(OH)2 of concentration 2.0 mol l–1.
2HCl + Ca(OH)2 → CaCl2 + 2H2O
Calculate the concentration of the HCl solution in mol l –1.
15. 2.5 l of a solution of NaOH is exactly neutralised by 1.5 l of a solution of HCl of concentration 1.0 mol l–1.
HCl + NaOH → NaCl + H2O
Calculate the concentration of the NaOH solution in mol l –1.
MOLAR GAS VOLUME
One mole of all gases at the same temperature and pressure will have the same volume.
The volume taken up by 1 mole of a gas is called the molar volume.
Since there is 1 mole of gas in each volume there is also the same number of particles. The value of the molar gas volume is approx 22.4 litres mol -1 depending on the temperature and pressure of the environment.
Comparing gas volumes
Since the larger ‘balloon’ has twice the volume it must have twice the number of moles of gas within it. It must also have twice as many particles.
1. Comparing volumes
➢ Which has the greatest volume, 20g of NO2 or 20g of Cl2?
➢ Which has the greatest volume, 22g of of CO2 or 22g of SO2 ?
➢ Which has more particles, 20g of neon or 46g of sulphur dioxide?
2. Calculating the molar volume
➢ If 4 g of methane occupies 6 litres , what is the molar volume?
➢ If 1.1 g of CO2 occupies 0.625 litres, what is the molar volume?
3. Finding out the molar volume by experiment.
Density = mass/volume
Or mass = density x volume
Using this equation complete the following table.
|Gas |H2 |CH4 |N2 |O2 |Ar |CO2 |
|Gram formula mass/g | | | | | | |
|Density at 0oC and 1 atm pr/g l-1 | | | | | | |
| |0.09 |0.71 |1.25 |1.43 |1.78 |1.98 |
|Molar volume at 0oC and 1 atm pr/l| | | | | | |
Determination of molar volume of butane by experiment
Predicting the volume of gas that will be formed in a reaction.
Example 1. What volume of hydrogen would be produced if 20.0 cm3 H2SO4, concentration
0.5 mol l-1 reacts completely with excess zinc? (molar gas volume is 22.4 litres mol -1)
Zn + H2SO4, → ZnSO4, + H2
Example 2 What volume of gas will be produced when 3.705 g copper carbonate decomposes? (molar gas volume is 22.4 litres mol -1)
CuCO3 → CuO + CO2
CONFIRMING THE PREDICTION EXPERIMENTALLY
When copper carbonate reacts with sulphuric acid, carbon dioxide gas is produced. If we know how many moles of calcium carbonate we start with, we can work out the expected volume of gas.
CuCO3 + H2 SO4 → CuSO4 + CO2 + H20
1 mol 1 mol
1. Weigh out about 0.5g of solid. Record the mass accurately below
2. Assuming molar gas volume is 22. 4 litres mol -1 calculate the expected volume of gas.
3. Set up a boiling tube fitted with a bung and delivery tube, to collect gas over water in an inverted measuring cylinder. Add the weighed solid to the boiling tube. Add 10 cm3 of 2 mol l-1 H2SO4 , quickly stopper the boiling tube and collect the gas.
6. Record the volume of gas produced.
7. How accurate was your prediction? Account for any difference between the expected value and that obtained.
The molar volume of any gas is 24 litres at s.t.p.
16. What volume would the following amounts of gas occupy?
a) 2 moles of helium
b) 0.1 moles of oxygen
c) 5.5 moles of nitrogen
d) 2.4 g of ozone
e) 0.88 g of carbon dioxide
f) 32 g of sulphur trioxide
17. What volume (in l) of carbon dioxide would be produced by completely reacting 60 g of carbon with oxygen?
C + O2 ( CO2
18. What volume (in l) of hydrogen would be produced by completely reacting 60 cm3 of hydrochloric acid of concentration 1.2 mol l–1 with zinc?
Zn + 2HCl ( ZnCl2 + H2
19. What volume (in l) of carbon dioxide would be produced by completely reacting 10g of calcium carbonate with hydrochloric acid?
CaCO3 + 2HCl ( CaCl2 + H2O + CO2
20. What volume (in l) of hydrogen would be produced by completely reacting 60 cm3 of hydrochloric acid of concentration 1.2 mol l–1 with zinc?
Zn + 2HCl ( ZnCl2 + H2
21. In the reaction of lithium with water, what mass of lithium (in grams) would be required to produce 600 cm3 of hydrogen?
2Li + 2H2O ( 2LiOH + H2
22. Calculate the volume of oxygen that would be required to react completely with 1.0 l of methane.
CH4 + 2O2 ( CO2 +2H2O
23. Calculate the volume of oxygen that would be required to react completely with 5.0 l of ethane.
C2H4 + 3O2 ( 2CO2 +2H2O
THE IDEA OF EXCESS
When a reaction takes place between 2 reactants, it is very unlikely that both of the substances are in exactly the right proportions and that they will both run out at the same time. Usually, one runs out before the other and this reactant limits how much product can be formed.
The reaction will be over once one of the reactants has been used up – the other one is said to be ‘in excess’. In order to minimise costs, industry aims to have the cheaper reactant in excess where possible.
Calculating excess
Example. 10g zinc was added to 25cm3 of HCl concentration 0.5 mol l-1. Show by calculation which one is in excess.
Zn + 2HCl → Zn Cl2 + H2
➢ The zinc costs more than the acid so this is not a very economical procedure. What could be done to make it more cost effective?
Example 2. 0.05g magnesium was added to 25cm3 of HCl concentration 1.0 mol l-1. Show by calculation which one is in excess.
. Mg + 2HCl → Mg Cl2 + H2
QUESTIONS
24. Iron(II) sulphide reacts with hydrochloric acid as follows:
FeS(s) + 2HCl(aq) ( FeCl2(aq) + H2S(g)
If 4.4g of iron(II) sulphide was added to 160cm3 of 0.5 mol l-1 hydrochloric acid, show by calculation which substance is in excess.
25. A student added 0.20g of silver nitrate, AgNO3, to 25 cm3 of water. This solution was then added to 20cm3 of 0.0010 mol l-1 hydrochloric acid. The equation for the reaction is:
AgNO3(aq) + HCl(aq) ( AgCl(s) + HNO3(aq)
Show by calculation which reactant is in excess.
26. Calcite is a very pure form of calcium carbonate which reacts with nitric acid as follows:
CaCO3(s) + 2HNO3(aq) ( Ca(NO3)2(aq) + H2O(l) + CO2(g)
A 2.14g piece of calcite was added to 50.0cm3 of 0.200 mol l-1 nitric acid in a beaker.
Calculate the mass of calcite, in grams, left unreacted.
27. Copper(II) oxide reacts with sulphuric acid as follows:
CuO(s) + H2SO4(aq) ( CuSO4(aq) + H2O(l)
1.6 g of copper(II) oxide is added to a beaker containing 50cm3 of 0.25 mol l-1 sulphuric acid.
Calculate the mass of copper(II) oxide remaining after the reaction was complete.
28. Lead reacts with hydrochloric acid as follows:
Pb(s) + 2HCl(aq) ( PbCl2(aq) + H2(g)
If 6.22g of lead was added to 50cm3 of 1 mol l-1 hydrochloric acid, calculate the mass of lead left unreacted.
29. A strip of zinc metal weighing 2.00 g is placed in an aqueous solution containing 10.00 g of silver nitrate. The reaction that occurs is
Zn(s) + 2AgNO3(aq) ( 2Ag(s) + Zn(NO3)2(aq)
(a) Determine which reactant is in excess.
(b) Calculate how many grams of silver will be formed.
30. A piece of lithium with a mass of 1.50 g is placed in an aqueous solution containing 6.00 g of copper (II) sulphate. The reaction that occurs is:
2Li(s) + CuSO4(aq) ( Cu + Li2SO4 (aq)
(a) Determine which reactant is in excess.
(b) Calculate how many grams of copper will be formed.
EQUILIBRIUM
Many reactions are reversible. A reversible reaction can reach equilibrium in a closed system. A reaction reaches equilibrium when the rate of the forward reaction equals the rate of the reverse reaction.
reactants ⇌ products
At equilibrium, the concentration of the products and the reactants will remain constant.
The concentration of reactants will probably not equal the concentration of the products.
Where an industrial process produces an equilibrium, costly reactants may not be completely converted into products; chemists try to manipulate the equilibrium to achieve the best possible conversion rate.
Catalysts increase the rate at which an equilibrium is formed but do not affect the equilibrium position.
The equilibrium position will be the same whether we start with only the products or only the reactants. This can be shown in the following experiment.
Iodine is soluble in both water and cyclohexane. Water and cyclohexane do not mix. Iodine sets up an equilibrium between the 2 layers.
Iodine ( cyclohexane) ⇌ Iodine ( aq)
Draw diagrams to show
Iodine in cyclohexane/ water iodine(aq) /cyclohexane equilibrium
➢ Explain the changes that you see.
➢ Show on the graph the time at which equilibrium is reached.
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Le Chateliers Principle
An equilibrium will move to undo any change imposed upon it by temporarily favouring either the forward or backward reaction until a new equilibrium position is reached.
reactants ⇌ products
If the forward reaction is favoured we say the equilibrium has moved to the right.
If the reverse reaction is favoured we say the equilibrium has moved to the left.
Temperature may alter the position of equilibrium
N2O4 ⇌ 2NO2 ∆H = +ve
dinitrogen tetraoxide nitrogen dioxide
(colourless) (brown)
Increasing the temperature will cause the equilibrium to move to ........................... the temperature.
The .......................... reaction takes in energy so the equilibrium moves to the ............................ producing more .............................. and less ............................ So the reaction becomes ........................
Decreasing the temperature will cause the equilibrium to ........................ the temperature. The ...................... reaction gives out energy so the equilibrium shifts to the ............................ producing more ........................ and less ................................ So the colour becomes ............................
A mixture of cobalt chloride and conc HCl sets up the following equilibrium:
Co(H2O)6 2- + 4Cl- ⇌ CoCl4 - + 6 H2O ∆H = +ve
Pink Blue
If the temperature is increased the equilibrium will favour the .................... reaction because that will lower the temperature. The equilibrium move to the ....................., therefore the solution becomes ...................... in colour.
If the temperature is reduced the equilibrium will favour the .......................... reaction because that will increase the temperature. The equilibrium move to the ................, therefore the solution becomes ......................... in colour.
Concentration may alter the position of equilibrium
Add concentrated HCl to the equilibrium mixture
Co(H2O)6 2- + 4Cl- ⇌ CoCl4 - + 6 H2O
Pink Blue
Adding extra Cl- ions forces the equilibrium to try to remove these. The ____________ reaction is favoured because this uses up _________ ions.
The equilibrium has moved to the ____________ so the solution becomes ___________ in colour.
What would happen if extra CoCl4 - were added?
Add 10 cm3 iron (III) chloride to a test tube. Iron (III) ions are yellow. Add potassium thiocyanate solution until the solution goes orange. Red coloured iron thiocyanate ions form. The equilibrium position now lies in the middle, roughly equal amounts of both coloured ions are present.
Fe3+ + CNS- ⇌ [FeCNS]2+
yellow red
Divide the mixture between 4 test tubes, A, B, C and D
A Leave as control
B. Add FeCl3
C Add KCNS
D Add sodium chloride (removes Fe3+ by complexing)
a. Fully explain the colour changes you see in B
b. Fully explain the colour changes in C.
c. Fully explain the colour change in D.
ICl + Cl2 ⇄ ICl3
brown liquid yellow solid
Increasing the concentration of a chemical will cause the equilibrium to move to use up the chemical.
Increasing the concentration of chlorine will cause the equilibrium to move to ........ .... the chlorine. The .................. reaction uses up the chlorine so the equilibrium moves to the ............ producing ............... yellow solid and ................. brown liquid.
Decreasing the concentration of a chemical will cause the equilibrium to move to form the chemical.
Decreasing the concentration of chlorine will cause the equilibrium to move to............. the chlorine. The ........................ reaction produces chlorine so the equilibrium moves to the ............... producing ............... brown liquid and ................ yellow solid.
Effect of pressure ( Equilibria involving gases)
N2O4 ⇌ 2NO2
dinitrogen tetraoxide nitrogen dioxide
(colourless) (brown)
1 mol 2 mol
fewer particles more particles
[pic] Lower pressure [pic] Higher pressure
Increasing the pressure will cause the equilibrium to move to ............................. the pressure. The equilibrium will move to .............................. the number of gas particles.
The equilibrium moves to the .................. producing more ........................ and less ........................... so the colour ..................................
Decreasing the pressure will cause the equilibrium to move to ................ the pressure. The equilibrium will move to ...................... the number of gas particles.
