ChemistryFinal Exam ReviewName:

Review Sheet #1Handouts Section

Date Given:

A solid method for reviewing for the final is as follows: 1) review the following list of objectives, 2) look through and rework old tests and quizzes, 3) read through the book, 4) review old homework problems, 5) read through notes, 6) pay attention during review sessions, and 7) come to class with questions.

When studying, do not make the mistake of telling yourself that you know the material. Convince yourself by working out problems and testing yourself. Practice, practice, practice!!!

Extra Credit Opportunities:

(0-10+ points) For every double-sided sheet of paper that you use to study for the final (working out problems, re-writing notes, correcting test problems), you will receive one point credited toward your final exam score (10 point maximum).

(0-5 points) Book, Binder, Calculator checks: There will be 5 days that are dedicated to working on and presenting review projects. For each day that you bring your Book, Binder, AND Calculator to class on these days, you will receive one extra credit point on the Syllabus.

Chapter 7 - Chemical Reactions

  • Identify the reactants and products in a chemical reaction.
  • Rewrite a chemical equation from words into symbols.
  • Write and balance chemical equations.
  • Identify various types of chemical reactions:
  • combination, decomposition, single-replacement, double replacement, combustion

Chapter 8 – Stoichiometry

  • Calculate the amount of reactants required or product formed in a nonchemical process.
  • Interpret balanced chemical equations in terms of interacting moles, representative particles (molecules or formula units), masses, and volume.
  • Calculate stoichiometric quantities from balanced chemical equations using moles, mass, particles (molecules or units), and volumes of gases at STP.
  • *Identify the limiting reagent in a reaction and use it to calculate stoichiometric quantities and the amount of excess reagents.
  • *Calculate the theoretical yield, actual yield or Percent yield for a chemical reaction.

Chapter 9 - Gases

  • Describe the motion of particles of a gas according to the kinetic theory.
  • Interpret gas pressure in terms of kinetic theory.
  • Convert between units of gas pressure.
  • Show that temperature is a measure of the Kinetic Energy of particles in a substance
  • Describe the nature of a liquid in terms of the attractive forces between particles
  • Explain vaporization of liquids, using kinetic theory
  • Describe what happens on a particle level at the boiling point of a liquid.
  • Describe how you can distinguish between particles in solids from gases and liquids.
  • Explain sublimation
  • *Interpret the phase diagram of water.

Chapter 11 - Gases

  • Infer the effects of adding or removing gas and changing the dimensions of a container.
  • Calculate pressure or volume of a contained gas at constant temperature
  • Infer from kinetic theory the effect of temperature changes on a contained gas.
  • Calculate temperature or volume of a contained gas at constant pressure.
  • Calculate temperature or pressure of a contained gas at constant volume.
  • Calculate any variable using the ideal gas law.
  • Calculate the total pressure of a mixture of gases.
  • Explain why equal volumes of gases contain the same number of particles.
  • *Explain Graham's Law and calculate masses of gases.

Chapter 10 - Thermodynamics

  • Distinguish among various forms of energy: stored energy, work, and heat.
  • Understand the law of conservation of energy
  • Explain the heat capacity of objects and express it in standard units of heat.
  • Identify different substances using specific heat capacities.
  • Describe heat changes in terms of a system and its surroundings.
  • Calculate the heat changes that occur in chemical and physical processes.
  • Use a calorimeter in experiement to determine heat transfers and specific heats.
  • Construct equations that show the heat changes for chemical and physical processes.
  • Describe in words and with diagrams the heat changes that occur in melting, freezing, boiling, and condensing.
  • *Apply Hess’s law of heat summation to find heat changes for chemical and physical processes.
  • Calculate heat changes, using standard heats of formation.

Chapter 17 - Solutions

  • List three factors that determine how fast a soluble substance dissolves.
  • Explain the difference among saturated, unsaturated, and supersaturated solutions.
  • Define and work problems involving the molarity of a solution.
  • Distinguish between dilute and concentrated solutions.
  • Describe how to prepare dilute solutions from concentrated solutions of known molarity.
  • *Calculate percent by volume and percent by mass for solutions.
  • Describe three colligative properties (vapor pressure, boiling point elevation, and freezing point depression)
  • *Calculate the molality and mole fraction of a solution.
  • *Calculate new freezing and boiling points.
  • Distinguish between colloids and suspensions from solutions.

Chapter 18 - Reaction Rates and Equilibrium

  • Interpret and express the meaning of the rate of a chemical reaction.
  • Explain how the rate of a chemical reaction is influenced by the temperature, concentration, particle size of reactants, and catalysts using collision theory.
  • Define chemical equilibrium in terms of a reversible reaction.
  • Predict changes in the equilibrium position due to changes in concentration, temperature, and pressure using Le Chatelier’s principle.
  • Define free energy and use the concept to contrast spontaneous and nonspontaneous reactions.
  • Show how changes in entropy relate to a change of state, a change in temperature, and a change in the number of product particles compared with reactant particles.
  • Explain how changes in energy and changes in entropy both influence the spontaneity of a reaction.
  • *Write the equilibrium constant for a reaction and compute its value from experimental data.
  • *Compute the change in entropy of a reaction using standard entropies.
  • *Analyze the reaction mechanism for a reaction given a potential energy diagram.

