Second Semester Chemistry Review



Study Guide for Second Semester Final Exam in Chemistry

Students may have a 5x8 index card, hand written, both sides for the final exam

Cumulative Assessment of first and second semester Chemistry

Atomic and Molecular Structure

1. The periodic table displays the elements in increasing atomic number and shows how periodicity of the physical and chemical properties of the elements relates to atomic structure. As a basis for understanding this concept:

a. Students know how to relate the position of an element in the periodic table to its atomic number and atomic mass.

b. Students know how to use the periodic table to identify metals, semimetals, nonmetals, and halogens.

c. Students know how to use the periodic table to identify alkali metals, alkaline earth metals and transition metals, trends in ionization energy, electronegativity, and the relative sizes of ions and atoms.

d. Students know how to use the periodic table to determine the number of electrons available for bonding.

e. Students know how to relate the position of an element in the periodic table to its quantum electron configuration and to its reactivity with other elements in the table.

Chemical Bonds

2. Biological, chemical, and physical properties of matter result from the ability of atoms to form bonds from electrostatic forces between electrons and protons and between atoms and molecules. As a basis for understanding this concept:

a. Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

b. Students know chemical bonds between atoms in molecules such as H2 , CH4 , NH3 , H2 CCH2 , N2 , Cl2 , and many large biological molecules are covalent.

c. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction.

d. Students know how to draw Lewis dot structures.

e. Students know how to identify solids and liquids held together by van der Waals forces or hydrogen bonding and relate these forces to volatility and boiling/ melting point temperatures.

Conservation of Matter and Stoichiometry

3. The conservation of atoms in chemical reactions leads to the principle of conservation of matter and the ability to calculate the mass of products and reactants. As a basis for understanding this concept:

a. Students know how to describe chemical reactions by writing balanced equations.

b. Students know the quantity one mole is set by defining one mole of carbon 12 atoms to have a mass of exactly 12 grams.

c. Students know one mole equals 6.02x1023particles (atoms or molecules).

d. Students know how to determine the molar mass of a molecule from its chemical formula and a table of atomic masses and how to convert the mass of a molecular substance to moles, number of particles, or volume of gas at standard temperature and pressure.

e. Students know how to calculate the masses of reactants and products in a chemical reaction from the mass of one of the reactants or products and the relevant atomic masses.

Gases and Their Properties

4. The kinetic molecular theory describes the motion of atoms and molecules and explains the properties of gases. As a basis for understanding this concept:

a. Students know the random motion of molecules and their collisions with a surface create the observable pressure on that surface.

b. Students know the random motion of molecules explains the diffusion of gases.

c. Students know how to apply the gas laws to relations between the pressure, temperature, and volume of any amount of an ideal gas or any mixture of ideal gases.

d. Students know the values and meanings of standard temperature and pressure (STP).

e. Students know how to convert between the Celsius and Kelvin temperature scales.

f. Students know there is no temperature lower than 0 Kelvin.

g. Students know the kinetic theory of gases relates the absolute temperature of a gas to the average kinetic energy of its molecules or atoms.

h. Students know how to solve problems by using the ideal gas law in the form PV = nRT.

Acids and Bases

5. Acids, bases, and salts are three classes of compounds that form ions in water solutions. As a basis for understanding this concept:

a. Students know the observable properties of acids, bases, and salt solutions.

b. Students know acids are hydrogen-ion-donating and bases are hydrogen-ion-accepting substances.

c. Students know strong acids and bases fully dissociate and weak acids and bases partially dissociate.

d. Students know how to use the pH scale to characterize acid and base solutions.

e. * Students know how to calculate pH from the hydrogen-ion concentration.

Solutions

6. Solutions are homogeneous mixtures of two or more substances. As a basis for understanding this concept:

a. Students know the definitions of solute and solvent.

b. Students know how to describe the dissolving process at the molecular level by using the concept of random molecular motion.

c. Students know temperature, pressure, and surface area affect the dissolving process.

d. Students know how to calculate the concentration of a solute in terms of grams per liter, molarity, parts per million, and percent composition.

e. Students know the relationship between the molality of a solute in a solution and the solution's depressed freezing point or elevated boiling point.