The equilibrium moves to the .................. producing more ......... and less ............ so the colour ........................
Catalyst and equilibrium
A catalyst speeds up a reaction by lowering the activation energy. However in a reversible reaction it reduces the activation energy for both the forward and reverse reactions by hte same amount. Thus a catalyst speeds up both the reactions to the same extent.
The use of a catalyst does not change the equilibrium it only enables the position of equilibrium to be reached more quickly.
Equilibria Questions
1. The reaction between SO2 and O2 may be described as
2SO2(g) + O2(g) [pic] 2SO3(g) ΔH = -196 kJ mol-1
Select the two conditions that will favour a high yield of SO3.
A Removal of SO2
B Lower the pressure
C Lower the temperature
D Use a catalyst
E Addition of more O2
2. A chemical reaction has reached dynamic equilibrium at a certain temperature. Which one of following statements is incorrect?
A The reaction has stopped completely
B The concentrations of the reactants remains constant
C Products are continuously being formed
D The rate of the forward reaction is equal to the rate of the reverse reaction
3. Three moles of ethanol and three moles of ethanoic acid were reacted together according to The equation
C2H5OH + CH3COOH [pic] CH3COOC2H5 + H2O
At equilibrium, there was 2 moles each of ethyl ethanoate and water formed. What is the equilibrium constant for this reaction?
A 4
B 2
C 0.25
D 0.44
4. The diagram shows the concentrations of hydrogen, iodine and hydrogen iodide for the reaction between hydrogen and iodine.
Which of the following statements is incorrect?
A the equilibrium lies predominantly to the left
B at point A on the time axis, the concentration of all three gases is zero
C The reaction between the gases reaches equilibrium at point B.
D Adding more hydrogen at point D will alter the shape of the graph.
E At point C, the system is in a state of dynamic equilibrium.
5. Consider the following equilibrium reaction
[pic]
Which one of the following will cause a yellow colour to predominate?
A Addition of sodium hydroxide (NaOH)
B Addition of sodium chromate (Na2CrO4)
C Addition of hydrochloric acid (HCl)
D Removal of water
6. Which one of the following equilibrium reactions is not affected by a change in pressure?
A
B
C
D
7. What does it mean to say that a chemical reaction has reached equilibrium?
8. The forward reaction of the equilibrium system below is endothermic. Dilute HCl is added and the colour changes to blue. What colour change occurs when the mixture is cooled?
[pic]
9. The reaction of nitrogen dioxide (NO2) forming dinitrogen tetroxide (N2O4) is exothermic in the forward direction.
[pic]
State two conditions will cause the equilibrium mixture to go dark brown?
10. Ammonia is formed in the Haber process according to the
following balanced equation.
N2 + 3H2 [pic] 2NH3
The table shows the percentages of ammonia present at
equilibrium under different conditions of temperature T and
pressure P when hydrogen and nitrogen gases were mixed in
a 3:1 molar ratio. Is this an endothermic or exothermic
reaction? Give a reason for your answer.
11. The following equilibrium involves two compounds of phosphorus.
PCl3(g) + 3NH3 (g) [pic] P(NH2)3(g) + 3HCl(g)
(a) An increase in temperature moves the equilibrium to the left. What does this indicate about the enthalpy change for the forward reaction?
(b) What effect, if any, will an increase in pressure have on the equilibrium?
12. The balanced equation for a reaction at equilibrium is:
aA + bB [pic] cC + dD
(a) For this reaction, the equilibrium constant, K, can be defined as:
[pic]
where [A] represents the concentration of A, etc and a represents the number of moles of A,etc.
(i) Write down the expression for the equilibrium constant for the following equilibrium.
N2(g) + 3H2(g) [pic] 2NH3(g)
(ii) What will happen to the position of the equilibrium if the reaction is carried out over a catalyst?
(b) In industry, the reaction of nitrogen with hydrogen to produce ammonia by the Haber Process does not attain equilibrium.
Give one feature of the operating conditions which leads to the Haber Process not reaching equilibrium
13. When a yellow solution of iron (III) chloride (FeCl3)and a colourless solution of potassium thiocyanate (KCNS) were mixed in a test tube, a red colour appeared and the following equilibrium was established:
Fe3+(aq) + CNS-(aq) [pic] Fe(CNS)2+(aq)
yellow red
Explain:
(a) the effect on the Fe3+ ion concentration of adding KCNS to the equilibrium mixture
(b) why changing the pressure has no effect on this reaction.
14. Consider the following equilibrium reaction at room temperature used to dissolve iodine (I2)crystals in an aqueous solution of iodide ions (I-).
I2(aq) + I -(aq) [pic] I3-(aq)
State and explain the effect on the equilibrium concentration of triiodide ions of adding a substance that reacts with iodine, eg. starch.
PERCENTAGE YIELD
In a reversible reaction, there will be less than 100% conversion of reactants into products. The percentage yield is a measure of how much of a product is obtained compared to the amount expected if there was complete conversion.
• The actual yield is the amount that is obtained
• The theoretical yield is the amount that would be obtained assuming full conversion of the limiting reagent
• The percentage yield = actual yield x 100
theoretical yield
Example
Calcium benzoate, a food preservative (E213) can be made from the reaction between calcium carbonate and benzoic acid. If 2.44g of benzoic acid was reacted with excess calcium carbonate, what is the percentage yield if 1.75g of E213 was obtained?
Step 1 use the balanced equation to work out the theoretical yield
Benzoic acid + calcium carbonate ↔ calcium benzoate + water + carbon dioxide
2C7H6O2 + CaCO3 ↔ (C7H5O2 )2 Ca + H2O + CO2
Step 2 Use the actual yield from the question and the theoretical yield to calculate the percentage yield
The percentage yield = actual yield x 100
theoretical yield
Percentage yield of Zinc sulphate
Carry out the preparation of zinc sulphate
| |Mass (g) |
|Mass of zinc oxide used | |
|Mass of filter paper | |
|Mass of filter paper and dry zinc oxide | |
|Mass of unreacted zinc oxide | |
|Mass of evaporating basin | |
|Mass of evaporating basin + zinc sulphate | |
|Mass of zinc sulphate (actual yield) | |
Yield calculation
1. Write the balanced chemical equation. Show the mole ratio.
2. Work out theoretical yield of zinc oxide.
3. Use the actual yield from your results and the theoretical yield to calculate the percentage yield.
Do the calculations on yield on the worksheet
ATOM ECONOMY
The atom economy measures the proportion of the total mass of all starting materials successfully converted into the desired product. It can be calculated using the formula
Atom economy = mass of desired products x 100
Total mass of reactants
Green Chemistry - Comparing 2 ways to prepare methanol
1. Using steam reforming of methane
Step 1 CH4 + H2O ↔ CO + 3H2
Step 2 CO + 2H2 ↔ CH3OH
Assuming that both the reactants in step 1 are present in 1mol quantities and that there is complete conversion
Atom economy = mass of desired products x 100
Total mass of reactants
= 32 x 100 = 94%
16 + 18
2. Via incomplete combustion of methane
Step 1 CH4 + ½ O2 ↔ CO + 2H2
Step 2 CO + 2H2 ↔ CH3OH
➢ Work out the atom economy
The atom economy can help reveal how successfully all the products are converted into the desired product. This may reveal that the system is producing large amounts of unwanted waste products. One of the principles of Green Chemistry is that waste should be prevented by proper planning.
➢ Find out about the principles of Green Chemistry and how atom economy led to a change in the way ibuprofen is manufactured.
24. Percentage Yield and Atom Economy Questions
1. 20 g of lithium hydroxide was reacted with potassium chloride:
LiOH + KCl ( LiCl + KOH
(a) What is the theoretical yield of lithium chloride?
(b) If 6 g of lithium chloride was actually produced, what is the percentage yield?
2. The equation below shows the combustion of propanol:
C3H8 + 5 O2 ( 3 CO2 + 4 H2O
a) If you start with 5 grams of C3H8, what is the theoretical yield of water?
b) If the percentage yield was 75%, how many grams of water will actually be made?
3. In the reaction below, the theoretical yield was 10.7 g but the actual yield was 4.5 g.
Calculate the percentage yield.
Be + 2 HCl ( BeCl2 + H2
4. What is the theoretical yield of sodium oxide if you start with 20 grams of calcium oxide?
2 NaCl + CaO ( CaCl2 + Na2O
5. In the reaction below:
FeBr2 + 2 KCl ( FeCl2 + 2 KBr
a) What is the theoretical yield of iron (II) chloride if you start with 340 g of iron (II) bromide?
b) What is the percentage yield of iron (II) chloride if my actual yield is 40 g?
6. In the reaction below:
TiS + H2O ( H2S + TiO
What is the percentage yield of titanium (II) oxide if you start with 20 g of titanium (II) sulfide and the actual yield of titanium (II) oxide is 22 g?
7. In the reaction below:
U + 3 Br2 ( UBr6
What is the actual yield of uranium hexabromide if you start with 100 g of uranium and get a percentage yield of 83% ?
8. In the reaction below:
H2SO4 ( H2O + SO3
If you start with 89 kg of sulfuric acid and produce 71 kg of water, what is the percentage yield?
9. If, in the reaction below 32 kg of C2H6 produces 44 kg of CO2, what is the % yield?
2C2H6 + 7O2 ( 4CO2 + 6H2O
10. If, in the reaction below, 80 g of Cl2 produces 38 g of CCl4 what is the % yield?
CS2 + 3Cl2 ( CCl4 + S2Cl2
11. If, in the reaction below, 49 g of Fe3O4 produces a 78.25 % yield of Fe. How many grams are produced?
Fe3O4 + 4H2 ( 3Fe + 4H2O
12. If, in the reaction below, 40 tonnes of H2O produces 6.7 tonnes of HF what is the % yield?
CH3COF + H2O ( CH3COOH + HF
13. Calculate the atom economy for the production of lithium chloride assuming that all the reactants are converted into products.
LiOH + KCl ( LiCl + KOH
14. Calculate the atom economy for the production of titanium oxide assuming that all the reactants are converted into products.
TiS + H2O ( H2S + TiO
15. Calculate the atom economy for the production sulphur trioxide assuming that all the reactants are converted into products.
H2SO4 ( H2O + SO3
16. Which reaction below has the highest atom economy for producing water?
2C2H6 + 7O2 ( 4CO2 + 6H2O
C3H6 + 4½O2 ( 3CO2 + 3H2O
CHEMICAL ENERGY
Enthalpy (H) is a measure of the energy stored in a chemical.
MEASURING ENTHALPY CHANGES
The heat energy (Eh ) given to water by a burning fuel can be calculated using the formula
Eh = cm ∆T
c = Specific heat capacity of water
m = mass of water in Kg
∆T = change in temperature
ENTHALPHY OF COMBUSTION ALCOHOLS
The heat energy released when alcohols burn can be measured.
The enthalpy of combustion of a substance is the amount of energy given out when one mole of a substance burns in excess oxygen.
The enthalpy of combustion of a homologous series can be found using this method.
|Name of alcohol | | | | |
|Mass of burner before (g) | | | | |
|Mass of burner after (g) | | | | |
|Mass of alcohol used (g) | | | | |
|Mass of water heated (kg) | | | | |
|Temperature of water before (oC) | | | | |
|Temperature of water after (oC) | | | | |
|Change in temperature (oC) | | | | |
|Mass of one mole (g) | | | | |
|Enthalpy of combustion (kJmol-1) | | | | |
The heat energy gained by the water (Eh) is calculated using the formula:
Eh = c m ΔT
We assume that the heat energy released by the burning alcohol is gained only by the water.
The heat energy released on burning ……….. g of ……………anol ………….. kJ
So one mole .............. g of ................anol ....................kJ
The enthalpy of combustion of …………anol = -………….. kJ mol-1
(A negative sign is used because combustion is an exothermic reaction)
➢ Write a balanced equation to represent the enthalpy of combustion of ……anol
➢ Carryout similar calculations and write balanced equations for the other alcohols used.
➢ Compare your values with the values given in the data booklet. Explain the large differences and state how you could improve your experiment.
➢ Consider the possible source of error in the experiment
ENTHALPY OF SOLUTION
➢ What is the definition of the ‘Enthalpy of solution’?
➢ Write a balanced equation to represent
a. the enthalpy of solution of sodium hydroxide.
b. the enthalpy of solution of ammonium nitrate.
Carry out the practical measuring enthalpy of solution
| |mass (g) |start temp |final temp |change in temp |
| | |( oC) |( oC) |( oC) |
|sodium hydroxide | | | | |
|ammonium nitrate | | | | |
➢ Calculate Eh = cm ∆T for both solutes
➢ If we know how much energy is given out by a certain mass, we can work out ∆H.
