Ms. Agostine’s College Prep Chemistry Final Exam … Final Pkt.pdf · Ms. Agostine’s College Prep Chemistry Final Exam ... Calculate how much product is made from the limiting

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Ms. Agostine’s College Prep Chemistry Final Exam Review Packet for 2014-2015

ALL MAKE UP WORK MUST BE TURNED IN BY FRIDAY, JUNE 5, 2015, 2:35 PM

NO EXCEPTIONS!!!

FINAL EXAM REVIEW DAY: WEDNESDAY, JUNE 10, 2015

EXAM SCHEDULE:

THURSDAY JUNE 11: PERIOD 1 & 5

FRIDAY, JUNE 12: PERIOD 2 & 6

MONDAY, JUNE 15: PERIOD 3 & 7

TUESDAY, JUNE 16: PERIOD 4 & 8 (LAST DAY

OF SCHOOL)

EXAM SCHEDULE:

FIRST EXAM OF THE DAY: 7:40 AM – 10:30 AM

o (STUDENTS CANNOT LEAVE BEFORE 9:15 AM)

Early dismissal bus schedule will be posted at a later date

LUNCH WILL BE SERVED EVERYDAY FROM 10:00 AM – 11:40 AM

SECOND EXAM OF THE DAY: 11:45 AM – 2:35 PM

o (STUDENTS CANNOT LEAVE BEFORE 1:15 PM)

Late arrival bus schedule will be posted at a later date

LIBRARY AND CAFETERIA WILL BE OPEN DURING ALL TESTING TIMES

HELPFUL HINTS:

Reread textbook chapters!

Go over notes and highlight key terms in your chapter packets.

Review old tests (tests folders cannot leave the room), homework’s, quizzes,

and labs.

o Make a separate outline of what you didn’t understand well.

o Fill in the outline with info from text book, notes, etc.

Redo sample and practice problems as well as section reviews from the book

and end of chapters.

Always remember to check your units! Does everything cancel out?

Use the “E” in the calculator to represent “x 10 ^”

Ask questions; come to class prepared to review!

Email: [email protected] with questions… GOOD LUCK!

Go to www.DorettaAgostine.com and read through the notes

Go to the “Quia” page and play review games (link is on my web site)

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Chapters & Topics Covered on Final Exam

Chapter 8 – Chemical Reactions

Balancing Equations

Reaction Types

Predicting Reactions

Chapter 7 – Chemical Quantities

Mole concept

Mole calculations & road map

Percent composition

Empirical formula

Molecular formula

Chapter 9 – Stoichiometry

Mole Ratios

Limiting Reagent

Chapter 10 – States of Matter

Solids

Liquids

Gases

Phase Changes & Graphs

Chapter 11 – Thermochemistry

Heat & Specific Heat Calculations

Molar Heat of Fusion Calculations

Chapter 12 – Gas Laws

Boyle’s, Charles’, Gay Lussac’s Laws

Combined Gas Law & Ideal Gas Law

Dalton’s Law of Partial Pressure

Chapter 17/18 – Water & Aqueous Systems & Solutions

Water Properties

Solutions, Suspensions, Colloids

Molarity

Solubility

Chapter 20/21 – Acids & Bases

Properties of Acids and Bases

Types of Acids & Bases, Definitions

pH / pOH

Chapter 28 – Nuclear Chemistry

Radioactive Decay and Types of Decay

Radioactivity and Half-Life

Nuclear Equations

The final exam will consist of 70 multiple-choice questions and 12 open-ended questions or

problems.

YOU MUST LEGITIMATELY GIVE YOUR BEST EFFORT TO BE ELIGIBLE FOR THE CURVE! NO BLANKS,

NO ABBACADABBA PATTERNS!

