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Name: _____________________________________________________________________________________ Date: _______________ Hour: _________

Chemistry Final Exam Review: 2017 - RAEB

Many concepts and skills from first semester will be needed to be successful on the final exam. In addition to this

semester’s material you should also review key 1st semester topics. Common topics from 1st semester that you

should review include: naming compounds and writing formulas, dimensional analysis, and significant figure rules.

This review is worth 5% extra credit to your final exam. It is due the day you are scheduled to take your final.

Your work must be hand written and it must be your own work. Copying answers from another student will not

help you to be successful on this exam. Be sure to answer every question. For calculation questions, you must

show your work to receive credit.

Chapter 11:

1) What are the different types of reactions?

2) How do you determine if a precipitate has formed in a reaction?

3) What is a catalyst?

4) Rewrite the word equation as balanced chemical equations. a. hydrogen + sulfur → hydrogen sulfide

b. iron(III) chloride + calcium hydroxide → iron(III) hydroxide + calcium chloride

5) Balance the following equations.

a. SO2 + O2 → SO3

b. Fe2O3 + H2 → Fe + H2O

c. P + O2 → P4O10

d. Al + N2 → AlN 6) Complete and balance this equation for a combination reaction.

Be + O2 →

7) Complete and balance this decomposition reaction.

HI →

8) Complete the equations for these single replacement reactions in aqueous solution. Balance each equation.

a. Fe(s) + Pb(NO3)2(aq) →

b. Cl2(aq) + NaI(aq) →

c. Ca(s) + H2O(l) →

9) Write the products of these double replacement reactions. Then balance each equation.

a. NaOH(aq) + Fe(NO3)3(aq) →

Note: Iron(III) Hydroxide is a precipitate.

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b. Ba(NO3)2(aq) + H3PO4(aq) →

Note: Barium phosphate is a precipitate.

10) What are the five types of chemical reactions?

11) Classify each reaction and balance the equation.

a. C3H6 + O2 → CO2 + H2O

b. Al(OH)3 → Al2O3 + H2O

c. Li + O2 → Li2O

d. Zn + AgNO3 → Ag + Zn(NO3)2

12) Write the balanced net ionic equation for each reaction and identify the spectator ions in each reaction.

a. Pb(NO3)2(aq) + H2SO4(aq) → PbSO4(s) + HNO3(aq)

b. Pb(C2H3O2)2(aq) + HCl(aq) → PbCl2(s) + HC2H3O2(aq)

c. Na3PO4(aq) + FeCl3(aq) → NaCl(aq) + FePO4(s)

d. (NH4)2S(aq) + Co(NO3)2(aq) → CoS(s) + NH4NO3(aq)

13) Identify the reactants and products in each chemical reaction. a. Hydrogen gas and sodium hydroxide are formed when sodium is dropped into water.

b. In photosynthesis, carbon dioxide and water reacted to form oxygen gas and glucose.

14) What is the purpose of a catalyst?

Chapter 12:

1) What is the difference between limiting and excess reagent?

2) Interpret the equation for the formation of water from its elements in terms of number of molecules and moles, and volumes of gases at STP.

2H2(g) + O2(g) → 2H2O(g)

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3) Chemical reactions can be described in terms of what quantities?

4) What quantities are always conserved in chemical reactions?

5) Interpret the given equation in terms of relative numbers of representative particles, numbers of moles,

and masses of reactants and products. 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)

6) According to the following equation: 4Al(s) + 3O2(g) → 2Al2O3(s). a. How many moles of oxygen are required to react completely with 14.8 mol Al?

b. How many moles of Al2O3 are formed when 0.78 mol O2 reacts with aluminum?

7) Acetylene gas (C2H2) is produced by adding water to calcium carbide (CaC2). CaC2(s) + 2H2O(l) → C2H2(g) + Ca(OH)2(aq)

a. How many grams of acetylene are produced by adding water to 5.00 g CaC2?

b. Determine how many moles of CaC2 are needed to react with 49.0 g H2O.