Chemical Thermodynamics

7. Energy is exchanged or transformed in all chemical reactions and physical changes of matter. As a basis for understanding this concept:

a. Students know how to describe temperature and heat flow in terms of the motion of molecules (or atoms).

b. Students know chemical processes can either release (exothermic) or absorb (endothermic) thermal energy.

c. Students know energy is released when a material condenses or freezes and is absorbed when a material evaporates or melts.

d. Students know how to solve problems involving heat flow and temperature changes, using known values of specific heat and latent heat of phase change.

e. Students know how to apply Hess's law to calculate enthalpy change in a reaction.

f. Students know how to use the Gibbs free energy equation to determine whether a reaction would be spontaneous.

Reaction Rates

8. Chemical reaction rates depend on factors that influence the frequency of collision of reactant molecules. As a basis for understanding this concept:

a. Students know the rate of reaction is the decrease in concentration of reactants or the increase in concentration of products with time.

b. Students know how reaction rates depend on such factors as concentration, temperature, and pressure.

c. Students know the role a catalyst plays in increasing the reaction rate.

d. Students know the definition and role of activation energy in a chemical reaction.

Chemical Equilibrium

9. Chemical equilibrium is a dynamic process at the molecular level. As a basis for understanding this concept:

a. Students know how to use Le Chatelier's principle to predict the effect of changes in concentration, temperature, and pressure.

b. Students know equilibrium is established when forward and reverse reaction rates are equal.

c. Students know how to write and calculate an equilibrium constant expression for a reaction.

Nuclear Processes

11. Nuclear processes are those in which an atomic nucleus changes, including radioactive decay of naturally occurring and human-made isotopes, nuclear fission, and nuclear fusion. As a basis for understanding this concept:

a. Students know some naturally occurring isotopes of elements are radioactive, as are isotopes formed in nuclear reactions.

b. Students know the three most common forms of radioactive decay (alpha, beta, and gamma) and know how the nucleus changes in each type of decay.

Other Things to Know

1. Calculate denisty.

2. Significant figures.

3. Mettalic Bonds.

4. VESPR models.

5. Properties of ionic solids.

6. Phase Diagrams.

7. Energy Diagrams.

Chemistry Final Review (from 1st semester)

This review might give you an idea of some areas you may need to study. It is not a complete list. If you discover sections that you’ve struggled with during the school year you should return to your notes, the book, and the numerous worksheets you have completed.

Chapter 1 Introduction to Chemistry

1. scientific method

2. safety

3. equipment

Chapter 2 Data Analysis

1. Significant figures

2. Scientific notation

Chapter 11 The Mole

1. What is the numerical value of Avogadro’s number?

2. What information is provided by the formula for potassium chromate (K2CrO4)?

3. Determine the number of atoms or molecules in each of the following:

a. 0.250 moles of silver

b. 8.56x10-3 moles of sodium chloride

c. 25.8 grams of He gas

d. 1.24 grams of Ca3(PO4)2

4. Convert the following to mass in grams

a. 4.22x10-21 molecules of glucose (C6H12O6)

b. 0.425 moles of iodide gas

5. Express the composition of each as the mass percent of its elements (% composition)

a. sucrose (C12H22O11)

b. magnetite (Fe3O4)

6. What is an empirical formula?

7. Compare a molecular formula to an empirical formula.

8. What is a hydrate?

Chapter 4 Structure of the Atom

1. What flaws exist in Dalton’s, Thomson’s, Rutherford’s and Bohr’s model of the atom?

2. What is the (a) mass number, (b) atomic number of an isotope having 11 protons and 13 neutrons in its nucleus?

3. Describe the evolution of the model of the atom.

Chapter 5 Electrons in Atoms

1. Write electron configurations and energy level diagrams for the following elements: Ca, Zn, Cs, Ca+2, Br-1, S, and Ir.

And vocab……

Chapter 6 & 7 Development of the Modern Periodic Table and the Elements

1. How do chemists use the periodic law to classify elements?

2. What determines the vertical arrangement of the periodic table? Horizontal?

3. What property do the noble gases share?

4. Why is beryllium, a highly reactive metal? Placed in-group 2?

5. Describe the development of the modern periodic table. Include contributions made by Lavoisier, Newlands, Mendeleev, and Moseley.