➢ How do your answers compare to the known values?
Chemical Energy Questions
Enthalpy
1. Calculate the quantity of heat required to raise the temperature of
(a) 1 kg of water by 150C
(b) 250 cm3 of water by 7.80C
(c) 3 litres of water by 280C
2. A pupil burned 2.4g of sulphur in air to heat 150g of water. The temperature of the water increased from 15.30C to 36.00C.
Calculate the value for the enthalpy of combustion of sulphur using these experimental results.
3. A pupil found the enthalpy of combustion of propan-1-ol using the following apparatus.
(a) In addition to the initial and final temperatures of the water, what other measurements would the pupil have made.
(b) Describe a change that could be made to the experimental procedure in order to achieve more accurate results.
(c) The table shows the enthalpies of combustion of three alcohols.
|Alcohol |Enthalpy of combustion/kJ mol-1 |
|methanol |-715 |
|ethanol |-1371 |
|propan-1-ol |-2010 |
Why is there a regular increase in enthalpies of combustion from methanol to ethanol to propan-1-ol?
4. When 3.6 g of butanal (relative formula mass =72) was burned, 134 kJ of energy was released.
From this result, what is the enthalpy of combustion in kJ mol-1?
5. In an experiment the burning of 0.980g of ethanol resulted in the temperature of 400cm3 of water rising from 14.2oC to 31.6oC.
Use this information to calculate the enthalpy of combustion of ethanol.
6. Calculate the enthalpy change for each of the following experiments.
(a) When 1 g of potassium carbonate dissolved in 10 cm3 of water the temperature increased by 5.60C.
(b) When 1 g of sodium nitrate dissolved in 10 cm3 of water the temperature fell by 5.60C.
7. From the results of question 6 calculate the enthalpy of solution for
a) potassium carbonate (6a)
b) sodium nitrate (6b)
8. The enthalpy change when 1 mole of sodium carbonate dissolves in water is 24.6 kJ mol-1. Calculate the mass of sodium carbonate which would produce a temperature rise of 9.20C when added to 25cm3 of water.
9. 2g of sodium hydroxide, NaOH, is dissolved in 0.125 kg of water causing the temperature to rise from 19oC to 23oC.
Calculate the enthalpy of solution of sodium hydroxide.
10. 14.9g of potassium chloride, KCl, is dissolved in 200cm3 of water causing the temperature to fall from 19.5oC to 15.5oC.
Calculate the enthalpy of solution of potassium chloride.
11. A student dissolved 10.0g of ammonium chloride in 200cm3 of water and found that the temperature of the solution fell from 23.2oC to 19.8oC.
Calculate the enthalpy of solution of ammonium chloride.
12. A pupil added 50cm3 of NaOH(aq) to 50cm3 HCl(aq). Each solution had a concentration of 2.0 mol l-1. The temperature rise was 13.5oC.
Calculate the enthalpy of neutralisation.
13. 40cm3 of 1 mol l-1 of nitric acid, HNO3, and 40cm3 of 1 mol l-1 sodium hydroxide, NaOH, both at room temperature of 19oC were mixed and the temperature increased to 25.8oC.
Calculate the enthalpy of neutralisation.
Enthalpy Diagrams
1. a) Copy the diagrams below and mark with an arrow:- i) the activation energy EA.
ii) the enthalpy change (
b) State whether each reaction is endothermic or exothermic.
c) Calculate the value of (H and EA for each reaction
2. Copy the axes below and sketch potential energy diagrams for the following reactions, labelling the axes.
a) (H = -15 kJ mol-1 EA = 20 kJ mol-1
b) (H = +20 kJ mol-1 EA = 35 kJ mol-1
[pic]
3. Two chemicals A and B react in solution to form C. The reaction has an activation energy of 150 kJ mol-1. If hydrogen ions are used as a catalyst the activation energy is 50 kJ mol-1. The enthalpy change for the reaction is -125 kJ mol-1.
Present this information as a potential energy diagram using the template below.
Use a solid line for the uncatalysed reaction and a dotted line for the catalysed reaction.
4. The graph shows the potential energy diagram for a urease catalysis of urea.
(a) What is the enthalpy change for the reaction?
(b) Acid is a less effective catalyst than urease for this reaction.
Add a curve to the potential energy diagram to show the hydrolysis when acid is used as the catalyst.
5. Hydrogen peroxide, H2O2, decomposes very slowly to produce water and oxygen.
(a) The activation energy (EA) for the reaction is
75 kJ mol-1 and the enthalpy change ((H) is
–26 kJ mol-1.
Copy the diagram and use the information above to complete the potential energy diagram for the reaction using a solid line.
(b) When a catalyst is used the activation energy is
reduced by 30 kJ mol-1.
Add a dotted line to the diagram to show the path of reaction when a catalyst is used.
HESS’S LAW
Hess Law
The overall enthalpy change for a reaction is the same whichever route is taken.
∆H = ∆H1 + ∆H2 + ∆H3
Carry out the practical Hess’s Law
Potassium hydroxide pellets can be neutralised by adding hydrochloric acid. This is the direct route, ∆H1. Alternatively, the pellets can be dissolved in water (enthalpy of solution ∆H2) and then the solution can be neutralised with hydrochloric acid (enthalpy of neutralisation, ∆H3).
KOH(s K Cl (aq) + H2O
KOH(aq)
It can be seen that ∆H1 = ∆H2 + ∆H3
Identify the types of enthalpy change involved.
∆H2 KOH(s) → K+ (aq) + OH - (aq)
∆H3 K+ (aq) + OH - (aq) + HCl → K+ Cl - (aq) + H2O
∆H1 KOH(s) + HCl → K+ Cl - (aq) + H2O
Carry out the practical and record the results in your jotter. Calculate ∆H in your jotter using the results.
∆H1=
∆H2 =
∆H3 =
1. Show by calculation how well your results verified Hess's Law.
2. What were the main sources of error?
Calculations based on Hess’s Law
Using the enthalpy of combustion, work out the enthalpy of formation of compounds.
The enthalpy of formation is the energy needed to make one mole of a compound from its’ elements in their standard state.
The enthalpy of formation of methane can be represented by the equation:
C (s) + 2H2 (g) → CH4(g)
e.g. find the enthalpy of formation of ethane using the enthalpies of combustion
Step 1
Write desired “key equation”
C(s) + 2H2(g) → CH4 (g)
Step 2
Write out the equations for enthalphy of combustion C(s) H2(g) CH4(g) and the values from databook..
C(s) + O2(g) → CO2(g) ΔH = -394 kJ
H2(g) + ½ O2(g) → H2O(l) ΔH = -286 kJ
CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH = -891 kJ
Step 3
Compare with desired equation and rearrange.
C(s) + O2(g) → CO2(g) ΔH = -394 kJ
2H2(g) + O2(g) → 2H2O(l) ΔH = (2x-286) kJ *
CO2(g) + 2H2O(l) → CH4(g) + 2O2(g) ΔH = +891 kJ **
*We needed 2H2 so double equation. (if you double equation double the energy change).
**We needed CH4 on the right hand side of the arrow so equation reversed. (if you reverse a reaction then you must change the sign).
Step 4
Cancel and add to produce the key equation
C(s) + O2(g) → CO2(g) ΔH = -394 kJ
2H2(g) + O2(g) → 2H2O(l) ΔH = (2x-286) kJ
CO2(g) + 2H2O(l) → CH4(g) + 2O2(g) ΔH = +891 kJ
C(s) + 2H2(g) → CH4 (g) ΔH = -75kJmol-1
➢ Calculate the enthalpy of formation of propane using the enthalpies of combustion.
➢ Calculate the enthalpy of formation of ethanol using the enthalpies of combustion.
➢ Calculate the enthalpy of formation of ethanoic acid using the enthalpies of combustion
Hess law can also be used to work out the enthalpy for an overall reaction from known equations.
e.g. What is the value of ∆H for the reaction
FeO + CO → Fe + CO2
CO2 + C → 2CO ∆ H = a
FeO + C → Fe + CO ∆ H = b
The first equation is reversed as the target equation needs to have CO as a reactant and carbon dioxide as a product so ∆ H = –a
The second equation goes as written because the target equation needs FeO as a reactant so ∆ H = b
Therefore ∆H = b + (- a) = b - a
Hess’s Law
1. What is the relationship between a, b, c and d? Answer in the form a = ………………………
S(s) + H2(g) ( H2S(g) (H = a
H2(g) + ½ O2 (g) ( H2O(l) (H = b
S(s) + O2(g) ( SO2(g) (H = c
H2S(g) + 1 ½ O2(g) ( H2O(l) +SO2(g) (H = d
2. The enthalpy changes for the formation of one mole of aluminium oxide and one mole of iron(III) oxide are shown below.
2Al(s) + 1½O2(g) ( Al2O3(s) (H = -1676 kJ mol-1
2Fe(s) + 1½O2(g) ( Fe2O3(s) (H = -825 kJ mol-1
Use the above information to calculate the enthalpy change for the reaction:
2Al(s) + Fe2O3(s) ( Al2O3(s) + 2Fe(s)
3. The equation for the enthalpy of formation of propanone is:
3C(s) + 3H2(g) + ½O2(g) C3H6O(l)
Use the following information on enthalpies of combustion to calculate the enthalpy of formation of propanone.
C(s) + O2(g) CO2(g) (H = -394 kJmol-1
H2(g) + ½O2(g) H2O(l) (H = -286 kJmol-1
C3H6O(l) + 4O2(g) 3CO2(g) + 3H2O(l) (H = -1804 kJmol-1
4. The equation below represents the hydrogenation of ethene to ethane.
C2H4(g) + H2(g) C2H6(g)
Use the enthalpies of combustion of ethene, hydrogen and ethane from page 9 of the data booklet to calculate the enthalpy change for the above reaction.
5. Calculate a value for the enthalpy change involved in the formation of one mole of hydrogen peroxide from water (ΔH3). The enthalpy change when hydrogen forms hydrogen peroxide is -188 kJ mol-1 and the enthalpy of combustion of hydrogen to form water is -286 kJ mol-1.
6. Calculate a value for the enthalpy change involved in the decomposition of nitrogen dioxide to nitrogen monoxide given the following information.
Equation (a) N2(g) + O2(g) ( 2NO(g) ΔH = +181 kJ
equation (b) N2(g) + 2O2(g) ( 2NO2(g) ΔH = +68 kJ
BOND ENTHALPY
For a diatomic molecule, XY, the molar bond enthalpy is the energy required to break one mole of bonds. For a diatomic molecule, molar bond enthalpies can be measured directly. Bond breaking is an endothermic process.
Bond making is an exothermic process. The energy required to make one mole of bonds is the same as the bond enthalpy, but has a negative value.
e.g. energy to break one mole of H-H bonds = kJmol-1
energy to make one mole of H-H bonds = kJmol-1
Where bonds are present in a molecule with more than 2 atoms, the average bond enthalpy is worked out. The bond enthalpy will be affected by the environment the bond is in, so a C-H bond in methane may have a slightly different bond enthalpy from one in propene.
∆H can be calculated from bond enthalpies using the equation
(H = ( (H bonds broken + ( (H bonds made
e.g 2C + 2 H2 ( C 2H 4
sublime 2C 2 (715) make C=C -602
break 2( H-H) 2 ( 432) make 4 (C-H) -4(414)
2294 -2258
(H = ( (H bonds broken + ( (H bonds made
= 2294 + ( -2258 ) = 36 kJ mol-1
The equation can also be used to work out the average bond enthalpy. The enthalpy of formation of methane is -75 kJ mol-1. The example shows how the average bond enthalpy for a C-H bond in methane can be calculated:
C + 2 H2 ( CH 4
sublime C 715 make 4 C-H = 4x
break 2 H-H 2( 432)
+ 1579
(H = ( (H bonds broken + ( (H bonds made
- 75 = 1579 + 4x
-1674 = 4x
-413.5 = x
( bond enthalpy for C-H bond is 413.5 kJ mol-1
Bond Enthalpies
1. In the presence of bright light, hydrogen and bromine react. One step in the reaction is shown below.
H2(g) + Br(g) → HBr(g) + H(g)
The enthalpy change for this step can be represented as
A (H-H bond enthalpy) + (Br-Br bond enthalpy)
B (H-H bond enthalpy) − (Br-Br bond enthalpy)
C (H-H bond enthalpy) + (H-Br bond enthalpy)
D (H-H bond enthalpy) − (H-Br bond enthalpy).