8) How many molecules of oxygen are produced by the decomposition of 6.54 g of potassium chlorate (KClO3)?

KClO3(s) → 2KCl(s) + 3O2(g)

9) The equation for the combustion of carbon monoxide is: 2CO(g) + O2(g) → 2CO2(g). How many liters of oxygen are required to burn 3.86 L of carbon monoxide?

10) Consider the following equation: CS2(l) + 3O2(g) → CO2(g) + 2SO2(g). Calculate the volume of sulfur dioxide produced when 27.9 mL O2 reacts with carbon disulfide.

11) The equation for the complete combustion of ethane (C2H4) is: C2H4(g) + 3O2(g) → 2CO2(g) + 2H2O(g). If 27.0 mol C2H4 is reacted with 6.30 mol O2, identify the limiting reagent.

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12) The following equation shows the incomplete combustion of ethane: C2H4(g) + 3O2(g) → 2CO(g) + 2H2O(g). If 2.70 mol C2H4 is reated with 6.30 mol O2.

a. Identify the limiting reagent. b. Calculate the moles of water produced.

13) When 84.8 g of iron (III) oxide reacts with an excess of carbon monoxide, iron is produced. Fe2O3(s) + 3CO(g) → 2Fe(s) + 3CO2(g)

What is the theoretical yield of iron?

14) If 50.0 g of silicon dioxide is heated with an excess of carbon, 27.9 g of silicon carbide is produced. SiO2(s) + 3C(s) → SiC(s) + 2CO(g)

What is the percent yield of this reaction?

15) In a chemical reaction, how does an insufficient quantity of a reactant affect the amount of product formed?

16) The reaction of fluorine with ammonia produces dinitrogen tetrafluoride and hydrogen fluoride. 5F2(g) + 2NH3(g) → N2F4(g) + 6HF(g)

a. If you have 66.6 g NH3, how many grams of F2 are required for complete reaction?

b. How many grams of NH3 are required to produce 4.65 g HF

c. How many grams of N2F4 can be produced from 225 g F2?

17) Lithium nitride reacts with water to form ammonia and aqueous lithium hydroxide. Li3N(s) + 3H2O(l) → NH3(g) + 3LiOH(aq)

a. What mass of water is needed to react with 32.9 g Li3N?

b. When the above reaction takes place, how many molecules of NH3 are produced?

c. Calculate the number of grams of Li3N that must be added to an excess of water to produce 15.0 L NH3 (at STP).

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18) What is the significance of the limiting reagent in a reaction? What happens to the amount of any reagent

that is present in an excess?

19) In a reaction chamber, 3.0 mol of aluminum is mixed with 5.3 mol Cl2 and reacts. The reaction is described by the following balanced chemical equation: 2Al + 3Cl2 → 2AlCl3.

a. Identify the limiting reagent for the reaction. b. Calculate the number of moles of product formed.

c. Calculate the number of moles of excess reagent remaining after the reaction. Chapter 13:

1. What pressure (in kilopascals) does a gas exert when at 352 mm of Hg?

2. What pressure (in atmospheres) does a gas exert when at 352 mm of Hg?

3. Describe the basic assumption of the kinetic theory of gases.

4. In terms of kinetic energy, explain how a molecule in a liquid evaporates.

5. With regards to vapor pressure, what must occur for a liquid to boil?

6. What two phases are in equilibrium at a substance’s melting point?

7. What two phases are in equilibrium at a substance’s boiling point?

8. How do the melting points of ionic solids and molecular solids differ? Explain your answer.

9. Describe what is happening at the molecular level when dynamic equilibrium occurs in a closed container. (Hint: what are the molecules doing while at dynamic equilibrium?)

10. Name one physical property that you could use to distinguish between molecular and ionic solids.

11. What does the term STP mean?

12. What is the effect of temperature on vapor pressure?

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13. Explain the term Normal Boiling Point and what must the air pressure be at Normal Boiling Point.

