Chemistry Unit 6- Essential Knowledge



Chemistry Unit 6- Essential Knowledge

Unit 6 Problem Sets

(Answers for all problems, additional practice worksheets and a practice test are all eventually posted on Blackboard)

Instructions: Do Problems as Assigned. They must be done on a separate sheet of paper in PENCIL. Your full name (first and last), the date, and your class period must be in the upper right corner of your paper. Answers to assigned problems must be legibly printed (no cursive). Typing is only permitted with special permission.

SHOWING YOUR WORK IS REQUIRED FOR ALL MATH PROBLEMS!!!!!!!!

At the end of every problem set you must write in ink the honor pledge: “On my honor as a W-L student, I have neither given nor received aid on this assignment.” And sign your name.

Failure to follow this procedure will result in your paper being returned to you with no credit given.

|Topic |Essential Knowledge |

|Elements and the Periodic |Electronegativity is the ability of an atom in a bond to attract electrons. The electronegativity of an element can be |

|Table |judged from its position on the periodic table. |

| |The electronegativity increases across a period of the periodic table. (The atomic radius decreases, which means that the |

| |valence electrons are closer to the nucleus and held more tightly by the nucleus.) |

|1.6 |The electronegativity decreases down a group of the periodic table. (The valence electrons are further away from and more |

| |loosely held by the nucleus. It is more difficult for the nucleus to attract electrons.) |

| | |

| |Problem Set 1.6 |

| |1. Using the words “bond”, “electrons” and “attract, to DEFINE electronegativity. |

| |2. EXPLAIN how electronegativity and atomic radius are related. |

| |3. Which group/family on the periodic table has no electronegativity because their valence shell (octet) is filled? |

| |_________________ |

| |4. Which is the most electronegative element? |

| |F or Cl F or Li Cl or Si |

| |O or C |

| |Li or K Na or Mg |

|Compounds and Bonding |A covalent bond consists of electrons shared between atoms. This sharing is not always equal, because different atoms have |

|2.6 |different electronegativities. Unequal sharing results in a polar bond. |

| |The more electronegative atom in a covalent bond will attract the electrons more strongly and this will result in it having a|

| |slight negative charge. The less electronegative atom will therefore be slightly deficient in electrons and so will have a |

| |slight positive charge. |

| |A covalent bond in which the atoms do not share their electrons equally and have slight electrical charges is known as a |

| |polar covalent bond. (example: HF) |

| |A covalent bond in which the atoms share their electrons equally and do not have slight electrical charges is known as a |

| |nonpolar covalent bond (example: F2). Only two of the same atom can share electrons equally. |

| | |

| |Problem Set 2.6 |

| |1. Match each of the the models of electron clouds below with one of these three types of bonding: ionic, polar covalent or |

| |non-polar covalent. |

| |[pic] |

| |2. Label each as ionic, polar covalent or nonpolar covalent |

| |a. NaCl _____________ b. N2 ___________________c. HCl ____________d. O2 ___________________ |

|Kinetic Theory |The Heat of Reaction is the net amount of energy released or absorbed (∆ H) in a chemical reaction can be calculated from the|

| |balanced equation. The value for ∆ H is given in kilojoules (kJ). |

|3.6 |(H = HProd – HReact |

| |Exothermic reactions, (H (–) and the HReact > HProd |

| |Endothermic reactions, (H (+) and the HProd > HReact |

| |Temperature is the measure of the average kinetic energy (movement) of the gas particles. This must be expressed in Kelvin |

| |(K) for all Gas Law calculations. |

| |273K = 0o C (K=oC + 273) |

| | |

| |Problem Set 3.6 |

| |1. 0oC = ___________K |

| |2. 25oC = __________K |

| |3. ______oC = 373 K |

| |4. What is another name for 0 K? ___________________________________ |

| |5. If the energy of the products is higher than the energy of the reactants then the reaction is ______________ and ΔH is |

| |________________. |

| |6. If the energy of the products is lower than the energy of the reactants then the reaction is ______________ and ΔH is |

| |________________. |

| |7. For the reaction, ___Fe + ___O2 ( ___Fe2O3 + 1644kJ |

| |Balance the equation. |

| |Is this reaction endothermic or exothermic? ________________ Why? |

| |How much heat would be produced if only 1 mole of iron was burned? |

| |How much heat would be produced if only 1 mole of oxygen gas were consumed? |

| |If 5.6g of iron were burned, how many moles of iron reacted? |

| |How much heat would be produced from the 5.6g reacted in e? |

| |How much heat would be produced if 1.50g of iron III oxide, Fe2O3 were produced? |