6. Compare the modern periodic table to Mendeleev

7. Sketch a simplified version of the periodic table and indicate the location of groups, periods, metals, nonmetals, and metalloids.

8. How does the shielding effect alter atomic size?

9. What happens to electron affinity values as you more left to right across a period?

10. Which has the larger radius?

a. Mg or Si c. Ti or Cr

b. Li or Cs d. Mg or Mg+2

11. Which has the larger first ionization energy?

a. Ba or Bi c. Br or Kr

b. Al or Ti d. P or O

12. Which has the more negative value for the first electronegativity

a. B or F b. Cl or S

13. For the element aluminum, the number 26.982 represents the _________________.

a. atomic mass unit and atomic number

b. number of atoms in a mole and the number of grams in a mole

c. mass of an atom in amu and the number of grams in a mole

d. atomic mass and the molarity of its solution

Chapter 8 Ionic Compounds

1. Why do atoms bond?

2. Define, explain, and give examples of the octet rule and valence electrons.

3. What is bond energy?

4. Describe, explain and draw ionic bonding.

5. Briefly explain how malleability and ductility of metals are explained by metallic bonding.

6. Using electron dot structure, diagram the formation of an ionic bond between potassium and iodine.

7. How do you name salts and write formulas? Try these!

Sodium phosphate = ? Ba(OH)2 = ?

Ferric sulfate = ? CrF2 = ?

Calcium iodide = ? Mg3N2 = ?

Silver (I) bromide = ? K2O = ?

Copper (II) chloride = ? CaCl2 = ?

Sodium acetate = ? KNO3 = ?

8. Calculate the molar mass of Pb3(PO4)2.

Second Semester Chemistry Review

Chapter 9 Covalent Bonding

1. Write the formulas for the following compounds

a. Calcium chloride

b. Cobalt (II) carbonate

c. Chromium (III) bromide

d. carbon tetrachloride

e. barium nitrate

f. ammonium hydroxide

2. Write the names for each of these compounds

a. AlF3 d. PCl5

b. Li2CO3 e. P2O5

c. Zn(NO3)2 f. CuSO4

3. What is the difference between a chemical symbol and a chemical formula?

4. Describe the formation of single, double, and triple covalent bonds

5. Compare and contrast sigma and pi bonds.

8. What is the VSEPR model?

9. What are the bond angles in a molecule with a tetrahedral shape?

10. What is hybridization?

11. What elements exist as diatomic molecules?

12. Explain the difference between an empirical formula and a molecular formula.

13. List three differences between molecular and ionic compounds, and explain how they relate to the differences in bond types.

14. Compare the behavior of bonding electrons in a nonpolar covalent bond with that of bonding electrons in a polar covalent bond.

15. Draw Lewis structures for iodine monochloride, ICl, bromine gas, Br2, carbon tetrachloride, CCl4, and hydrogen bromide, HBr. Then identify the shapes.

Chapter 10 Chemical Reactions

1. Balance the following equations: (then label each equation as combustion, synthesis, single-displacement, double displacement, or decomposition.)

a. WO3 + H2 W + H2O

b. PbCl + HNO3 Pb(NO3)2 + HCl

c. RbBr + AgCl AgBr + RbCl

d. HfCl3 + Al HfCl2 + AlCl2

e. Zn + CrCl3 CrCl2 + ZnCl2

f. BaCO3 + C + H2O CO + Ba(OH)2

2. Substitute symbols and formulas for names and write balanced equations for each of the following reactions

a. ammonium nitrite decomposes to nitrogen gas and water

b. sulfuric acid decomposes to water and sulfur trioxide

c. ammonium nitrate decomposes to water and nitrogen dioxide

d. chromium displaces hydrogen to form hydrochloric acid, the chromium (II) chloride as the other product

e. Barium hydroxide reacts with carbon dioxide to from barium carbonate and water