Chapter 14:

1. In a given container, name two ways in which you can increase and decrease pressure.

2. Distinguish between Boyle’s law and Charles’s Law

3. What is the combined gas law?

4. How is the ideal gas law used?

5. What is the law of partial pressures?

6. List 3 factors that can affect gas pressure.

7. If the temperature is constant, what change in volume would cause the pressure of an enclosed gas to be reduced to ¼ of its original value? Which gas law are you using in this problem?

8. The pressure on 2.50 L of N2O changes from 105 kPa to 40.5 kPa. If the temperature remains constant, what will the new volume be? Which gas law are you using in this problem?

9. If a sample of gas occupies 6.80 L at 325°C, what will its volume be at 25°C if the pressure remains constant? Which gas law are you using in this problem?

10. The pressure in a car tire is 198 kPa at 27°C. After a long drive, the pressure is 225 kPa. In a rigid container where volume remains constant, what is the temperature of the air in the tire? Which gas law are you using in this problem?

11. A gas at 155 kPa and 25°C has an initial volume of 1.00 L. The pressure of the gas increases to 605 kPa as the temperature is raised to 125°C. What is the new volume? Which gas law are you using in this problem?

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12. How are pressure and volume of a gas related at constant temperature? Explain using the terms Direct or Inverse relationship.

13. How are pressure and temperature of a gas related at constant volume? Explain using the terms Direct or Inverse relationship.

14. How are temperature and volume of a gas related at constant pressure? Explain using the terms Direct or Inverse relationship.

15. Write the formulas for Boyle’s Law, Charles’ Law, and Gay Lussac’s Law.

16. When the temperature of a rigid hollow sphere containing 685 L of helium gas is held at 621 K, the pressure of the gas is 1.89 x 103 kPa. How many moles of helium does the sphere contain? Which gas law are you using in this problem?

17. A container has a volume of 2.20 L. How many grams of nitrogen gas will the container hold at a pressure of 102 kPa and a temperature of 37°C?

18. What pressure is exerted by 0.450 moles of a gas at 25°C if the gas is a 0.650 L container?

19. In your own words, state Dalton’s Law of partial pressures.

20. How is the partial pressure of a gas in a mixture calculated?

21. Determine the total pressure of a gas mixture that contains oxygen, nitrogen, and helium if the partial pressures are Poxygen= 20.0 kPa, Pnitrogen=46.7 kPa, and Phelium=26.7 kPa.

22. Explain why heating a contained gas that is held at a constant volume increases its pressure. Chapter 15:

1. What causes the high surface tension and low vapor pressure of water?

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2. What effect does a surfactant have on the surface tension of water?

3. What is the percent by mass of water in CuSO4•5H2O

4. In the formation of a solution, how does the solvent differ from the solute?

5. Why are all ionic compounds electrolytes?

6. Which of the following substances dissolve to a significant extent in water? Explain your answer in terms of polarity.

a. CH4

b. KCl

c. He

d. MgSO4

e. Sucrose

f. NaHCO3

7. How does a suspension differ from a solution?

8. What distinguishes a colloid from a suspension and a solution?

9. Why is water an excellent solvent for most ionic and polar covalent compounds but not for a nonpolar compound?

10. What is the main distinction between an aqueous solution of a strong electrolyte and an aqueous solution of a weak electrolyte?

11. Write formulas for these hydrates: a. Sodium sulfate hexahydrate

b. Calcium chloride trihydrate

c. Barium hydroxide octahydrate

12. Name each hydrate:

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a. SnCl4•5H2O

b. FeSO4•7H2O c. BaBr2•4H2O

d. FePO4•4H2O

13. Arrange colloids, suspensions, and solutions in order of increasing particle size.

14. What is hydrogen bonding and how does it work?

15. What is surface tension?

16. What is a surfactant?

17. What is the difference between a solute, a solvent, a solution, and a suspension?

18. What is the Tyndall effect?

19. What is Brownian motion?

Chapter 16:

1. The solubility of a gas in water is 0.16 g/L at 104 kPa. What is the solubility when the pressure of the gas is increased to 288 kPa? Assume constant temperature.