You MUST turn in the correct textbook or the appropriate $$ amount when you take the final

exam

o Textbooks are due the day of the final. They will NOT be accepted earlier.

o Cash or Check made out to “Abington Senior High School”

o (Text = $65, lab manual = $13)

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REVIEW CONCEPTS AND PRACTICE PROBLEMS

I. Chapter 8: Chemical Reactions

A. Describing Chemical Change

1. Using chemical equations to describe a chemical change

a) Reactants Products

b) Coefficients

c) Subscripts

d) Chemical formulas

e) Chemical nomenclature

f) Symbols for solid, liquid, gas and aqueous

(1) What is the difference between liquid and aqueous?

g) Law of Conservation of Mass and Matter

B. Types of Chemical Reactions (know general reactions: A + B, AB + CD…)

1. Combination/Synthesis reactions

2. Decomposition reactions

3. Combustion reactions (need which diatomic element?)

4. Single-replacement reactions

5. Activity series of a metal

6. Double-replacement reactions

7. Solubility table

8. Precipitates

9. Aqueous solutions

C. Sample Problems:

1. Complete, balance, and type the following. Use your Activity Series and

Solubility Table.

a) sodium + oxygen

b) electrolysis of water

c) lithium hydroxide + iron (III) nitrate

d) calcium hydroxide and phosphoric acid

2. Write the complete ionic equation and the net ionic equation for: Barium

nitrate plus potassium phosphate.

complete:

net:

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II. Chapter 7: Chemical Quantities

A. A Mole: A Measurement of Matter

1. What is a mole?

a. Avogadro’s Number

b. Representative particle (mo, at, fu)

c. Diatomic elements

d. Calculate the Molar Mass (where do you look?)

i. Gram atomic mass

ii. Gram molecular mass

iii. Gram formula mass

B. Mole-Mass and Mole-Volume Relationships

1. Mole-mass conversions

2. Mole-volume conversions

3. STP

a. Temp

b. Pressure

4. Molar volume of a gas

5. Mole Road Map (draw it!)

C. Percent Composition and Chemical Formulas

1. Percent composition

2. What do all the percentages of each element have to add up to?

3. Empirical formula

4. Calculating molecular formulas from empirical formulas: what is the whole

number multiple?

D. Sample Problems

1. How many representative particles are:

a) 3.00 mol Sn, b) 0.400 mol KCl, c) 7.50 mol SO2

2. Calculate the molar mass of the following and identify if it is a GAM, GMM, or

a GFM

b) H3PO4, b) N2O3, c)CaCO3, d) (NH4)2SO4, e) C4H9O2

3. How many moles is 15.5 g of silicon dioxide?

4. How many liters does 7.6 mol of argon gas occupy at STP?

5. What is the density of butane gas (C3H8) in units of g/mL at STP? (Hint: start

with the molar mass)

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6. How many liters does 18.1 g of oxygen gas occupy at STP?

7. What is the percent composition of the elements in ammonium oxalate?

8. What is the empirical formula of a substance that is 25.9% nitrogen and 74.1%

by mass?

III. Chapter 9: Stoichiometry

A. Arithmetic of Equations

i. Coefficients tell you the relative number of moles of reactants and

products

ii. All equations must be balanced to abide by the law of conservation of

matter

iii. The number and kinds of atoms in the reactants equation the number and

kinds of atoms in the products.

B. Chemical Calculations

i. Mole-Mole problems

ii. Mole-Mass Problems and Mass-Mole problems

iii. Mass-Mass problems

iv. Volume-volume problems

v. Particle-particle problems

vi. Mixed problems (mass-volume, volume-particle…)

C. Limiting Reagent and Percent Yield

i. Limiting reagent is used up first

ii. Must use the coefficients of the balanced equation to determine the

limiting reagent.

iii. Calculate how much product is made from the limiting reagent

iv. Percent yield equation:?

1. Theoretical yield

v. Actual yield

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D. Sample Problems: Balance each equation and then answer the question below.

C2H6 + O2 CO2 + H2O

a) If you had 8 moles of ethane, how many moles of oxygen would you

need?

b) If you had 3 L of oxygen gas, how many liters of carbon dioxide would you

make?

c) If you had 65 g of ethane, how many grams of water would you make?

d) If you started with 3 g of ethane and 9 g of oxygen, which would be the

limiting reagent?

e) How many grams of water would I make from that limiting reagent in d?