Chapter 19 - Acids and Bases

  • List properties of acids and bases
  • Name an acid and base when given the formula and vice versa.
  • Classify a solution as neutral, acidic, or basic, given the [H+] or [OH-] concentration.
  • Calculate the pH of a solution given the [H+] or [OH-] concentration and vice versa.
  • Define and give examples of Arrhenius acids and bases.
  • Classify substances as acids or bases using the Bronsted-Lowry theory.
  • Identify conjugate acid-base pairs using the Bronsted-Lowry tehory.
  • Describe at least two methods used to measure pH.
  • Demonstrate your knowledge of acid-base reactions by completing and balancing a neutralization reaction.
  • Explain the steps of a titration.

The test will consist of the following types of questions:

multiple choice (approximately 75 questions) = 150 points

short answer problems (approximately 10 questions for Honors only) = 50 points/ 200 total pts Honors

The test will be worth 150 points (Honors - 200) and will count for approximately 20% of your final grade. Unless a specific circumstance arises, you are expected to finish the exam during the allotted exam time.

UNDER NO CIRCUMSTANCE WILL YOU BE ALLOWED TO FINISH THE EXAM ON ANOTHER DAY.

Chapter 7 Objectives

(10) 16. Write a skeleton equation for the following reactions. Show all symbols of state. For a bonus point, go ahead and balance!

a) solid iron combines with oxygen gas to form solid iron(III)oxide.

b) aqueous sodium chloride reacts with aqueous calcium oxide to produce aqueous sodium oxide and solid calcium chloride.

(9) 17. Balance the following equations.

a) ______CS2(s) + ______O2(g) → ______CO2(g) + ______SO2(g)

b) ______KBrO3(s) → ______KBr(s) + ______O2(g)

c) ______Ca(s) + ______H3PO4(aq) →______Ca3(PO4)2(s) + ______H2(g)

(9) 18. For each equation in question 17, identify the type of reaction.

a)______b)______

c)______

(12) 19. Predict the products and balance each of the following reactions:

Decomposition: a) H2O →

Single Replacement: b) Al(s) + Fe(NO3)2→

Double Replacement: c) ZnI2(aq) + NaOH(aq) 

Chapter 8 Objectives

8. Interpreting Chemical Equations (10 pts)

a. Balance the following equation.

______Ca + ______O2  ______CaO

b. Interpret the above equation in terms of representative particles (atoms, molecules, and/or formula units.)

9. Mole-Mole Calculations (10 pts) ☺

a Balance the following equation.

______N2 + ______H2  ______NH3

b. If 12.0 moles of H2(g) react with excess N2(g), how many moles of ammonia are produced.

10. Mole-Mass Calculations (8 pts) ☺

CaC2(s) + 2H2O(l)  C2H2(g) + Ca(OH)2(s)

Given the balanced equation above for the reaction of calcium carbide and water, how many grams of calcium carbide, CaC2 will react with 3 moles of water?

11. Mass-Mass Calculations (10 pts) ☺

How many grams of carbon dioxide are formed when 50 grams of Fe2O3react with excess carbon monoxide?

Fe2O3 + 3 CO  2 Fe + 3 CO2

12. Other Stoichiometric Calculations (8 pts) ☺

If aluminum reacts with oxygen according to the following,

4Al(s) + 3O2(g) →2Al2O3(s)

what volume (in Liters) of oxygen gas will be needed to form 1.00 gram of aluminum oxide at STP? (1 mole = 22.4 L at STP)

Chapter 9 Objectives

Write a one paragraph response in which you distinguish between gases, liquids, and solids in terms of the kinetic theory.

a) A gas is at a pressure of 3.0 atm. What is this pressure in kilopascals?

b) A gas is at a pressure of 410 mmHg. What is this pressure in atm?

c) A gas is at a pressure of 600 kPa. What is this pressure in mmHg?

Chapter 11 Objectives

A sample of nitrogen occupies a volume of 0.600 L at 273K. What volume would the gas occupy at 373K if the pressure remains constant?

A gas has a pressure of 600 kPa at 200 K. What will its pressure be at 400 K, if the volume does not change?

A sample of hydrogen occupies a volume of 2.0 L at a pressure of 400 kPa. If the temperature of the gas is kept constant, what would be the new volume of the gas at 600 kPa?

A sample of oxygen gas occupies a volume of 1.25 L at -23°C. If the pressure remains constant, what is the new Celsius temperature of the gas if its volume decreases to 0.925 L?

Combined and Ideal Gas Laws (use R = 0.0821 L atm / mol K, or R = 8.314 L kPa / mol K)

A gas occupies a volume of 0.5 L at 300 K and 100.0 kPa. What is the volume of the gas at conditions of STP (0°C, 101.3 kPa)?