2. Use the information in the table to calculate the enthalpy change for the following reaction:
H2(g) + Cl2(g) ( 2HCl(g)
|Bonds |ΔH to break bond (kJ mol-1) |
|H-H |432 |
|Cl-Cl |243 |
|H-Cl |428 |
3. Using the bond enthalpy values from your data booklet, calculate the enthalpy changes for the following reactions:
(a) CH4(g) + 2O2(g) ( CO2(g) + 2H2O(g)
(b) C3H8(g) + 5O2(g) ( 3CO2(g) + 4H2O(g)
(c) C3H6(g) + H2(g) ( C3H8(g)
(d) N2(g) + 2O2(g) ( 2NO2(g)
Redox Reactions
Displacement reactions are examples of redox
1. What is meant by
oxidation
reduction
redox
The substance that is oxidised ‘gives’ electrons to the substance that is reduced.
2. Writing a redox reaction by combining half equations.
If magnesium is added to a solution containing Fe(III) ions, a displacement reaction occurrs.
➢ Write the equation for the oxidation of Magnesium metal
➢ Write the equation for the reduction of iron (III) ions to Fe atoms
➢ Multiply through so the number of electrons is the same on both sides Write the overall redox reaction.
➢ How many electrons were transferred?
In your jotter, write redox equations for
a. Displacement of silver ions by zinc
b. Displacement of copper ions by aluminium.
In your jotter, identify the 2 half equations and say which is oxidation and which is reduction for
a. Zn + 2Ag + → 2Ag + Zn 2+
b. I2 + SO3 2- + H20 → 2I – + SO4 2- + 2H +
c. 2I – + 2Fe 3+ → I2 + 2Fe 2+
OXIDISING AGENTS AND REDUCING AGENTS
Oxidising agents remove electrons from other chemicals forcing them to lose electrons. Oxidising agents are found on the bottom left hand side of the ECS. Oxidising agents are reduced in the reaction.
1. Acidified potassium dichromate is an excellent oxidising agent e.g. oxidation of alcohol in the breathalyzer test. The dichromate ion is reduced in the process. Write the half equation for the dichromate ion when it acts as an oxidising agent.
Reducing agents donate electrons to other chemicals, causing them to become reduced. The reducing agent is itself oxidized in the process. Reducing agents are found at the top right of the ECS.
2. Write the half equation for lithium acting as a reducing agent.
2. In the equations below, circle the reducing agent and underline the oxidising agent.
Zn + 2Ag + → 2Ag + Zn 2+
I2 + SO3 2- + H20 → 2I – + SO4 2- + 2H +
2I – + 2Fe 3+ → I2 + 2Fe 2+
2Br – + SO4 2- + 2H + → Br 2 + SO3 2- + H20
MnO4 – + 8H + + 5Fe 2+ → Mn2+ + 4H20 + 5Fe 3+
Elements as oxidising or reducing agents
The elements with low electronegativities (metals) tend to form ions by losing electrons (oxidation) and so can act as reducing agents. The strongest reducing agents are found in group 1.
The elements with high electronegativities (non-metals) tend to form ions by gaining electrons (reduction) and so can act as oxidising agents. The strongest oxidising agents are found in group 7.
Comparing the strength of oxidising agents from group 7.
Which halogen would you expect to be the best oxidising agent? Why?
Answer the questions
o Which halogen solution is the strongest bleaching agent?
o Which halogen is the most reactive?
o Write symbol equation for the reaction of chlorine with potassium bromide.
o Rewrite the above equation omitting spectator ions
o Write a half equation for the oxidising agent in the above reaction
Molecules and group ions can act as oxidising and reducing agents
Eg MnO4- (aq) + 8H+ (aq) + 5e- → Mn2+ (aq) + 4H2O(l)
1. Is MnO4- acting as a reducing or oxidising agent in the reaction with glycerol?
2. The dichromate and permanganate ions are strong oxidising agents in acidic solutions. How is this shown in the half equation?
Hydrogen peroxide is an example of a molecule which is a strong oxidising agent. Carbon monoxide is an example of a gas that can be used as a reducing agent. Oxidising and reducing agents can be selected using an electrochemical series from a databook or can be identified in the equation showing a redox reaction.
3. Carry out the blue bottle experiment.
Methylene blue is reduced by dextrose to a colourless compound. This can be oxidized back to the blue form by a gas. Which gas is acting as the oxidizing agent?
Explain why is it necessary to periodically remove the stopper?
4. Elephants toothpaste. Write the equation for the decomposition of hydrogen peroxide.
A slight brown colour can be seen as the iodide ions are converted to iodine molecules by the hydrogen peroxide. Is hydrogen peroxide acting as a reducing agent or an oxidizing agent?
H2O2 + 2H+ + 2 I- → 2 H2O + I2-
EVERYDAY USES FOR STRONG OXIDISING AGENTS
Oxidising agents are widely employed because of the effectiveness with which they can kill fungi, and bacteria and can inactivate viruses. The oxidation process is also an effective means of breaking down coloured compounds making oxidising agents ideal for use as "bleach" for clothes and hair.
Research one chemical used in the bleaching of clothes or hair and one that is used in antibacterial solutions.
HOW TO WRITE HALF EQUATIONS
Oxidation and reduction reactions can be represented by ion-electron equations.
When molecules or group ions are involved, if the reactant and product species are known, a balanced ion-electron equation can be written by adding appropriate numbers of water molecules, hydrogen ions and electrons.
The sequence is usually:
• Balance the atoms apart from oxygen and hydrogen.
• Balance the oxygen and hydrogen atoms by adding hydrogen ions/water
• Balance the charges by adding electrons to the most positive side
e.g. CIO3 - → CI2
Balance the equation for the element that is oxidised/reduced
2CIO3 - → CI2
Add enough H+ ions to convert O to water. Balance this part:
2CIO3 - + 6H + → CI2 + 3H20
Total up the charges on each side and add electrons to the most positive side to equalise the charge.
2 - + 6+ = 4+ → no charge
therefore add 4 electrons to the left hand side.
CIO3 - + 6H + + 4e - → CI2 + 3H20
Oxidising and Reducing Agents
1. Which of the following is a redox reaction?
A NaOH + HCl ( NaCl + H2O
B Zn + 2HCl ( ZnCl2 + H2
C NiO + 2HCl ( NiCl2 + H2O
D CuCO3 + 2HCl ( CuCl2 + H2O + CO2
2. During a redox process in acid solution, iodate ions, IO3- (aq) , are converted into iodine, I2 (aq).
IO3- (aq) ( I2 (aq)
The numbers of H+ (aq) and H2O(l) required to balance the ion-electron equation for the formation of 1 mol of I2(aq) are, respectively
A 6 and 3
B 3 and 6
C 12 and 6
D 6 and 12
3. Iodide ions can be oxidised using acidified potassium permanganate solution. The equations are:
2I-(aq) ( I2(aq) + 2e-
MnO4-(aq) + 8H+(aq) + 5e- (Mn2+(aq) + 4H2O(l)
How many moles of iodide ions are oxidised by one mole of permanganate ions?
A 1.0
B 2.0
C 2.5
D 5.0
4. What is the significance of the acronym ‘OILRIG’ when explaining a redox process.
5. What is meant by a spectator ion?
6. For the following displacement reactions write down the relevant ion-electron equations and use them to work out the redox equation. Do not include the spectator ions.
(a) copper metal reacts with silver(I) nitrate solution to form copper (II) nitrate solution and silver.
(b) chromium metal reacts with nickel (II) sulphate solution to form chromium (III) sulphate solution and nickel.
(c) magnesium metal displaces aluminium from aluminium (III) oxide.
(d) copper is displaced from a solution of copper (II) sulphate by sodium metal.
7. Give the names of two strong oxidising agents and give two uses of each.
8. The ion-electron equations below represent the reduction and oxidation reactions which take place when an acidified solution of dichromate ions react with sulphite ions.
Cr2O72-(aq) + 14H+(aq) + 6e- 2Cr3+(aq) + 7H2O(l)
SO32-(aq) + H2O(l) SO42-(aq) + 2H+(aq) + 2e-
Write the REDOX equation for this reaction.
9. Sulphur dioxide is added to wine as a preservative. A mass of 20 to 40 mg of sulphur dioxide per litre of wine will safeguard the wine without affecting its taste.
(a) Describe clearly, with full experimental detail, how 0.05 mol l-1 iodine solution would be
diluted to give 250 cm3 of 0.005 mol l-1 solution.
(b) The equation for the reaction which takes place is:
SO2(aq) + I2(aq) + 2H2O(l) ( 4H+(aq) + SO42-(aq) + 2I-(aq)
(i) The indicator used in this reaction causes a change from blue to colourless at the end point. Name a substance which could be used as this indicator.
(ii) Write the ion-electron equation for the reduction reaction taking place.
10. (a) In acid solution, iodate ions, IO3-(aq), are readily converted into iodine. Write an ion-electron equation for this half-reaction.
(b) Use the equation to explain whether the iodate ion is an oxidizing or a reducing agent.
ANALYSIS - CHROMATOGRAPHY
Chemical analysis has a wide range of applications. In industry it is used to check the composition and purity of reactants and products.
Chromatography separates compounds according to their relative affinity for the ‘mobile phase’ and the ‘stationary phase’.
The mobile phase is a liquid or a gas. The size of molecules and their polarity may affect how soluble they are in the mobile phase
The stationary phase may be paper, silica gel, or an inert packing material. The size and polarity of the compounds may affect their affinity for the stationary phase.
Every compound will have a unique ‘fingerprint’ for each type of chromatography
➢ How are compounds identified in paper chromatography or TLC?
➢ How are compounds identified in column chromatography methods?
Which graph, A, B or C trace shows the smallest molecules? How do you know?
Chromatography
1. Use the diagram showing a paper chromatography experiment to
define the following terms:
(a) mobile phase
(c) stationary phase
(d) Rf value
2. Compare and explain the speed at which the following move up the paper in paper
chromatography.
(a) Large molecules compared with small molecules.
(b) A polar solvent compared with a non-polar solvent.
3. An organic chemist is attempting to synthesise a fragrance compound by the following chemical reaction.
compound X + compound Y → fragrance compound
After one hour, a sample is removed and compared with pure samples of compounds X and Y using thin-layer chromatography.
Which of the following chromatograms shows that the reaction has produced a pure sample of the fragrance compound?
4. Describe how chromatography can be used to identify the amino acids that make up a protein.
5. Label the parts A – F on the gas chromatography equipment below:
6. In terms of gas liquid chromatography
(a) what is the mobile phase?
(b) what is the stationary phase?
(c) why is the injection port heated?
(d) explain what is meant by retention time.
7. Give 3 different uses of gas liquid chromatography.
8. (a) Which gases are usually used as carrier gases in gas chromatography?
(b) Explain why these particular gases are used.
9. If the stationary phase in gas chromatography is non-polar, how would the retention times of polar and non-polar samples in the column compare to each other?
10. A technician analyses a mixture of hydrocarbons using gas chromatography.
She first calibrates the equipment using standard hydrocarbons. The retention times of these hydrocarbons are shown in the table.
|hydrocarbon |formula |retention time in minutes |
|methane |CH4 |1.7 |
|ethane |C2H6 |2.2 |
|propane |C3H8 |3.5 |
|butane |C4H10 |4.0 |
|pentane |C5H12 |7.4 |
The technician then analyses the mixture of hydrocarbons. The recorder print out from this analysis is shown below.
(a) How does the recorder print out show that butane has the highest concentration?
(b) Use data in the table to draw a conclusion relating the formula of each hydrocarbon to its retention time.
CHemical Analysis – volumetric titration
Volumetric analysis involves using a solution of known concentration to determine the amount of another substance present. The volume of the reactant needed to complete the reaction is determined by titration. An indicator may be needed to show the ‘end point’ or the point at which the reaction is just complete
A solution of accurately known concentration is known as a standard solution.
➢ Calculate the exact concentration of a solution of sodium hydroxide made by dissolving 4.02 g of NaOH in 500.0 cm3 of water in a standard flask.
REdox Titrations
Iron tablets contain iron(II) sulphate. Potassium permanganate can be used to analyse the tablets to determine (find out) how much iron a tablet contains.
In the reaction, the oxidising agent potassium permanganate oxidises iron(II) ions to iron(III) ions.
What to do.
Accurately weigh out one iron tablet.
Dissolve the tablet in about 100 cm3 water using heater and a stirrer. Transfer the solution to a 250.0 cm3 standard flask. Wash the beaker repeatedly and transfer the washings to the standard flask. Top up the level to the 250.0cm3 mark.