2. What factor determines whether a substance will dissolve in a specific solvent?

3. What 3 factors determine the rate at which a solute dissolves?

4. What units are usually used to express the solubility of a solute?

5. A solution has a volume of 2.0 L and contains 36.0 g of glucose (C6H12O6). If the molar mass of glucose is 180 g/mol, what is the molarity of the solution?

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6. How many moles of ammonium nitrate are in 335 mL of 0.425 M NH4NO3?

7. How many milliliters of a solution of 4.00 M KI are needed to prepare 250.0 mL of 0.760 M KI?

8. If 10 mL of propanone is diluted with water to a total solution volume of 200 mL, what is the percent by volume of the propanone in the solution?

9. What are two ways of expressing the concentration of a solution as a percent?

10. How many grams of K2SO4 would you need to prepare 1500 mL of 5% K2SO4 (m/v) solution?

11. An equal number of moles of KI and MgI2 are dissolved in equal volumes of water. Which solution has the higher:

a. Boiling point?

b. Vapor pressure?

c. Freezing point?

12. Define the following terms: solubility, saturated solution, and unsaturated solution.

13. Knowing the molarity of a solution is more meaningful than knowing whether a solution is dilute or concentrated. Explain.

14. What are colligative properties? Identify 3 colligative properties and explain why each occurs.

15. Compare the terms molarity and dilution.

Chapter 17:

1. When solid barium hydroxide octahydrate (Ba(OH)2 • 8H2O) is mixed in a beaker with solid ammonium thiocyanate (NH4SCN) a reaction occurs. The beaker becomes very cold. Is this reaction exothermic or endothermic?

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2. When 435J of heat is added to 3.4g of olive oil at 21°C the temperature will increase to 85°C. What is the specific heat of olive oil?

3. How much heat is required to raise the temperature of 250.0g of mercury 52°C?

4. How do endothermic processes differ from exothermic processes?

5. On what factors does the heat capacity of an object depend?

6. When 50.0mL of water containing 0.50 mol HCl at 22.5°C is mixed with 50.0mL of water containing 0.50 mol NaOH in a calorimeter, the temperature of the solution increases to 26.0°C. How much heat (in kJ) was released by this reaction?

7. When carbon disulfide is formed from its elements, heat is absorbed. Calculate the amount of heat (in kJ) absorbed when 5.66g of carbon disulfide is formed. C(s) + 2S2(s) → CS2(l) ∆H= 89.3kJ

8. How are enthalpy changes treated in a chemical reaction?

9. How many kilojoules of heat are required to melt a 10.0g popsicle at 0°C? Assume the popsicle has the same molar mass and heat of fusion as water.

10. How much heat is absorbed when 63.7g H20(l) at 100°C at 101.3kPa is converted to steam at 100°C? Express your answer in kj.

11. How much heat in (kJ) is released when 0.667 mol of NaOH(s) is dissolved in water?

12. How many moles of NH4NO3(s) must be dissolved in water so that 88.0kJ of heat is absorbed from the water?

13. Define potential energy in terms of chemistry.

14. Why do you think it is important to define system and surroundings?

15. Describe the sign convention that is used thermochemical calculations.

16. What is the function of a calorimeter?

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17. What information is given in a thermochemical equation?

18. What is the difference between specific heat capacity and heat capacity?

19. What is the difference between a calorie and a joule?

20. What is the difference between endothermic and exothermic?

Chapter 18:

1. How is the rate of a chemical reaction expressed?

2. What are four factors that affect the rate of a chemical reaction?

3. Does every collision between reacting particles lead to products? Explain.

4. How is the equilibrium position of this reaction affected by the following changes? C(s) + H20(g) + heat CO(g) + H2(g)

a. Lowering the temp b. Increasing the pressure c. Removing hydrogen d. Adding water vapor

5. The reversible reaction N2(g) + 3H2(g) 2NH3(g) produces ammonia, which is a fertilizer. At equilibrium, a 1-L flask contains 0.15mol H2, 0.25mol N2, and 0.10mol NH3. Calculate Keq for the reaction.