3. Predict the products in the following reaction. Then write and balance the equations.

a. the decomposition of copper (II) oxide

b. react copper metal with silver nitrate to form a copper (II) compound and another product.

c. magnesium metal with oxygen gas

d. calcium hydroxide and sulfuric acid

e. zinc metal and sulfuric acid

4. Write the ionic and net ionic equations for the following. Use your book to identify the ppt, and label the spectators ions.

a. AgNO3(aq) + BaCl2(aq) AgCl(s) + Ba(NO3)2(aq)

b. Mg(s) + HCl(aq) MgCl2(aq) + H2(g)

c. Na2SO4(aq) + BaBr2(aq) NaBr(aq) + BaSO4(s)

5. Name the evidence for a chemical reaction and give examples.

Chapter 12 Stoichiometry

1. Define stoichiometry, limiting reactants, excess reactants and percent yield.

2. What mass of sodium oxide is produced by the reaction of 1.44 g of sodium with oxygen?

3. How much lead (II) nitrate is needed to react with sodium chromate to produce 4.62 kg of lead (II) chromate?

Pb(NO3)2 + Na2CrO4 2NaNO3 + PbCrO4

4. What quantity of hydrogen gas (in grams) is formed when 0.85 g of lithium reacts with water?

2Li + 2H2O 2LiOH + H2

5. What mass (g) of water is given off when 192 g of Cr7H14 burns completely in air?

2Cr7H14 + 21O2 14CO2 + 14H2O

6. How many moles of cesium xenon heptafluoride (CsXeF7) can be produced from the react ion of 12.5 moles of cesium fluoride with 10.0 moles of xenon hexafluoride? Limiting reactant?

CsF(s) + XeF6(s) CsXeF7(s)

7. Iron is obtained commercially by the reaction of hematite (Fe2O3) with carbon monoxide. How many grams of iron are produced if 25.0 moles of hematite react with 30.0 moles of carbon monoxide? Limiting reactant?

Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

8. Upon heating, calcium carbonate decomposes to produce calcium oxide and carbon dioxide.

a. Determine the theoretical yield of CO2 if 235.0 grams of CaCO3 is heated

b. What is the percent yield of CO2 if 97.5 grams CO2 is collected?

Chapter 13 States of Matter

1. Use the kinetic-molecular theory to explain the compression and expansion of gases.

2. Compare diffusion and effusion. Explain the relationship between the rates of these processes and the molar mass of gas.

3. Why are dipole-dipole forces stronger than dispersion forces for molecules of comparable mass?

4. Which of the molecules listed below can form hydrogen bonds? For which of the molecules would dispersion forces be the only intermolecular force? Give reasons for your answers.

a. H2 b. NH3 c. HCl d. HF

5. Explain how hydrogen bonds affect the viscosity of liquids. How does a change in temperature affect viscosity?

6. What effect does soap have on the surface tension of water?

7. How are unit cell and crystal lattice related?

8. What is the difference between a molecular solid and a covalent network solid?

9. What information does a phase change supply?

10. What is the major difference between the processes of melting and freezing?

11. Explain what the triple point and the critical point on a phase diagram represent.

Chapter 14 Gases and Condensation

1. Given the setup in the figure (to right), what would be the pressure of the gas if Ptam were 745 tore and the change in the mercury column is 23.5 mm high?

2. Consider the manometer, on left, first constructed by Robert Boyle. When change in height =40 mm, what is the pressure of the gas trapped container. The temp. Is constant, and atmospheric pressure is Ptam = 766.50 mm Hg.

Answer each of the following with true or false. If statement is false, correct it.

3. One atmosphere equals 760 torr.

4. Boyles' law states that the volume of a gas varies inversely with the pressure under which it is measured.

5. Charles' law states that the volume of a gas varies inversely with the pressure under which it is measured.

6. STP stands for standard temperature, 273 K, and pressure 1 atm.

7. PV=nRT

8. Temp. in C are changed to the Kelvin scale by adding 100: T = t + 100.

9. A mole of ideal gas contains 6.02x1023 molecules and occupies 22.4 liters at STP.

10. In a mixture of gases the total pressure is the sum of the partial pressures of the components.

11. According to Graham's law of diffusion, gases with larger molecular weights diffuse through small openings more rapidly than gases with smaller molecular weights.