How many moles of N2 are in a flask with a volume of 0.200 L at a pressure of 350 kPa and a temperature of 400 K?

What volume is occupied by 0.75 moles of any gas at STP?

How many moles of Cl2 are in 19.3 L of chlorine gas at STP?

Chapter 10 Objectives

Determine the specific heat of a material if a 10-gram sample absorbed 100 J as it was heated from 20ºC to 60ºC.

How much heat is absorbed by a 200 gram piece of granite as energy from the sun causes its temperature to change from 20°C to 50°C? The specific heat of granite is 0.803 J/g °C.

Given the equation 3CO(g) + Fe2O3(s)  2Fe(s) + 3CO2(g) + 24.7 kJ, how much heat is released when 3.0 mol of Fe2O3 react?

How much heat is needed to melt 2.0 moles of ice? ∆Hfusion for water = 6.01 kJ/mol

How many moles of NH4NO3 would need to be dissolved in order for the solution to absorb 125 kJ of energy? ∆Hsolution = 25.7 kJ/mol

Use standard enthalpies of formation from Table C-13 (attached) to calculate ΔHreaction for each of these reactions.

a. 2H2S(g) + 3O2(g) →2H2O(g) + 2SO2(g)

Chapter 17 Objectives

Calculate the molarity of a solution that contains 50.0 g of NaCl per 0.6 L of solution.

How many moles of solute are present in 3.5 L of a 0.50 M LiNO3 solution? How many grams of solute are there?

What is the percent (m/v) of a water solution that contains 80 g of NaOH, and that has a volume of 350 mL?

How many milliliters of alcohol are in 500 mL of 75.0% (v/v) solution?

How would you prepare 500 mL of 0.5M HCl solution from a 6.0 M HCl stock solution?

Calculate the mole fraction of ethanol in a solution containing 1.50 moles of ethanol and 4.50 moles of water.

Calculate the molality of a solution prepared by dissolving 200 g of AlCl3 in 1250 g of water.

Calculate the boiling point of a solution that contains 0.900 mol of CaCl2 dissolved in 3000 g of water. (Kb for water = 0.512 ºC/m and normal b.p. = 100ºC)

Suppose a solution of Kool-Aid and water can hold 75 grams of Kool-Aid for every 100 mL of water. If you dissolve 65 grams of Kool-Aid in 100 mL of water, what kind of solution (saturated, unsaturated, supersaturated) do you have? Explain.

Chapter 18 Objectives

2NOCl(g) ↔ 2NO(g) + Cl2(g)

a. Write the equilibrium constant expression (Keq) for the above reaction.

b. With the following concentrations, calculate the Keq for the above reaction: [NOCl] = 0.50M, [NO] = 1.0M, and [Cl2] = 0.25M.

Determine whether entropy is increasing or decreasing in each of the following reactions.

a) CaCO3(s)  CaO(s) + CO2 (g)

b) 2H2(g) + O2(g)  2H2O (l)

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

4HCl(g) + O2(g) ↔ 2H2O(g) + 2Cl2(g) + energy

Stress/Disturbance / Predicted Shift / Explanation
Increasing temperature
(note that reaction is endothermic)
Increasing pressure
Adding H2O
Removing O2(g)

Chapter 19 Objectives

Write the formulas for the following acids or bases.

hydroiodic acidsulfuric acidlithium hydroxidecalcium hydroxide

Name the following acids or bases from their formulas.

HClHNO3NaOHMg(OH)2

What is the pH of a solution in which [H+] = 4.0 x 10-10 M?

If the [OH-] = 1 x 10-2 M, what is the pH of the solution?

What is the [H+] concentration if the pH of a solution equals 4.7?

Write complete balanced equations for the following acid-base reactions. Also, what is the name of the salt produced in each reaction?

a. HCl(aq) + NaOH (aq) →b. Ca(OH)2(aq) + H3PO4(aq) 

Identify the conjugate acid/base pairs in the following:

HNO3 (aq) + NH3 (aq) → NH4+ (aq) + NO3- (aq)

What is the molarity of sodium hydroxide if 20.0 mL of the NaOH solution is neutralized by the following solutions?

a. 40.0 mL of 1.00M HClb. 60.0 mL of 0.50M H2SO4

How many millilters of a 2.0 M solution of HCl are needed to titrate the following solutions?

a. 200 mL of 3.0M KOHb. 125 mL of 0.60M Ca(OH)2

The test will consist of the following types of questions:

multiple choice (approximately 75 questions) = 150 points

short answer problems (approximately 10 questions for Honors only) = 50 points/ 200 total pts Honors

The test will be worth 150 points (Honors - 200) and will count for approximately 20% of your final grade. Unless a specific circumstance arises, you are expected to finish the exam during the allotted exam time.

UNDER NO CIRCUMSTANCE WILL YOU BE ALLOWED TO FINISH THE EXAM ON ANOTHER DAY.

Chem2ndFinalExamReviewSheet2005_06.doc, 11/14/18