|The Mole and Stoichiometry |The molar coefficients from balanced chemical equations are used to predict the masses of reactants or products. |

| |At STP (standard temperature and pressure, 0o Celsius and 1 atmosphere) the volume of 1 mole of any gas is 22.4 liters. If |

|4.6 |the temperature and/or pressure are changed, then the volume will also change. For chemical reactions at STP, the molar |

| |volume of a gas (1 mole = 22.4L) can be used to predict the volume of gas produced. |

| |The theoretical yield of a reaction is the maximum amount of product that can be made if the reaction is 100% efficient and |

| |there are no errors. |

| |%yield = 100 x actual mass of product (given)/theoretical mass of product (calculated) |

| | |

| |Problem Set 4.6 |

| |1. For the reaction, ___Fe2O3 ( ___Fe + ___O2, |

| |Assume that 20.0 g of iron oxide, Fe2O3,is decomposed. |

| |Balance the reaction. |

| |Calculate the moles of iron oxide, Fe2O3, decomposed. |

| |Determine the theoretical moles of iron that would be produced. |

| |What theoretical mass of iron is produced? |

| |When a class of students did this lab, they found that the average amount of Fe produced was 12.0g. What was their |

| |theoretical yield of iron? _______ |

| |2. What is STP ? __________________________________________ |

| |3. 1 mole of a gas at STP = ____________liters |

| |4. 33.6 liters of a gas at STP = ___________ moles |

| |5. 3 moles of a gas at STP = _____________liters |

| |6. 0.02 moles of a gas at STP = ___________ liters |

| |7. 3 moles of He(g) at STP = _____________liters |

| |8. 11.2 liters of Ne(g) at STP = ___________ moles |

| |9. Balance this equation and use it for a. and b. |

| |N2 + H2 ( NH3 (unbalanced) |

| |a. How many liters of ammonia gas (NH3) could be formed from 5.6 grams of N2 ? |

| |b. How many grams of H2 would be needed to make 44.8 liters of ammonia at STP? |

| |In the following problems, calculate how much of the indicated product is made. Show all your work. |

| |10. LiOH + HBr ( LiBr + H2O |

| |If you start with ten grams of lithium hydroxide, how many grams of lithium bromide will be produced? |

| |11. C2H4 + 3 O2 ( 2 CO2 + 2 H2O |

| |If you start with 45 grams of ethylene (C2H4), how many grams of carbon dioxide will be produced? |

| |12. 2 HCl + Na2SO4 ( 2 NaCl + H2SO4 |

| |If you start with 20 grams of hydrochloric acid, how many grams of sulfuric acid will be produced? |

|Chemical Reactions |In a single replacement reaction one element takes the place of another in a compound. The general form for a single |

| |replacement reaction is |

|5.6 |A + BX ( AX + B |

| |Where A and B are elements and BX and AX are ionic compounds. |

| |An example is: Cu + FeCl2 ( CuCl2 + Fe. |

| |In a double replacement reaction the positive portions of two ionic compounds are interchanged. The general form for a double|

| |replacement reaction is |

| |AX + CY ( AY + CX |

| |Where the cations A and C in two ionic compounds that swap partners. |

| |An example is: BaCl2 + Na2O ( BaO + 2NaCl |

| |Problem Set 5.6 |

| |1. Classify each of the following as a single replacement or a double replacement reaction. |

| |a. MX + Z ( MZ + X |

| |b.MX + YZ ( MZ + YX |

| |c.2KI + CaO ( K2O + CaI2 |

| |d.BaS + MgO ( BaO + MgS |

| |e.H2SO4 + FeS ( H2S + FeSO4 |

| |f.2NH4Cl + F2 ( Cl2 + 2NH4F |

| |2. Give the formula for the following compounds. |

| |a. Barium sulfate b. Ammonium Phosphate |

| |c. Sodium Phosphide d. Dinitrogen Pentoxide |

| |e. Nitrogen dioxide f. Calcium Carbonate |

| |g. Lead (III) oxide h. Copper (II) nitrate |

| |3. Name the following compounds |

| |a. Fe SO4 b. NH4Cl c. K2O d. CH4 |

| |e. NH3 f. H2O |

| |4. Predict the products of the following double replacements reactions |

| |a. ____Ca(OH)2 + ____HF ( |

| |b. ____Pb(NO3)2 + ____ K2CrO4 ( |

| |c. ____NaC2H3O2 + ____ H2SO4 ( |

| |d. ____Cu(OH)2 + ____H3PO4 ( |

| |e. ____AgNO3 + ____ Na2CO3 |

|Solutions |The general rule for predicting solubility is “like dissolves like”. Water is a polar substance, so it can dissolve ionic and|