12. Below the critical temperature it is impossible to liquefy a gas regardless of pressure.

Complete the following statements with one of the following:

a. increases b. decreases c. remains the same

13. If the temp. of a gas is increased and the pressure on the system is unchanged, the volume of the gas ___________.

14. If the gas is enclosed n a rigid container and the container is heated, the pressure exerted by the gas __________.

15. A gas is allowed to expand from 1 liter to 22.4 liters the number of moles of gas __________.

16. As a gas is heated, the average kinetic energy of the molecules ___________.

17. A weather balloon is released from a station in Texas. As the balloon rises, its size __________ due to decreased atmospheric pressure.

18. A gas in a 1 liter container is heated from 0 C to 100 C. Simultaneously the volume of the container increases the 2 liters. The pressure ____________

19. a vessel containing argon at 0.10 ATM is heated from 0 C to 100 C. If the volume does not change, what is the final pressure?

20. A 75 ml sample of gas is heated at constant pressure from 0 C to 33 C. What is the volume at 33 C?

21. A gas in a 1.00 liter container is allowed to expand to a volume of 5.00 liters. If the initial pressure is 748 torr, what is the final pressure if the temp. does not change?

22. What volume will 16 g of oxygen gas occupy at STP?

23. How many moles of gas at 1.00 ATM and 25 C are contained in a 5.00 liter vessel?

24. A balloon containing 1.00 of He at STP is purchased in an air-conditioned store, in which the temp. is 258 C, and carried outside where it heats to 45 C. If there is no pressure change, what is the final volume?

Chapter 15 Solutions

1. What composition and properties characterize a true solution?

2. Define solute and solvent.

3. What is the most common solvent? Why is it the most common?

4. Describe the process by which water dissolved an ionic compound. How does entropy change in this process?

5. Why is a solution likely to form when a polar solute and polar solvent are combined? Why is a solution unlikely to form when a polar solute and a nonpolar solvent are combined or when a nonpolar solute and polar solvent are combined?

6. How does evaporation differ from boiling?

7. What effect does evaporation from the surface of a liquid have on the average kinetic energy of the molecules remaining in the liquid phase?

8. How can you make water boil without heating it?

9. Predict the solubility of the first substance in the second on the basis of comparative polarities.

a. RbF in ethanol d. NCl3 in C6H6

b. CuS in water e. gasoline in water

c. ethanol in water f. benzene in hexane

10. What is a saturated solution? Describe the equilibrium process that takes place in a saturated solution.

11. Under what conditions can a supersaturated solution form?

12. Discuss three actions that can increase the rate of solution of a solute in a solvent. Explain how each of these actions works to increase the rate of solution.

13. What effect would an increase in pressure have on the solubility of a solid in a liquid? What effect would it have on the solubility of a gas in a liquid?

14. Calculate the molarity of the following solution.

a. 31.1 g of Al2 (SO4) 3 in 1000 ml solution

b. 48.4 g of CaCl2 in 100 ml solution

c. 313.5 g of LiClO3 in 250 ml solution

15. How many moles of Pb(NO3)2 are in 40.0 ml of a 0.250M solution?

16. Explain why vapor pressure of a solvent decreases as a nonvolatile solute is added to the solution.

17. Give the definition of boiling point and freezing point.

18. Why what happens chemically when salt is put on icy roads.

19. Calculate the freezing point and boiling point of a solution that contains 55.4 grams NaCl and 42.3 g KBr dissolved in 750.3 mL of water.

20. The solubility of a certain salt is 40 g/100g of water at 25 C. If a solution is made by adding 360 g to 1 L of water, is this a saturated solution?

21. What effect does adding a solute to a solvent have on the....

.....vapor pressure of the solvent?

......freezing point of the solvent?

.....boiling point of the solvent?

Chapter 16 Energy and Chemical Change

1. When roller coaster cars reach the top of a slope, there seems to be no stopping the exciting downhill plunge. Many chemical reactions proceed this way, too. Energy is added to the reactants, which then react in a way that seems almost unstoppable. How does burning a sheet of newspaper fit the roller-coaster model?