Fill the burette with acidified potassium permanganate. Using a pipette, transfer 25.0 cm3 of the iron solution into a conical flask. Titrate the solution until the pink colour of the permanganate ion is just seen. Repeat at least twice to get 2 concordant accurate titres.
| |Rough titre |Accurate titre 1 |Accurate titre 2 |Accurate titre 3 |
|Initial burette reading (cm3) | | | | |
|Final burette reading (cm3) | | | | |
|Volume used (cm3) | | | | |
Average of the concordant titres (within 0.2cm3) =
Mass of one tablet =
1 Write the 2 half equations. Multiply through so the number of electrons is the same on both sides of the equation. Combine the 2 half equations to give the overall redox equation.
2. Calculate the number of moles of iron sulphate in solution
3. Calculate the mass of iron sulphate in the tablet
4. Why is no indicator needed for the reaction.
Redox titration to determine vitamin C content
Volumetric Analysis
1. 25 cm3 of a solution of sodium hydroxide was added to a flask and titrated with a 0.2 mol l-1 solution of hydrochloric acid.
HCl + NaOH → NaCl + H2O
The experiment was carried out three times and the volumes of HCl titrated in each experiment are shown in the table.
|Titration |Volume of 0.2 mol l-1 solution of HCl (cm3) |
|1 |11.3 |
|2 |10.4 |
|3 |10.6 |
Calculate the concentration of the NaOH solution in mol l –1.
2. 20 cm3 of a solution of potassium hydroxide was added to a flask and titrated with a 0.1 mol l-1 solution of hydrochloric acid.
HCl + KOH → KCl + H2O
The experiment was carried out three times and the volumes of HCl titrated in each experiment are shown in the table.
|Titration |Volume of 0.2 mol l-1 solution of HCl (cm3) |
|1 |20.6 |
|2 |19.9 |
|3 |20.0 |
Calculate the concentration of the KOH solution in mol l –1.
3. 10 cm3 of a solution of KOH was added to a flask and titrated with a 0.05 mol l-1 solution of H2SO4.
H2SO4 + 2KOH → K2SO4 + 2H2O
The experiment was carried out three times and the volumes of HCl titrated in each experiment are shown in the table.
|Titration |Volume of 0.2 mol l-1 solution of HCl (cm3) |
|1 |15.9 |
|2 |15.2 |
|3 |15.3 |
Calculate the concentration of the KOH solution in mol l –1.
4. Rhubarb leaves contain oxalic acid, (COOH)2. A pupil found that it required 17 cm3 of 0.001 mol l-1 of sodium hydroxide to neutralise 25 cm3 of a solution made from rhubarb
leaves. Calculate the concentration of oxalic acid in the solution given that the equationfor the reaction is:
(COOH)2 + 2NaOH ( Na2(COO)2 + 2H2O
5. Acidified potassium permanganate can be used to determine the concentration of hydrogen peroxide solution; the solutions react in the ratio of
2 mol of potassium permanganate: 5mol of hydrogen peroxide.
In an analysis it is found that 16.8 cm3 of 0.025 mol l-1 potassium permanganate reacts exactly with a 50 cm3 sample of hydrogen peroxide solution. What is the concentration, in mol l-1 of the hydrogen peroxide solution?
6. Iodine reacts with thiosulphate ions as follows:
I2(aq) + 2S2O32-(aq) 2I -(aq) + S4O62-(aq)
In an experiment it was found that 1.2 x 10-5 mol of iodine reacted with 3.0 cm3 of the sodium thiosulphate solution. Use this information to calculate the concentration of the thiosulphate solution in mol l-1.
7. Vitamin C, C6H8O6, is a powerful reducing agent. The concentration of vitamin C in a solution can be found by titrating it with a standard solution of iodine, using starch as an indicator. The equation for the reaction is:
C6H8O6(aq) + I2(aq) C6H6O6(aq) + 2H+(aq) + 2I-(aq)
A vitamin C tablet was crushed and dissolved in some water. The solution was then transferred to a standard 250 cm3 flask and made up to the 250 cm3 mark with distilled water.
In one investigation it was found that an average of 29.5 cm3 of 0.02 mol l-1 iodine solution was required to react completely with 25.0 cm3 of vitamin C solution.
Use this result to calculate the mass, in grams, of vitamin C in the tablet.
8. Hydrogen sulfide, H2S, can cause an unpleasant smell in water supplies. The concentration of hydrogen sulfide can be measured by titrating with a chlorine standard solution.
The equation for the reaction taking place is
4Cl2(aq) + H2S(aq) + 4H2O(l) → SO42−(aq) + 10H+(aq) + 8Cl−(aq)
50·0 cm3 samples of water were titrated using a 0∙010 mol l−1 chlorine solution.
(a) Name an appropriate piece of apparatus which could be used to measure out the water samples.
(b) What is meant by the term standard solution?
(c) An average of 29·4 cm3 of 0∙010 mol l−1 chlorine solution was required to react completely with a 50·0 cm3 sample of water. Calculate the hydrogen sulfide concentration, in mol l−1,present in the water sample. Show your working clearly.
9. A compound known as ethylenediaminetetraacetic acid(EDTA) is useful for measuring the quantities of certain metal ions in solution. For example, Ca2+ ions and EDTA react in a 1 mol:1 mol ratio.
It is found that 14.6 cm3 of 0.1 mol l-1 EDTA reacts exactly with a 25cm3 sample of a solution containing Ca2+ ions. Calculate the concentration, in mol l-1, of the calcium ion solution.
Exam questions
3.1 Getting the Most from Reactants
1. How many moles of oxygen atoms are in 0.5 mol of carbon dioxide?
A 0.25
B 0.5
C 1
D 2
2. A fullerene molecule consists of 60 carbon atoms.
Approximately how many such molecules are present in 12 g of this type of carbon?
A 1.0 × 1022
B 1.2 × 1023
C 6.0 × 1023
D 3.6 × 1025
3. Which of the following gases would contain the greatest number of molecules in a 100 g
sample, at room temperature?
A Fluorine
B Hydrogen
C Nitrogen
D Oxygen
4. A mixture of potassium chloride and potassium carbonate is known to contain
0·1 mol of chloride ions and 0·1 mol of carbonate ions. How many moles of potassium ions are present?
A 0·15
B 0·20
C 0·25
D 0·30
5. The mass of 1 mol of sodium is 23 g. What is the approximate mass of one sodium
atom?
A 6 × 1023 g
B 6 × 10–23 g
C 3.8 × 10–23 g
D 3.8 × 10–24 g
6. Which of the following gas samples has the same volume as 7 g of carbon monoxide?
(All volumes are measured at the same temperature and pressure.)
A 1 g of hydrogen
B 3.5 g of nitrogen
C 10 g of argon
D 35 . 5 g of chlorine
7. In which of the following pairs do the gases contain the same number of oxygen atoms?
A 1 mol of oxygen and 1 mol of carbon monoxide
B 1 mol of oxygen and 0.5 mol of carbon dioxide
C 0.5 mol of oxygen and 1 mol of carbon dioxide
D 1 mol of oxygen and 1 mol of carbon dioxide
8. The Avogadro Constant is the same as the number of
A molecules in 16 g of oxygen
B ions in 1 litre of sodium chloride solution, concentration 1 mol l–1
C atoms in 24 g of carbon
D molecules in 2 g of hydrogen.
9. Which of the following contains one mole of neutrons?
[pic]
10. The Avogadro Constant is the same as the number of
A ions in 1 mol of NaCl
B atoms in 1 mol of hydrogen gas
C electrons in 1 mol of helium gas
D molecules in 1 mol of oxygen gas.
11. The Avogadro Constant is the same as the number of
A molecules in 16 g of oxygen
B electrons in 1 g of hydrogen
C atoms in 24 g of carbon
D ions in 1 litre of sodium chloride solution, concentration 1 mol l–1.
12. Avogadro’s Constant is the same as the number of
A molecules in 16.0 g of oxygen
B atoms in 20.2 g of neon
C formula units in 20.0 g of sodium hydroxide
D ions in 58.5 g of sodium chloride.
13. Which of the following has the largest volume under the same conditions of temperature and pressure?
A 1 g hydrogen
B 14 g nitrogen
C 20·2 g neon
D 35·5 g chlorine
14. The equation for the complete combustion of propane is:
C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l)
30cm3 of propane is mixed with 200cm3 of oxygen and the mixture is ignited. What is the volume of the resulting gas mixture? (All volumes are measured at the same temperature and pressure.)
A 90cm3
B 120cm3
C 140cm3
D 210cm3
15. 20cm3 of butane is burned in 150 cm3 of oxygen.
C4H10(g) + 6O2(g) → 4CO2(g) + 5H2O(g)
What is the total volume of gas present after complete combustion of the butane?
A 80cm3
B 100cm3
C 180cm3
D 200cm3
16. 2NO(g) + O2(g) → 2NO2(g)
How many litres of nitrogen dioxide gas could theoretically be obtained in the reaction of 1 litre of nitrogen monoxide gas with 2 litres of oxygen gas? (All volumes are measured under the same conditions of temperature and pressure.)
A 1
B 2
C 3
D 4
17. 2NO(g) + O2(g) → 2NO2(g)
How many litres of nitrogen dioxide gas would be produced in a reaction, starting with
a mixture of 5 litres of nitrogen monoxide gas and 2 litres of oxygen gas? (All volumes are measured under the same conditions of temperature and pressure.)
A 2
B 3
C 4
D 5
18. What volume of oxygen (in litres) would be required for the complete combustion of a
gaseous mixture containing 1 litre of carbon monoxide and 3 litres of hydrogen?
(All volumes are measured at the same temperature and pressure.)
A 1
B 2
C 3
D 4
19. 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(l)
ethyne
What volume of gas would be produced by the complete combustion of 100 cm3 of ethyne gas? All volumes were measured at atmospheric pressure and room temperature.
A 200 cm3
B 300 cm3
C 400 cm3
D 800 cm3
20. 20cm3 of ammonia gas reacted with an excess of heated copper(II) oxide.
3CuO + 2NH3 → 3Cu + 3H2O + N2
Assuming all measurements were made at 200 °C, what would be the volume of gaseous
products?
A 10cm3
B 20cm3
C 30cm3
D 40cm3
21. Calcium carbonate reacts with nitric acid as follows.
CaCO3(s) + 2HNO3(aq) → Ca(NO3)2(aq) + H2O(l) + CO2(g)
0·05 mol of calcium carbonate was added to a solution containing 0·08 mol of nitric acid. Which of the following statements is true?
A 0·05 mol of carbon dioxide is produced.
B 0·08 mol of calcium nitrate is produced.
C Calcium carbonate is in excess by 0·01 mol.
D Nitric acid is in excess by 0·03 mol.
22. A mixture of magnesium bromide and magnesium sulfate is known to contain 3 mol
of magnesium and 4 mol of bromide ions. How many moles of sulfate ions are present?
A 1
B 2
C 3
D 4
23. 5 g of copper is added to excess silver nitrate solution. The equation for the reaction that takes place is:
[pic]
After some time, the solid present is filtered off from the solution, washed with water,
dried and weighed. The final mass of the solid will be
A less than 5 g
B 5g
C 10g
D more than 10 g.
24. A pupil added 0·1 mol of zinc to a solution containing 0·05 mol of silver(I) nitrate.
Zn(s) + 2AgNO3(aq) → Zn(NO3)2(aq) + 2Ag(s)
Which of the following statements about the experiment is correct?
A 0·05 mol of zinc reacts.
B 0·05 mol of silver is displaced.
C Silver nitrate is in excess.
D All of the zinc reacts.
25. 0·5 mol of copper(II) chloride and 0·5 mol of copper(II) sulphate are dissolved together in water and made up to 500 cm3 of solution. What is the concentration of Cu2+(aq) ions in the solution in mol l–1?
A 0·5
B 1·0
C 2·0
D 4·0
26. 10 g of magnesium is added to 1 litre of 1 mol l–1 copper(II) sulphate solution and the
mixture stirred until the reaction is complete. Which of these is a result of this reaction?
A All the magnesium reacts.
B 63 . 5 g of copper is displaced.
C 2 mol of copper is displaced.
D The resulting solution is colourless.
27. Ammonia is manufactured from hydrogen and nitrogen by the Haber Process.
3H2(g) + N2(g) [pic] 2NH3(g)
If 80 kg of ammonia is produced from 60 kg of hydrogen, what is the percentage yield?
[pic]
28. Two identical samples of copper(II) carbonate were added to an excess of 1 mol l–1
hydrochloric acid and 1 mol l–1 sulphuric acid respectively.
Which of the following would have been different for the two reactions?
A The pH of the final solution
B The volume of gas produced
C The mass of water formed
D The mass of copper(II) carbonate dissolved
29. Which of the following is the best description of a feedstock?
A A consumer product such as a textile, plastic or detergent.
B A complex chemical that has been synthesised from small molecules.
C A mixture of chemicals formed by the cracking of the naphtha fraction from oil.