| |polar solutes. |

|6.6 |Oil is non-polar, so oil will not dissolve in water. Oil and water do not mix but different oils will mix with each other |

| |because all oils are nonpolar. Nonpolar solutes will dissolve in nonpolar solvent. |

| |Ammonia (NH3) dissolves in water, therefore it must be a polar molecule. |

| |Temperature can have a great effect on solubility. Solubility graphs relate solubility to temperature. For most solid |

| |solutes, as temperature increases, the solubility increases. |

| |For most gas solutes, solubility decreases as temperature increases. |

| | |

| |Problem Set 6.6 |

| |Which of the following will result in solution formation? |

| |Nonpolar covalent with nonpolar covalent. |

| |Nonpolar covalent with polar covalent. |

| |Polar covalent with polar covalent. |

| |Nonpolar covalent with ionic. |

| |Polar covalent with ionic. |

| |Indicate whether each of the following pairs would form a solution when mixed together. |

| |CH4 and H2O b. H2S and H2O |

| |c. NaCl and C4H10 d. CH4 and C4H10 ________________ |

| |e. KI and CH3OH |

| |Air is mainly composed of three gases, CO2, O2 and N2. Explain why air is not very soluble in water. |

|Experimentation |Mixtures can be separated based on the physical properties of the components of the mixture. |

|7.6 |Centrifugation (using a centrifuge) separates mixtures of undissolved solids from liquids by density. |

| |Distillation separates mixtures of liquids by boiling point |

| |Filtration separates mixtures of undissolved solids from liquids by particle size. |

| |Chromatography separates mixtures based on attractive forces between particles. |

| | |

| |Problem Set 7.6 |

| |Differences in what physical property/properties make each of the following separations possible? |

| |Distillation |

| |Centrifugation |

| |Filtration |

| |Chromatography |

Mastery Knowledge – Unit 6

1. The forces of attraction between neighboring molecules are called intermolecular forces. These attractions are much weaker than either an ionic or a covalent bond.

2. According to the kinetic-molecular theory, the state of a substance at room temperature depends on the strength of the attractions between its particles. Substances with very weak intermolecular attractions are gases at room temperature, while those with moderate attractions are liquid, and those with stronger attractions are solids at room temperature.

3. The three types of intermolecular forces, in order of increasing strength, are van der Waals, dipole-dipole, and hydrogen bonding. Molecules that have stronger intermolecular attractions will have higher boiling points.

4. Electrons in an atom can shift so as to concentrate on one side of the nucleus causing one end of the atom to be relatively more positive than the opposite end, which becomes relatively more negative. This brief shift of electrons is a temporary dipole. When one atom with a temporary dipole causes the formation of a temporary dipole on a neighboring atom it is called an induced dipole. This is the premise behind the Van der Waals Forces.

5. Van der Waals forces (also known as dispersion forces) are due to attractions caused when temporary dipoles are induced throughout a substance. Van der Waals forces are present in all molecular substances. The larger the electron cloud of an atom, the more easily it is distorted into a dipole. Thus, larger atoms have stronger van der Waals forces. Because the seven diatomic molecules (H2, N2, O2, and group 17, the halogens) are nonpolar molecules their only intermolecular attraction is due to van der Waals forces. As you move down Group 17, the physical state at room temperature changes from gaseous fluorine (F2) and chlorine (Cl2), to liquid bromine (Br2), to solid iodine (I2).

6. A permanent dipole is a result of a covalent bond between two different atoms. Because the two atoms are not identical, there is a difference in their electronegativities. As a result, each atom acquires a partial charge, and is called a polar molecule. Dipole-dipole forces occur when polar molecules are attracted to one another.