2. What is the minimum starting energy of a reaction called?

3. What is the deference between an exothermic reaction and endothermic reaction?

4. State the law of conservation of energy.

5. What is the relationship between the temperature of an object and the average kinetic energy of its particles?

6. How much heat is required to raise the temperature of 250 g of water from 50 C to 75 C?

7. If 700 g of water at 90 C loses 3500 calories of heat, whit will its final temperature be?

8. A quantity of water is heated from 10 C to 50 C. During the process, 12,000 calories of heat is added to the water. How many grams of water are heated?

9. Potassium bromide has a standard enthalpy of formation of -393.798 kJ/mole. Explain what this statement means.

10. Given a reaction for which H = +2.0 kcal and S = +23.0 kcal/mole at 25 C. Explain by calculating G, how this reaction can occur even though the H is unfavorable.

11. What is H for the reaction Fe2O3(s) + 3CO(g) 2Fe(s) + 3CO2(g)

given:: Hf for Fe2O3(s) is -196.5 kcal/mole

Hf for CO(g) is -26.4 kcal/mole

Hf for CO2(g) is -94.1 kcal/mole

12. What is a primary cause of an increase in the entropy of a system?

13. A chemical reaction tends to proceed spontaneously in the direction of a(n) ______________ in enthalpy and a(n) ______________ in entropy.

14. All spontaneous reactions show a net decrease in __________.

15. When equilibrium exists between 2 opposing reactions, the value of G is __________.

16. State Hess's law. Use and example to explain it.

17. Will the change in entropy for each of the following be positive or negative?

a. dry ice sublimes

b. an ice cube melts

c. raindrops form in a cloud

d. carbohydrates are broken down in cells to release carbon dioxide and water

e. a solution of sodium chloride and a solution of silver nitrate are combined, crystals of

silver chloride settle to the bottom of the container

f. helium under pressure in a cylinder is used to inflate a balloon.

18. In the following situations, tell whether H and TS each have positive or negative values

a. a reaction is exothermic and involves a decrease in order

b. a reaction is endothermic and involves a decrease in order

c. a reaction is exothermic and involves a increase in order

d. a reaction is endothermic and involves a increase in order

In which of these situations is it unclear whether the reaction will occur spontaneously? How can it be determined in these cases whether a reaction will be spontaneous?

19. Relate Gibbs free energy in terms of entropy, enthalpy, and temperature. Write the equation, and then explain it in words.

Chapter 17 Reaction Rates (Kinetics)

1. How is rate usually defined?

2. How does an activated complex form?

3. In terms of activation energy, account for the kinetic stability of a substance.

4. Explain how concentration of reactions affects reaction rate.

5. How does temperature affect the rate at which activated complexes are formed? Describe the two main mechanisms that affect this rate.

6. What does a catalyst do? In general. How does a catalyst carry out its function?

7. A certain reaction takes place in three steps as follows.

a. which is the rate determining step?

b. if you were to examine the contents of the reaction vessel after the reaction had begun but before completion, what relative amounts of each substance would you expect to find?

8. Enzymes in living organisms are examples of _____

9. What is the reaction mechanism? Explain why increasing the concentration of a reactant doesn't always increase the rate of the reaction.

10. Identify all the letters in the following graph, and tell whether the reaction is exo, or endothermic. Describe how a catalyst would change the graph

Chapter 18 Chemical Equilibrium

1. How does the concept of reversibility explain the establishment of equilibrium?

2. Define a reaction that is at equilibrium.

3. Describe Le Chatelier's principle.

4. How does a system at equilibrium respond to a stress? What factors are considered to be stresses on an equilibrium system?

5. Use Le Chatelier’s principle to predict how each of these changes would affect the ammonia equilibrium system

NH2(g) + 3H2(g) ↔ 2NH3(g)

a. removing hydrogen from the system

b. adding ammonia to the system

c. adding hydrogen to the system

Chapter 19 Acids and Bases

1. List some properties of acids and bases.

2. Fundamentally, what determines the strength of an acid or base?

3. _______________ occurs when water causes the molecules of a covalent substance to form ions.