6. For the same mixture, under the same conditions described in problem 7, calculate Keq for 2NH3(g) N2(g) + 3H2(g). How is the Keq for a forward reaction related to the Keq for a reverse reaction?

7. At 750 the following reaction reaches equilibrium in a 1-L flask. H2(g) + CO2 H20(g) + CO(g) Analysis of the equilibrium mixture gives the following results: H2 = 0.053mol, CO2 = 0.053mol, H2O =

0.047mol, and CO = 0.047mol. Calculate the Keq for the reaction.

8. How do the amounts of reactants and products change after a reaction has reached chemical equilibrium?

9. What are the three stresses that can upset the equilibrium of a chemical reaction?

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10. Using the following equilibrium constants for several reactions, determine in which reactions the products are favored? Why?

a. Keq = 1×102 b. Keq = 0.003 c. Keq = 3.5

11. What are two characteristics of spontaneous reactions?

12. What two factors determine the spontaneity of a reaction?

13. How is the rate of a reaction influenced by a catalyst? How do catalyst make this possible?

14. What is LeChatelier’s principle? Use it to explain why carbonated drinks go flat when their containers are left open.

15. The products in a spontaneous process are more ordered that the reactants. Is the entropy change favorable or unfavorable?

16. Predict the direction of the entropy change in each reaction a. CaCO3(s) → CaO(s) + CO2(g) b. NH3(g) + HCl(g) → NH4Cl(s)

17. What is the difference between spontaneous and nonspontaneous reactions?

18. What is entropy?

19. Why do reactions a equilibrium tend to shift back and forth when pressure or other factors change?

20. What is the formula to determine the equilibrium constant for a reaction? Chapter 19

1. Identify the following acids as monoprotic, diprotic, or triprotic. Explain your reasoning.

a. H2CO3

b. H3PO4

c. HCl

d. H2SO4

2. Classify each solution as acidic, basic, or neutral.

a. [H+] = 6.0 x 10-10 M

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b. [OH-] = 3.0 x 10-2 M

c. [H+] = 2.0 x 10-7 M

d. [OH-] = 1.0 x 10-7 M

3. Find the pH of each solution.

a. [H+] = 1 x 10-4 M

b. [H+] = 0.0015 M

4. Calculate the [H+] for each solution.

a. pH = 5.00

b. pH = 12.83

5. Calculate the pH of each solution.

a. [OH-] = 4.3 x 10-5 M

b. [OH-] = 4.5 x 10-11 M

6. What is the relationship between [H+] and [OH-] in an aqueous solution?

7. What are the hydroxide-ion concentrations for solutions with the following pH values?

a. 6.00

b. 9.00

c. 12.00

8. How many moles of potassium hydroxide are needed to completely neutralize 1.56 mol of phosphoric acid?

9. How many moles of sodium hydroxide are required to neutralize 0.20 mol of nitric acid?

10. What is the molarity of H3PO4 if 15.0 mL is completely neutralized by 38.5 mL of 0.150 M NaOH?

11. What are the products of a reaction between an acid and a base?

12. Write complete balanced equations for the following acid-base reactions.

a. H2SO4(aq) + KOH(aq)

b. H3PO4(aq) + Ca(OH)2(aq)

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c. HNO3(aq) + Mg(OH)2(aq)