D A chemical from which other chemicals can be extracted or synthesised.
30. Which of the following compounds is a raw material in the chemical industry?
A Ammonia
B Calcium carbonate
C Hexane
D Nitric acid
31. The mean bond enthalpy of the N−H bond is equal to one third of the value of ΔH for
which change?
[pic]
32. Ammonia is made by the Haber Process.
N2(g) + 3H2(g) [pic] 2NH3(g)
The equilibrium position lies to the left. Which line in the table is correct?
[pic]
33. The flow chart summarises some industrial processes involving ethene.
[pic]
The feedstocks for ethene in these processes are
A ethane and glycol
B ethane and ethanol
C glycol and poly(ethene)
D glycol, poly(ethene) and ethanol.
34. Polylactic acid is used to make a biodegradable polymer. Polylactic acid can be manufactured by either a batch or a continuous process. What is meant by a batch process? (1)
35. Magnesium metal can be extracted from sea water.
An outline of the reactions involved is shown in the flow diagram.
[pic]
(a) Why can the magnesium hydroxide be easily separated from the calcium chloride at Stage 1? (1)
(b) Name the type of chemical reaction taking place at Stage 2. (1)
(c) Give two different features of this process that make it economical. (2)
36. Cerium metal is extracted from the mineral monazite.
The flow diagram for the extraction of cerium from the mineral is shown below.
[pic]
(a) Name the type of chemical reaction taking place in Step A. (1)
(b) In Step B, cerium hydroxide is heated to form cerium oxide, Ce2O3, and compound Z. Name compound Z. (1)
(c) In Step C, cerium metal is obtained by electrolysis. What feature of the electrolysis can be used to reduce the cost of cerium production? (1)
37. Ozone can be produced in the laboratory by electrical discharge.
3O2(g) → 2O3(g)
Calculate the approximate number of O3(g) molecules produced from one mole of O2(g) molecules. (1)
38. Chlorine gas can be produced by heating calcium hypochlorite, Ca(OCl)2, in dilute
hydrochloric acid.
Ca(OCl)2(s) + 2HCl(aq) → Ca(OH)2(aq) + 2Cl2(g)
Calculate the mass of calcium hypochlorite that would be needed to produce 0·096 litres of chlorine gas. (Take the molar volume of chlorine gas to be 24 litres mol–1.)
Show your working clearly. (2)
39. A student bubbled 240 cm3 of carbon dioxide into 400cm3 of 0.10 mol l–1 lithium hydroxide solution.
The equation for the reaction is:
2LiOH(aq) + CO2(g) → Li2CO3(aq) + H2O(l)
Calculate the number of moles of lithium hydroxide that would not have reacted.
(Take the molar volume of carbon dioxide to be 24 litres mol–1.)
Show your working clearly. (2)
40. (a) In the lab, nitrogen dioxide gas can be prepared by heating copper(II) nitrate.
Cu(NO3)2(s) → CuO(s) + 2NO2(g) + ½ O2(g)
Calculate the volume of nitrogen dioxide gas produced when 2.0g of copper(II) nitrate is completely decomposed on heating. (Take the molar volume of nitrogen dioxide to be 24 litres mol–1.) Show your working clearly. (2)
(b) Nitrogen dioxide has a boiling point of 22 °C. Complete the diagram to show how nitrogen dioxide can be separated and collected. (1)
[pic]
41. Sherbet contains a mixture of sodium hydrogencarbonate and tartaric acid. The
fizzing sensation in the mouth is due to the carbon dioxide produced in the
following reaction.
[pic]
In an experiment, a student found that adding water to 20 sherbet sweets produced
105 cm3 of carbon dioxide. Assuming that sodium hydrogencarbonate is in excess, calculate the average mass of tartaric acid, in grams, in one sweet.
(Take the molar volume of carbon dioxide to be 24 litre mol–1.)
Show your working clearly. (2)
42. The nutritional information states that 100 g of margarine contains 0.70 g of sodium. The sodium is present as sodium chloride (NaCl). Calculate the mass of sodium chloride, in g, present in every 100 g of margarine. (1)
43. Hydrogen fluoride gas is manufactured by reacting calcium fluoride with concentrated sulphuric acid.
CaF2 + H2SO4 ( CaSO4 + 2HF
What volume of hydrogen fluoride gas is produced when 1.0 kg of calcium fluoride reacts completely with concentrated sulphuric acid?
(Take the molar volume of hydrogen fluoride gas to be 24 litres mol–1.)
Show your working clearly. (2)
44. Methanamide, HCONH2, is widely used in industry to make nitrogen compounds.
It is also used as a solvent as it can dissolve ionic compounds.
[pic]
(a) Why is methanamide a suitable solvent for ionic compounds? (1)
(b) In industry, methanamide is produced by the reaction of an ester with ammonia.
[pic]
(i) Name the ester used in the industrial manufacture of methanamide. (1)
(ii) Calculate the atom economy for the production of methanamide. (1)
(c) In the lab, methanamide can be prepared by the reaction of methanoic acid
with ammonia.
[pic]
When 1·38 g of methanoic acid was reacted with excess ammonia, 0·945 g of methanamide was produced. Calculate the percentage yield of methanamide. Show your working clearly. (2)
45. Aspirin, a common pain-killer, can be made by the reaction of salicylic acid with
ethanoic anhydride.
[pic]
(a) Calculate the atom economy for the formation of aspirin using this method.
Show your working clearly. (2)
(b) In a laboratory preparation of aspirin, 5·02 g of salicylic acid produced 2·62 g
of aspirin. Calculate the percentage yield of aspirin. Show your working clearly. (2)
46. From the 1990s, ibuprofen has been synthesised by a three step process. The equation below shows the final step of the synthesis.
[pic]
What is the atom economy of this step? (1)
47. One of the chemicals released in a bee sting is an ester that has the structure shown.
[pic]
This ester can be produced by the reaction of 2-methylbutan-1-ol with ethanoic acid.
If there is a 65% yield, calculate the mass of ester produced, in grams, when 4.0 g of the alcohol reacts with a slight excess of the acid.
(Mass of one mole of the alcohol = 88 g; mass of one mole of the ester = 130 g)
Show your working clearly. (2)
48. Ammonia is produced in industry by the Haber Process.
N2(g) + 3H2(g) [pic] 2NH3(g)
Under certain conditions, 500 kg of nitrogen reacts with excess hydrogen to produce
405 kg of ammonia. Calculate the percentage yield of ammonia under these conditions. Show your working clearly. (2)
49. Ethane-1,2-diol is produced in industry by reacting glycerol with hydrogen.
[pic]
Excess hydrogen reacts with 27·6 kg of glycerol to produce 13·4 kg of ethane-1,2-diol.
Calculate the percentage yield of ethane-1,2-diol. Show your working clearly. (2)
50. Sulphur trioxide can be prepared in the laboratory by the reaction of sulphur dioxide
with oxygen.
2SO2(g) + O2(g) [pic] 2SO3(g)
The sulphur dioxide and oxygen gases are dried by bubbling them through concentrated
sulphuric acid. The reaction mixture is passed over heated vanadium(V) oxide.
Sulphur trioxide has a melting point of 17 °C. It is collected as a white crystalline solid.
(a) Complete the diagram to show how the reactant gases are dried and the product is collected. (2)
[pic]
(b) Under certain conditions, 43·2 tonnes of sulphur trioxide are produced in the reaction of 51·2 tonnes of sulphur dioxide with excess oxygen. Calculate the percentage yield of sulphur trioxide. Show your working clearly. (2)
3.2 Equilibria
1. A catalyst is used in the Haber Process.
N2(g) + 3H2(g) [pic] 2NH3(g)
Which of the following best describes the action of the catalyst?
A Increases the rate of the forward reaction only
B Increases the rate of the reverse reaction only
C Increases the rate of both the forward and reverse reactions
D Changes the position of the equilibrium of the reaction
2. In which of the following systems will the equilibrium be unaffected by a change in
pressure?
A 2NO2(g) [pic]N2O4(g)
B H2(g) + I2(g) [pic] 2HI(g)
C N2(g) + 3H2(g) [pic] 2NH3(g)
D 2NO(g) + O2(g) [pic] 2NO2(g)
3. A few drops of concentrated sulphuric acid were added to a mixture of 0·1 mol of
methanol and 0·2 mol of ethanoic acid. Even after a considerable time, the reaction mixture was found to contain some of each reactant. Which of the following is the best explanation for the incomplete reaction?
A The temperature was too low.
B An equilibrium mixture was formed.
C Insufficient methanol was used.
D Insufficient ethanoic acid was used.
4. Which line in the table shows the effect of a catalyst on the reaction rates and position of equilibrium in a reversible reaction?
[pic]
5. The following equilibrium exists in bromine water.
[pic]
The red colour of bromine water would fade on adding a few drops of a concentrated
solution of
A HCl
B KBr
C AgNO3
D NaOBr.
6. A catalyst is added to a reaction at equilibrium. Which of the following does not apply?
A The rate of the forward reaction increases.
B The rate of the reverse reaction increases.
C The position of equilibrium remains unchanged.
D The position of equilibrium shifts to the right.
7. In which of the following reactions woulD an increase in pressure cause the equilibrium position to move to the left?
[pic]
8. If ammonia is added to a solution containing copper(II) ions an equilibrium is set up.
Cu2+(aq) + 2OH–(aq) + 4NH3(aq) [pic] Cu(NH3)4(OH)2(aq)
(deep blue)
If acid is added to this equilibrium system
A the intensity of the deep blue colour will increase
B the equilibrium position will move to the right
C the concentration of Cu2+(aq) ions will increase
D the equilibrium position will not be affected.
9. Steam and carbon monoxide react to form an equilibrium mixture.
CO(g) + H2O(g) [pic] H2(g) + CO2(g)
Which of the following graphs shows how the rates of the forward and reverse reactions
change when carbon monoxide and steam are mixed?
[pic]
10. 2SO2(g) + O2(g) [pic] 2SO3(g)
The equation represents a mixture at equilibrium. Which line in the table is true for the mixture after a further 2 hours of reaction?
[pic]
11. In which of the following would an increase in pressure result in the equilibrium position
being moved to the left?
[pic]
12. Nitrogen dioxide gas can be prepared in different ways. It is manufactured industrially as part of the Ostwald process. In the first stage of the process, nitrogen monoxide is produced by passing ammonia and oxygen over a platinum catalyst.
NH3(g) + O2(g) ( NO(g) + H2O(g)
(a) Balance the above equation. (1)
(b) Platinum metal is a heterogeneous catalyst for this reaction. What is meant by a heterogeneous catalyst? (1)
(c) The nitrogen monoxide then combines with oxygen in an exothermic reaction to form nitrogen dioxide.
2NO(g) + O2(g) [pic] 2NO2(g)
What happens to the yield of nitrogen dioxide gas if the reaction mixture is cooled? (1)
13. Atmospheric oxygen, O2(g), dissolves in the Earth’s oceans forming dissolved oxygen, O2(aq), which is essential for aquatic life. An equilibrium is established.
O2(g) + (aq) [pic] O2(aq) ΔH = –12·1 kJ mol–1
(a) (i) What is meant by a reaction at “equilibrium”?
(ii) What would happen to the concentration of dissolved oxygen if the temperature of the Earth’s oceans increased?
(b) A sample of oceanic water was found to contain 0·010 g of dissolved oxygen.
Calculate the number of moles of dissolved oxygen present in the sample. (1)
14. When cyclopropane gas is heated over a catalyst, it isomerises to form propene gas and an equilibrium is obtained.
[pic]
The graph shows the concentrations of cyclopropane and propene as equilibrium is established in the reaction.
[pic]
(a) Mark clearly on the graph the point at which equilibrium has just been reached. (1)
(b) Why does increasing the pressure have no effect on the position of this equilibrium? (1)
(c) The equilibrium can also be achieved by starting with propene.
[pic]
Using the initial concentrations shown, sketch a graph to show how the concentrations of propene and cyclopropane change as equilibrium is reached for the reverse reaction (1)
[pic]
15. Tetrafluoroethene, C2F4, is produced in industry by a series of reactions. The final reaction in its manufacture is shown below.
2CHClF2(g) [pic] C2F4(g) + 2HCl(g)
The graph shows the variation in concentration of C2F4 formed as temperature is increased.
[pic]
(a) What conclusion can be drawn about the enthalpy change for the formation of tetrafluoroethene? (1)
(b) Sketch a graph to show how the concentration of tetrafluoroethene formed would vary with increasing pressure. (1)
[pic]
16. Ammonia is produced in industry by the Haber Process.
N2(g) + 3H2(g) [pic] 2NH3(g)
(a) State whether the industrial manufacture of ammonia is likely to be a batch or a continuous process. (1)
(b) The graph shows how the percentage yield of ammonia changes with temperature at a pressure of 100 atmospheres.