7. A hydrogen bond is not a bond but an extremely strong dipole-dipole attraction that occurs in molecules that contain hydrogen bonded to nitrogen, oxygen or fluorine. This is because these (N, O, and F) are the three most electronegative nonmetals, whereas hydrogen is one of the least electronegative nonmetals. Because of the large difference in electronegativities in the F-H, O-H, and N-H bonds, these bonds are very polar, giving rise to large partial positive,(+ and partial negative, (-charges. A hydrogen bond is considerably stronger than all other intermolecular forces.

Mastery Knowledge Problem Set

Intermolecular Forces (forces that hold polar and non polar covalent molecules together so that liquids and solids can form)

1. Weak Intermolecular forces are found in the ________ phase of matter.

2. Moderate Intermolecular Forces are found in the ________phase of matter.

3. Strong Intermolecular Forces are found in the _________ phase of matter.

4. Molecules with strong intermolecular forces will have ____________ boiling points.

5. Three types of intermolecular forces are :

6. Define Temporary Dipole:

7. Define Induced Dipole:

8. Define Permanent Dipole:

9. What are Van der Waals forces (known as London Disperson Forces)?

10. What are Dipole-Dipole Forces?

11. What are Hydrogen “Bonds”: (why is “bonds” in quotation marks)?

12. What type of intermolecular forces act on non polar covalent molecules?

13. What type of intermolecular forces act between polar molecules that do not contain hydrogen? _________________and ___________________________

14. What type of intermolecular forces act between polar molecules that contain Hydrogen bonded to F, N and O? ______________________

|NF3 |NaCl |CH4 |HF |NO3- |

|Name |Name |Name |Name |Name |

|Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |

|If covalent, polar or |If covalent, polar or |If covalent, polar or nonpolar |If covalent, polar or nonpolar |If covalent, polar or |

|nonpolar bonds? |nonpolar bonds? |bonds? |bonds? |nonpolar bonds? |

|Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |

|N2 |SiBr4 |PO43- |MgO |H2O |

|Name |Name: |Name |Name |Name |

|Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |

|If covalent, polar or |If covalent, polar or |If covalent, polar or nonpolar |If covalent, polar or nonpolar |If covalent, polar or |

|nonpolar bonds? |nonpolar bonds? |bonds? |bonds? |nonpolar bonds? |

|Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |

|C2H4 |C3H8 |BF3 (exception be careful!) |I2 |CCl2H2 |

|Ionic or Covalent? |Ionic or Covalent? |Ionic or Covalent? |Name |Ionic or Covalent? |

|If covalent, polar or |If covalent, polar or |If covalent, polar or nonpolar |Ionic or Covalent? |If covalent, polar or |

|nonpolar bonds? |nonpolar bonds? |bonds? |If covalent, polar or nonpolar |nonpolar bonds? |

|Lewis Dot Structure |Lewis Dot Structure |Lewis Dot Structure |bonds? |Lewis Dot Structure |

| | | |Lewis Dot Structure | |

Review for Unit 6

1. Define Electronegativity

2. Describe the trends in electronegativity seen on the periodic table

3. If you were an atom with a large electronegativity value, how would you act around another atom?

4. Fill in the table below:

|Type of Bond |Electronegativity Difference |What happens with electrons |Example |

|1. | | | |

|2. | | | |

|3. | | | |

5. What is the following picture of?

6. What type of bonds are in the molecule to the right?

7. Define (H? What are the units of (H?

8. What type of reaction is the following equation and where would you place heat in this reaction?

[pic]

9. Use the following equation to answer the following questions.

2NH3 ( N2 + 3H2

How many moles of nitrogen gas would you get if you started with 36.75 moles of ammonia?

How many moles of hydrogen gas would you get if you started with 1 mole of ammonia?

How much ammonia do you need to make 42.88 moles of hydrogen?

10. Define and give an example of a single replacement reaction.

11. Define and give an example of a double replacement reaction.

12. Define and give an example of a synthesis reaction.

13. Define and give an example of a decomposition reaction.

14. Define and give an example of a combustion reaction.

15. Why don’t oil and water mix?

16. Fill out the solubility table below by placing an X if the substance will dissolve.

| |Polar |Non Polar |Alcohol |

|Polar | | | |

|Non Polar | | | |

|Alcohols | | | |

|Ionic compounds | | | |

17. List and describe 4 separation techniques.

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