4. An acid reacts with active metals to produce ___________

5. Write the general equation for a neutralization reaction

6. Name these acids

a. HF d. HNO3

b. HCl f. H3PO4

c. H2SO4

7. What is the pH of

a. 0.0000040 M HCl c. 0.000072 KOH

b. 0.0003 M H2SO4 d. 0.00083 Ca(OH)2

8. Determine the hydrogen ion concentration in a solution whose hydroxide ion concentration is 4.49x10-4 M. What is the pH of the solution?

9. The lower the pH, the ________ acidic a substance is.

10. A solution whose pH is 8 has a hydronium ion concentration of _______.

11. What is a titration? Why is a standard solution necessary for a titration?

12. What volume of 0.196 M LiOH is required to neutralize 25.0 ml of 0.413 M HBr?

13. Find the [H+] concentration of the following solutions

a. pH = 3.20 d. pOH = 2.52

b. pH = 10.2 e. pOH = 7.25

c. pH = 6.61 f. pOH = 4.97

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Avogadro’s number

Empirical formula

Hydrate

Molar mass

Mole

Molecular formula

Percent composition

Mass Number

Neutron

Nucleus

Proton

Transmutation equation

Atomic Mass

Atomic Mass Unit

Atomic Number

Dalton’s Atomic Theory

Isotope

Actinides

Alkali metals

Alkaline earth metals

Group

Halogen

Inner Transition Metal

Ion

ionization energy

Metal

Metalloid

Noble Gas

Nonmetal

Octet Rule

Period

Periodic Law

Representative Element

Transition Element

Transition Metal

Coordinate covalent bond Molecule

Covalent bond Pi bond

Endothermic Polar covlent

Exothermic Resonance

Hybridization Sigma bond

Lewis structure Structural formula

VSEPR model

Aqueous solution

Chemical equation

Chemical reaction

Coefficient

Combustion reaction

Ionic equation

Decomposition reaction

Double replacement reaction

Net ionic equation

Precipitate

Product

Reactant

Single replacement reaction

Solute

Solvent

Spectator ion

Synthesis reaction

Avogadro’s number

Empirical formula

Hydrate

Molar mass

Mole

Molecular formula

Percent composition

amorphous solid

barometer

boiling point

condensation

crystalline solid

Dalton’s law of partial pressures

diffusion

dipole-dipole forces

dispersion forces

evaporation

freezing point

Graham’s law of effusion

hydrogen bond

kinetic-molecular theory

melting point

phase diagram

pressure

sublimation

surface tension

surfactant

temperature

triple point

unit cell

vaporization

vapor pressure

viscosity

molar volume

pascal

partial pressure

STP

Diffusion

Effusion equilibrium

Graham’s law of effusion

Ideal gas law

Absolute temperature

Avogadro’s principle

Boyle’s law

Charles’s law

[pic]

Boiling point elevation

Brownian motion

Colligative property

colloid

concentration

freezing point depression

heat of solution

henry’s law

immiscible

insoluble

miscible

molality

molarity

mole fraction

osmotic pressure

saturated solution

solubility

soluble

salvation

supersaturated solution

suspension

tyndall effect

unsaturated solution

vapor pressure lowering

calorie

calorimeter

chemical potential energy

energy

enthalpy

enthalpy (heat) of reaction, formation, or combustion

entropy

free energy

heat

Hess’s Law

Joule

Law of conservation of energy

Law of disorder

Molar enthalpy (heat) of fusion

Molar enthalpy (heat) of vaporization

Specific heat

Spontaneous process

Themochemical equation

thermochemistry

Activated complex

Activation energy

Catalyst

Collision theory

Inhibitor

Instantaneous rate

Intermediate

Rate-determining step

Reaction mechanism

Reaction rate

X(Y; Y(Z; Z(final product

Fast slow fast

C

A

B

D

E

Chemical Equilibrium

Reversible Reaction

Le Chatelier’s principle

Acid-base indicator

Acidic solution

Basic solution

End point

Equivalence point

Neutralization reaction

pH

pOH

Titration

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