[pic]
(i) A student correctly concludes from the graph that the production of ammonia is an exothermic process. What is the reasoning that leads to this conclusion? (1)
(ii) Explain clearly why the industrial manufacture of ammonia is carried out at a pressure greater than 100 atmospheres. (2)
17. Rivers and drains are carefully monitored to ensure that they remain uncontaminated by potentially harmful substances from nearby industries. Chromate ions, CrO42–, are particularly hazardous.
When chromate ions dissolve in water the following equilibrium is established.
2CrO42–(aq) + 2H+(aq) [pic] Cr2O7 2–(aq) + H2O(l)
yellow orange
Explain fully the colour change that would be observed when solid sodium hydroxide is added to the solution. (2)
3.3 Chemical Energy
1. Which of the following represents an exothermic process?
A Cl2(g) → 2Cl(g)
B Na(s) → Na(g)
C Na(g) → Na+(g) + e–
D Na+(g) + Cl–(g) → Na+Cl–(s)
2. Which of the following equations represents an enthalpy of combustion?
[pic]
3. The enthalpy of combustion of methanol is –727 kJ mol–1. What mass of methanol has to be burned to produce 72·7 kJ?
A 3·2 g
B 32·0 g
C 72·7 g
D 727·0 g
4. 5N2O4(l) + 4CH3NHNH2(l) → 4CO2(g) + 12H2O(l) + 9N2(g) ΔH = −5116 kJ
The energy released when 2 moles of each reactant are mixed and ignited is
A 2046 kJ
B 2558 kJ
C 4093 kJ
D 5116 kJ.
5. Aluminium reacts with oxygen to form aluminium oxide.
2Al(s) + 1½O2(g) → Al2O3(s) ΔH = −1670 kJ mol−1
What is the enthalpy of combustion of aluminium in kJ mol−1?
A −835
B −1113
C −1670
D +1670
6.
What is the relationship between a, b, c and d?
A a = b + c – d
B a = d – b – c
C a = b – c – d
D a = d + c – b
7. Given the equations
[pic]
then, according to Hess’s Law
A c = a – b
B c = a + b
C c = b – a
D c = – b – a.
8.
What is the relationship between a, b, c and d?
A a = c + d − b
B a = b − c − d
C a = −b − c − d
D a = c + b + d
9. In the presence of bright light, hydrogen and chlorine react explosively. One step in the reaction is shown below.
H2(g) + Cl(g) → HCl(g) + H(g)
The enthalpy change for this step can be represented as
A (H-H bond enthalpy) + (Cl-Cl bond enthalpy)
B (H-H bond enthalpy) − (Cl-Cl bond enthalpy)
C (H-H bond enthalpy) + (H-Cl bond enthalpy)
D (H-H bond enthalpy) − (H-Cl bond enthalpy).
10. The energy changes taking place during chemical reactions have many everyday uses.
(a) Some portable cold packs make use of the temperature drop that takes place when the chemicals in the pack dissolve in water.
Name the type of reaction that results in a fall in temperature. (1)
(b) Flameless heaters are used by mountain climbers to heat food and drinks. The chemical reaction in a flameless heater releases 45 kJ of energy. If 200 g of water is heated using this heater, calculate the rise in temperature of the water, in °C. (1)
11. The compound diborane (B2H6) is used as a rocket fuel.
(a) It can be prepared as shown.
BF3 + NaBH4 → B2H6 + NaBF4
Balance this equation. (1)
(b) The equation for the combustion of diborane is shown below.
B2H6(g) + 3O2(g) → B2O3(s) + 3H2O(l)
Calculate the enthalpy of combustion of diborane (B2H6) in kJ mol−1, using the following data. (2)
[pic]
12. Hydrogen peroxide decomposes as shown:
H2O2(aq) → H2O(l) + ½ O2(g)
The reaction can be catalysed by iron(III) nitrate solution.
(a) What type of catalyst is iron(III) nitrate solution in this reaction? (1)
(b) In order to calculate the enthalpy change for the decomposition of hydrogen peroxide, a student added iron(III) nitrate solution to hydrogen peroxide solution.
[pic]
As a result of the reaction, the temperature of the solution in the polystyrene beaker
increased by 16 °C.
(i) What is the effect of the catalyst on the enthalpy change (ΔH) for the reaction? (1)
(ii) Use the experimental data to calculate the enthalpy change, in kJ mol–1, for the decomposition of hydrogen peroxide. Show your working clearly. (3)
13. A student used the simple laboratory apparatus shown to determine the enthalpy of combustion of methanol.
[pic]
(a) (i) What measurements are needed to calculate the energy released by the burning methanol? (1)
(ii) The student found that burning 0.370 g of methanol produces 3.86 kJ of energy.
Use this result to calculate the enthalpy of combustion of methanol. (1)
(b) A more accurate value can be obtained using a bomb calorimeter.
[pic]
One reason for the more accurate value is that less heat is lost to the surroundings than in the simple laboratory method. Give one other reason for the value being more accurate in the bomb calorimeter method. (1)
14. Mobile phones are being developed that can be powered by methanol. Methanol can be made by a two-stage process.
In the first stage, methane is reacted with steam to produce a mixture of carbon monoxide and hydrogen.
CH4(g) + H2O(g) [pic] CO(g) + 3H2(g)
Use the data below to calculate the enthalpy change, in kJ mol–1, for the forward reaction.
[pic]
Show your working clearly. (2)
15. The enthalpies of combustion of some alcohols are shown in the table.
[pic]
(a) Using this data, predict the enthalpy of combustion of butan-1-ol, in kJ mol–1. (1)
(b) A value for the enthalpy of combustion of butan-2-ol, C4H9OH, can be determined experimentally using the apparatus shown.
[pic]
Mass of butan-2-ol burned = 1·0 g
Temperature rise of water = 40 °C
Use these results to calculate the enthalpy of combustion of butan-2-ol, in kJ mol–1. (2)
(c) Enthalpy changes can also be calculated using Hess’s Law. The enthalpy of formation for pentan-1-ol is shown below.
5C(s) + 6H2(g) + ½ O2(g) → C5H11OH(l) ΔH = –354 kJ mol–1
Using this value, and the enthalpies of combustion of carbon and hydrogen from the data booklet, calculate the enthalpy of combustion of pentan-1-ol, in kJ mol–1. (2)
16. Different fuels are used for different purposes.
(a) Ethanol, C2H5OH, can be used as a fuel in some camping stoves.
[pic]
(i) The enthalpy of combustion of ethanol given in the data booklet is
−1367 kJ mol−1. Using this value, calculate the mass of ethanol, in g, required to raise the temperature of 500 g of water from 18 °C to 100 °C. Show your working clearly. (3)
(ii) Suggest two reasons why less energy is obtained from burning ethanol in the camping stove than is predicted from its enthalpy of combustion. (2)
(b) Petrol is a fuel used in cars.
[pic]
A car has a 50·0 litre petrol tank. Calculate the energy, in kJ, released by the complete combustion of one tank of petrol. (2)
17. Hydrogen has been named as a ‘fuel for the future’. In a recent article researchers
reported success in making hydrogen from glycerol:
C3H8O3(l) → CO2(g) + CH4(g) + H2(g)
(a) Balance this equation. (1)
(b) The enthalpy of formation of glycerol is the enthalpy change for the reaction:
3C(s) + 4H2(g) + 1½O2(g) → C3H8O3(l)
(graphite)
Calculate the enthalpy of formation of glycerol, in kJ mol–1, using information from the data booklet and the following data.
C3H8O3(l) + 3½O2(g) → 3CO2(g) + 4H2O(l) ΔH = –1654 kJ mol–1
Show your working clearly. (2)
18. The equation for the enthalpy of formation of ethyne is:
2C(s) + H2(g) ( C2H2(g)
Use the enthalpies of combustion of carbon, hydrogen and ethyne given in the data booklet to calculate the enthalpy of formation of ethyne, in kJ mol–1. Show your working clearly. (2)
19. Chloromethane can be produced by the reaction of methane with chlorine.
CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g)
Using bond enthalpies from the data booklet, calculate the enthalpy change, in
kJ mol−1, for this reaction. (2)
20. The production of hydrogen chloride from hydrogen and chlorine is exothermic.
H2(g) + Cl2(g) → 2HCl(g)
Using bond enthalpy values, calculate the enthalpy change, in kJ mol-1, for this reaction. (2)
21. When in danger, bombardier beetles can fire a hot, toxic mixture of chemicals at the
attacker. This mixture contains quinone, C6H4O2, a compound that is formed by the reaction of hydroquinone, C6H4(OH)2, with hydrogen peroxide, H2O2. The reaction is catalysed by an enzyme called catalase.
(a) Most enzymes can catalyse only specific reactions, eg catalase cannot catalyse the hydrolysis of starch. Give a reason for this. (1)
(b) The equation for the overall reaction is:
C6H4(OH)2(aq) + H2O2(aq) ( C6H4O2(aq) + 2H2O(l)
Use the following data to calculate the enthalpy change, in kJ mol–1, for the above
reaction.
C6H4(OH)2(aq) → C6H4O2(aq) + H2(g) ΔH = +177.4 kJ mol–1
H2(g) + O2(g) → H2O2(aq) ΔH = –191.2 kJ mol–1
H2(g) + ½ O2(g) → H2O(g) ΔH = –241.8 kJ mol–1
H2O(g) → H2O(l) ΔH = –43.8 kJ mol–1
Show your working clearly. (2)
22. (a) Methane is produced in the reaction of aluminium carbide with water.
Al4C3 + H2O → Al(OH)3 + CH4
Balance the above equation. (1)
(b) Silane, silicon hydride, is formed in the reaction of silicon with hydrogen.
Si(s) + 2H2(g) → SiH4(g)
silane
The enthalpy change for this reaction is called the enthalpy of formation of silane.
The combustion of silane gives silicon dioxide and water.
SiH4(g) + 2O2(g) → SiO2(s) + 2H2O(l) ΔH = –1517 kJ mol–1
The enthalpy of combustion of silicon is –911 kJ mol–1.
Use this information and the enthalpy of combustion of hydrogen in the data booklet to calculate the enthalpy of formation of silane, in kJ mol–1. Show your working clearly. (2)
23. (a) Hess’s Law can be verified using the reactions summarised below.
[pic]
(i) Complete the list of measurements that would have to be carried out in order to determine the enthalpy change for Reaction 2. (1)
[pic]
(ii) Why was the reaction carried out in a polystyrene cup? (1)
(iii) A student found that 1·08kJ of energy was released when 1·2 g of potassium hydroxide was dissolved completely in water. Calculate the enthalpy of solution of potassium hydroxide. (1)
(b) A student wrote the following incorrect statement.
The enthalpy of neutralisation for hydrochloric acid reacting with potassium hydroxide is less than that for sulphuric acid reacting with potassium hydroxide because fewer moles of water are formed as shown in these equations.
HCl + KOH → KCl + H2O
H2SO4 + 2KOH → K2SO4 + 2H2O
Explain the mistake in the student’s statement. (1)
3.4 Oxidising or Reducing Agents
1. The iodate ion, IO3–, can be converted to iodine.
Which is the correct ion-electron equation for the reaction?
A 2IO3–(aq) + 12H+(aq) + 12e– ( 2I–(aq) + 6H2O(l)
B IO3–(aq) + 6H+(aq) + 7e– ( I–(aq) + 3H2O(l)
C 2IO3–(aq) + 12H+(aq) + 11e– ( I2(aq) + 6H2O(l)
D 2IO3–(aq) + 12H+(aq) + 10e– ( I2(aq) + 6H2O(l)
2. Which of the following is a redox reaction?
A Mg + 2HCl ( MgCl2 + H2
B MgO + 2HCl ( MgCl2 + H2O
C MgCO3 + 2HCl ( MgCl2 + H2O + CO2
D Mg(OH)2 + 2HCl ( MgCl2 + 2H2O
3. The ion-electron equations for a redox reaction are:
[pic]
How many moles of iodide ions are oxidized by one mole of permanganate ions?
A 0.2
B 0.4
C 2
D 5
4. In which of the following reactions is the hydrogen ion acting as an oxidising agent?
[pic]
5. During a redox process in acid solution, iodate ions IO3−(aq) are converted into iodine I2(aq).
IO3−(aq) → I2(aq)
The numbers of H+(aq) and H2O(l) required to balance the ion-electron equation for the
formation of 1 mol of I2(aq) are, respectively
A 3 and 6
B 6 and 3
C 6 and 12
D 12 and 6.