13. Classify each compound as an Arrhenius acid or an Arrhenius base.

a. Ca(OH)2+ d. C2H5COOH

b. HNO3 e. HBr

c. KOH f. H2SO4

14. Label the conjugate acid-base pairs in each equation below

a. HNO3 + H2O H3O+ + NO3-

b. CH3COOH + H2O ⇆ H3O+ + CH3COO-

c. NH3 + H2O ⇆ NH4+ + OH-

d. H2O + CH3COO- ⇆ CH3COOH + OH-

15. What is the molarity of sodium hydroxide if 20.0 mL of the solution is neutralized by each of the following

1.00 M solutions?

a. 28.0 mL of HCl

b. 17.4 mL of H3PO4

16. What is an amphoteric substance?

17. Using the pH scale, what are acids, bases, and neutral numbers?

18. What is an Arrhenius acid and base?

19. What is a Bronsted-Lowry acid and base?

20. What is a conjugate acid and base?

21. What is a Lewis acid and base?

22. What is the difference between hydrogen ion concentration and hydroxide ion concentration?

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23. How do you determine pH using the hydrogen ion concentration of a solution?

24. What is a titration? Chapter 20

1. Determine what is oxidized and what is reduced in each reaction. Identify the oxidizing agent and reducing

agent in each case.

a. 2Na(s) + S(s) Na2S(s)

b. 4Al(s) + 3O2(g) 2Al2O3(s)

2. Identify these processes as either oxidation or reduction.

a. 2I- I2 + 2e-

b. Zn2+ + 2e- Zn

3. Define oxidation and reduction in terms of the gain or loss of electrons.

4. Use electron transfer or electron shift to identify what is oxidized and what is reduced in each reaction. Use

the electronegativity values in Table 6.2 in Chapter 6, for molecular compounds.

a. 2Na(s) + Br2(l) 2NaBr(s)

b. H2(g) + Cl2(g) 2HCl(g)

c. 2Li(s) + F2(g) 2LiF(s)

d. S(s) + Cl2(g) SCl2(g)

e. N2(g) + 2O2(g) 2NO2(g)

f. Mg(s) + Cu(NO3)2(aq) Mg(NO3)2(aq) + Cu(s)

5. Identify the reducing agent and the oxidizing agent for each problem in #4.

6. Determine the oxidation number of chlorine in each of the following substances.

a. KClO3 c. Ca(ClO4)2

b. Cl2 d. Cl2O

7. Identify which atoms are oxidized and which are reduced in each reaction.

a. 2H2(g) + O2(g) 2H2O(l)

b. 2KNO3(s) 2KNO2(g) + O2(g)

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8. How is a change in oxidation number related to the process of oxidation and reduction?

9. Identify which of the following are oxidation-reduction reactions. If a reaction is a redox reaction, name the

element oxidized and the element reduced.

a. Mg(s) + Br2(l) MgBr2(s)

b. H2CO3(aq) H2O(l) + CO2(g)

10. Balance each redox equation using oxidation-number change method.

a. KClO3(s) KCl(s) + O2(g)

b. HNO2(aq) + HI(aq) NO(g) + I2(s) + H2O(l)

11. Balance each redox equation using the oxidation-number change method.

a. Bi2S3(s) + HNO3(aq) Bi(NO3)3(aq) + NO(g) + S(s) + H2O(l)

b. SbCl5(aq) + KI(aq) SbCl3(aq) + KCl(aq) + I2(s)

12. Which of these statements is false?

a. The oxidation number of an uncombined element is zero.

b. The sum of the oxidation numbers of the atoms in a polyatomic ion must equal the charge of the ion.

c. Every element has a single oxidation number.

d. The oxidation number of oxygen in a compound or polyatomic ion is almost always -2.

13. Assign oxidation numbers to the atoms in the following ions:

a. Ca+2

b. Al2S3

c. Na2CrO4

d. V2O5

e. MnO4-

14. What is the difference between a redox reaction and a non-redox reaction?

Chapter 25

1. List the symbol for an alpha, beta and gamma particle. Which of the three has no mass or no charge?

2. 𝐾 → 𝑒 + ___________−10

1942 _____________________

3. 𝐿𝑖 → 𝐻𝑒 + ___________24

36 _____________________

4. The half-life of radon-222 is 3.8 days. How much of the 100.0-g sample is left after 15.2 days?