6. In which of the following reactions is a positive ion reduced?
A Iodide ( iodine
B Nickel(II) ( nickel(III)
C Cobalt(III) ( cobalt(II)
D Sulphate ( sulphite
7. In which reaction is hydrogen gas acting as an oxidising agent?
A H2 + CuO → H2O + Cu
B H2 + C2H4 → C2H6
C H2 + Cl2 → 2HCl
D H2 + 2Na → 2NaH
8. Iodide ions can be oxidised using acidified potassium permanganate solution.
The equations are:
[pic]
How many moles of iodide ions are oxidized by one mole of permanganate ions?
A 1.0
B 2.0
C 2.5
D 5.0
9. During a redox process in acid solution, iodate ions are converted into iodine.
2IO3–(aq) + 12H+(aq) + xe– → I2(aq) + 6H2O(l)
To balance the equation, what is the value of x?
A 2
B 6
C 10
D 12
10. The following reactions take place when nitric acid is added to zinc.
[pic]
How many moles of NO3– (aq) are reduced by one mole of zinc?
[pic]
11. One of the reactions taking place within a carbon monoxide sensor is
2CO + 2H2O → 2CO2 + 4H+ + 4e−
This reaction is an example of
A reduction
B redox
C oxidation
D hydration.
12. The concentration of ethanol in a person’s breath can be determined by measuring the voltage produced in an electrochemical cell.
[pic]
Different ethanol vapour concentrations produce different voltages as is shown in
the graph below.
[pic]
(a) Calculate the mass of ethanol, in g, in 1000 cm3 of breath when a voltage of 20 mV was recorded. (Take the molar volume of ethanol, C2H5OH, vapour to be
24 litres mol−1.) Show your working clearly. (3)
(b) The ion-electron equations for the reduction and oxidation reactions occurring in the cell are shown below.
O2 + 4H+ + 4e− → 2H2O
CH3CH2OH + H2O → CH3COOH + 4H+ + 4e−
Write the overall redox equation for the reaction taking place. (1)
13. Oxalic acid is found in rhubarb. The number of moles of oxalic acid in a carton of rhubarb juice can be found by titrating samples of the juice with a solution of potassium permanganate, a powerful oxidising agent.
The equation for the overall reaction is:
5(COOH)2(aq) + 6H+(aq) + 2MnO4 –(aq) → 2Mn2+ (aq) + 10CO2(aq) + 8H2O(l)
(a) Write the ion-electron equation for the reduction reaction. (1)
(b) Why is an indicator not required to detect the end-point of the titration? (1)
3.5 Chemical Analysis
1. An organic chemist is attempting to synthesise a fragrance compound by the following chemical reaction.
compound X + compound Y → fragrance compound
After one hour, a sample is removed and compared with pure samples of compounds X and Y using thin-layer chromatography. Which of the following chromatograms shows that the reaction has produced a pure sample of the fragrance compound?
[pic]
2. The alcohol content of wine was analysed by four students. Each student carried out the experiment three times.
[pic]
The most reproducible results were obtained by
A Student A
B Student B
C Student C
D Student D
3. 45 cm3 of a solution could be most accurately measured out using a
A 50 cm3 beaker
B 50 cm3 burette
C 50 cm3 pipette
D 50 cm3 measuring cylinder.
4. Aluminium carbonate can be produced by the following reaction.
2AlCl3(aq) + 3K2CO3(aq) → Al2(CO3)3(s) + 6KCl(aq)
The most suitable method for obtaining a sample of the aluminium carbonate is
A collection over water
B distillation
C evaporation
D filtration.
5. Seaweeds are a rich source of iodine in the form of iodide ions. The mass of iodine in a seaweed can be found using the procedure outlined below.
(a) Step 1
The seaweed is dried in an oven and ground into a fine powder. Hydrogen peroxide solution is then added to oxidise the iodide ions to iodine molecules.
The ion-electron equation for the reduction reaction is shown.
H2O2(aq) + 2H+(aq) + 2e– → 2H2O(l)
Write a balanced redox equation for the reaction of hydrogen peroxide with iodide ions. (1)
(b) Step 2
Using starch solution as an indicator, the iodine solution is then titrated with sodium thiosulphate solution to find the mass of iodine in the sample. The balanced equation for the reaction is shown.
2Na2S2O3(aq) + I2(aq) → 2NaI(aq) + Na2S4O6(aq)
In an analysis of seaweed, 14.9cm3 of 0.00500 mol l–1 sodium thiosulphate solution was required to reach the end-point. Calculate the mass of iodine present in the seaweed sample. Show your working clearly. (3)
6. Oxalic acid is found in rhubarb. The number of moles of oxalic acid in a carton of rhubarb juice can be found by titrating samples of the juice with a solution of potassium permanganate, a powerful oxidising agent.
The equation for the overall reaction is:
5(COOH)2(aq) + 6H+(aq) + 2MnO4 –(aq) → 2Mn2+ (aq) + 10CO2(aq) + 8H2O(l)
(a) Write the ion-electron equation for the reduction reaction. (1)
(b) Why is an indicator not required to detect the end-point of the titration? (1)
(c) In an investigation using a 500 cm3 carton of rhubarb juice, separate 25.0cm3 samples were measured out. Three samples were then titrated with 0.040 mol l–1 potassium permanganate solution, giving the following results.
[pic]
Average volume of potassium permanganate solution used = 26.9cm3.
(i) Why was the first titration result not included in calculating the average volume of potassium permanganate solution used? (1)
(ii) Calculate the number of moles of oxalic acid in the 500 cm3 carton of rhubarb juice.
Show your working clearly. (2)
7. The number of moles of carbon monoxide in a sample of air can be measured as follows.
Step 1 The carbon monoxide reacts with iodine(V) oxide, producing iodine.
5CO(g) + I2O5(s) → I2(s) + 5CO2(g)
Step 2 The iodine is then dissolved in potassium iodide solution and titrated
against sodium thiosulphate solution.
I2(aq) + 2S2O32–(aq) → S4O62–(aq) + 2I–(aq)
(a) Write the ion-electron equation for the oxidation reaction in Step 2. (1)
(b) Name a chemical that can be used to indicate when all of the iodine has been
removed in the reaction taking place in Step 2. (1)
(c) If 50.4cm3 of 0.10 mol l–1 sodium thiosulphate solution was used in a titration,
calculate the number of moles of carbon monoxide in the sample of air.
Show your working clearly. (2)
8. A major problem for the developed world is the pollution of rivers and streams by nitrite and nitrate ions. The concentration of nitrite ions, NO2 –(aq), in water can be determined by titrating samples against acidified permanganate solution.
(a) Suggest two points of good practice that should be followed to ensure that an accurate end-point is achieved in a titration.
(b) An average of 21·6cm3 of 0·0150 mol l–1 acidified permanganate solution was required to react completely with the nitrite ions in a 25·0 cm3 sample of river water. The equation for the reaction taking place is:
2MnO4– (aq) + 5NO2–(aq) + 6H+(aq) → 2Mn2+(aq) + 5NO3–(aq) + 3H2O(l)
(i) Calculate the nitrite ion concentration, in mol l–1, in the river water. Show your working clearly. (2)
(ii) During the reaction the nitrite ion is oxidised to the nitrate ion. Complete the ion-electron equation for the oxidation of the nitrite ions.
NO2–(aq) → NO3–(aq) (1)
9. (a) The concentration of chromate ions in water can be measured by titrating with a solution of iron(II) sulphate solution. To prepare the iron(II) sulphate solution used in this titration, iron(II) sulphate crystals were weighed accurately into a dry beaker. Describe how these crystals should be dissolved and then transferred to a standard flask in order to produce a solution of accurately known concentration. (2)
(b) A 50·0cm3 sample of contaminated water containing chromate ions was titrated and found to require 27·4 cm3 of 0·0200 mol l–1 iron(II) sulphate solution to reach the end-point. The redox equation for the reaction is:
3Fe2+(aq) + CrO42–(aq) + 8H+(aq) → 3Fe3+(aq) + Cr3+(aq) + 4H2O(l)
Calculate the chromate ion concentration, in mol l–1, present in the sample of water. Show your working clearly. (2)
10. Hydrogen sulfide, H2S, can cause an unpleasant smell in water supplies. The concentration of hydrogen sulfide can be measured by titrating with a chlorine standard solution. The equation for the reaction taking place is
4Cl2(aq) + H2S(aq) + 4H2O(l) → SO42−(aq) + 10H+(aq) + 8Cl−(aq)
50·0 cm3 samples of water were titrated using a 0∙010 mol l−1 chlorine solution.
(a) Name an appropriate piece of apparatus which could be used to measure out the water samples. (1)
(b) What is meant by the term standard solution? (1)
(c) An average of 29·4 cm3 of 0∙010 mol l−1 chlorine solution was required to react completely with a 50·0 cm3 sample of water. Calculate the hydrogen sulfide concentration, in mol l−1, present in the water sample. Show your working clearly. (3)
11. Zinc is an essential element for the body and is found in a variety of foods.
(a) The mass of zinc in four 100 g samples taken from a cheese spread was measured.
[pic]
Calculate the average mass of Zn, in mg, in 100 g of this cheese spread. (1)
(b) The recommended daily allowance of zinc is 9·5 mg for an adult male. 100 g of peanuts contains 3·3 mg of zinc. Calculate the mass of peanuts which would provide the recommended daily allowance of zinc. (2)
12. Solutions containing iodine are used to treat foot rot in sheep. The concentration of iodine in a solution can be determined by titrating with a solution of thiosulfate ions.
[pic]
(a) Write an ion-electron equation for the reaction of the oxidising agent in the titration. (1)
(b) Three 20·0 cm3 samples of a sheep treatment solution were titrated with 0·10 mol l–1 thiosulfate solution. The results are shown below.
[pic]
(i) Why is the volume of sodium thiosulfate used in the calculation taken to be
18·15 cm3, although this is not the average of the three titres in the table? (1)
(ii) Calculate the concentration of iodine, in mol l–1, in the foot rot treatment solution. Show your working clearly. (3)
(iii) Describe how to prepare 250 cm3 of a 0·10 mol l−1 standard solution of sodium thiosulfate, Na2S2O3. Your answer should include the mass, in g, of sodium thiosulfate required. (3)
13. A student carried out an investigation to measure the nitrite level in the school water supply. A compound, which reacts with the nitrite ions to form a product that absorbs light, is added to water samples. The higher the concentration of nitrite ions present in a water sample, the greater the amount of light absorbed.
(a) The student prepared potassium nitrite solutions of known concentration by diluting samples from a stock solution.
(i) Calculate the mass, in mg, of potassium nitrite, KNO2, needed to make 1 litre of stock solution with a nitrite ion concentration of 250 mg l–1. (2)
(ii) Describe how the weighed potassium nitrate is dissolved to prepare the stock solution to ensure that its concentration is accurately known. (2)
(iii) Why should the student use distilled or deionised water rather than tap water when dissolving the potassium nitrite? (1)
(iv) To prepare a solution with a nitrite ion concentration of 0·05 mg l–1 the student dilutes the stock solution. Why is this method more accurate than preparing a solution by weighing out potassium nitrite? (1)
(b) The graph below shows results for five solutions of potassium nitrite and a sample of distilled water.
[pic]
The results for four tap water samples are shown below.
[pic]
What is the concentration of nitrite ions, in mg l–1, in the tap water? (2)
14. Soft drinks contain many ingredients. Caffeine is added to some soft drinks. The concentration of caffeine can be found using chromatography. A chromatogram for a standard solution containing 50 mg l−1 of caffeine is shown below.
[pic]
Results from four caffeine standard solutions were used to produce the calibration graph below.
[pic]
Chromatograms for two soft drinks are shown below.
[pic]
(a) What is the caffeine content, in mg l−1 of soft drink X? (1)
(b) The caffeine content of the soft drink Y cannot be determined from its chromatogram. What should be done to the sample of soft drink Y so that the caffeine content could be reliably calculated? (1)
15. When forensic scientists analyse illegal drugs, anaesthetics such as lidocaine, benzocaine and tetracaine are sometimes found to be present. The gas chromatogram below is from an illegal drug.
[pic]
(a) The structures of lidocaine, benzocaine and tetracaine are shown below.
[pic]
Suggest why benzocaine has a shorter retention time than tetracaine. (1)
(b) Why is it difficult to obtain accurate values for the amount of lidocaine present in a sample containing large amounts of caffeine? (1)
(c) Add a peak to the diagram below to complete the chromatogram for a second sample that only contains half the amount of tetracaine compared to the first. (1)
[pic]
-----------------------
2.5 litres
Pupil booklet
+ H2O
+ HCl
5litres
1 mole carbon dioxide
1 mole of nitrogen
1 mole of oxygen
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+ HCl
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