8. Chlorobenzene (C6H5Cl) is used in the production of many important chemicals

such as aspirin, dyes, and disinfectants. One industrial method of preparing

chlorobenzene is the reaction between benzene (C6H6) and chlorine gas. The

products are chlorobenzene and hydrochloric acid. Two students were able to

make 36.8 g of the chlorobenzene in lab. What is their percent yield if the actual

yield is 39.0 g?

E. Chapter 10: States of Matter

a. Nature of Gases

i. Kinetic theory

1. gases are mostly empty space (this is why you can compress them)

2. collisions are elastic

ii. STP

iii. Volume of any gas at STP

b. Nature of Liquids

i. Vapor pressure

ii. Evaporation

iii. Vaporization

iv. Boiling point

v. Normal boiling point

vi. What happens to the boiling point of liquids at different elevations?

c. Nature of Solids

i. Melting point

ii. Crystalline structure of NaCl

iii. Allotropes of Carbon = 3 forms

d. Changes of State

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i. Phase change diagram

1. triple point

2. critical point

ii. Sublimation & deposition

e. Practice: Describe the kinetic theory and how it relates to gases.

i. Create a chart comparing and contrasting the 3 states of matter.

ii. Define and describe the following phase changes, include WHEN they will

occur and what is happening with the particles:

1. Melting

2. Freezing

3. Boiling

4. Evaporation

5. Condensation

6. Sublimation

f. Practice: Describe the major areas and points on a phase diagram. Draw a

sample phase diagram.

g. Practice: Explain the relationship between kinetic energy and temperature.

F. Chapter 12: The Behavior of Gases

a. The Properties of Gases

i. Kinetic molecular theory can be used to explain gas pressure, volume and

temperature

ii. Kelvin temperature and absolute zero

b. Factors Affecting Gas Pressure

i. Volume

ii. Pressure

iii. Temperature

iv. Number of moles

c. Gas Laws

i. Necessary units for temperature?

ii. Which are direct and inversely proportional relationships for:

1. Boyle’s Law

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2. Charles’ Law

3. Gay-Lussac’s Law

4. Combined Gas Law

d. Ideal Gases

i. Why were all the other equations technically incorrect?

ii. PV=nRT

iii. R = gas constant, Units?

e. Gas Molecules: Mixtures and Movements

i. Dalton’s Law of Partial Pressures

f. Practice: Explain and give the formula for each law:

i. Boyle’s Law

ii. Charles’ Law

iii. Gay-Lussac’s Law

iv. Combined Gas Law

v. Ideal Gas Law

vi.

g. Sample Problems

9. Pressure = 245 kPa, Volume changes from 47 to 74 mL. What is the new

pressure?

10. Temperature changes from 24 degree C to 35 degree C, if the initial volume

was 800 mL what is the new volume?

11. If the pressure drops from 760 torr to 600 torr, what happens to a gas that is 50 oC?

12. If a 1 L of a gas starts at STP, what is its new volume at 25 degree C and 770

torr?

13. If 45 L of natural gas, which is essentially methane, undergoes complete

combustion at 730 mm Hg and 20 degree C, how many grams of CO2 are

produced?

14. How many molecules are in 4.5 L of water at STP?

G. Chapter 17: Water & Aqueous Systems

a. Water molecule: draw it! Looks like who?

b. Polarity and solubility: Rule of thumb is = _________ dissolves __________

c. Properties

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i. Surface tension

ii. Specific heat capacity – why is it so high?

iii. Evaporation and condensation

iv. Why does ice float?

d. Aqueous solutions

i. Solute

ii. Solvent

iii. Solution

iv. Types: How do you know? How can you test for them?

1. electrolytes

2. nonelectrolytes

e. Heterogeneous Aqueous systems

i. Suspensions

ii. Colloids & Tyndall effect

H. Ch 18: Solutions

1. Factors that affect solubility

a. Surface area

b. Agitation/mixing

c. Temperature

d. Polarity: Like dissolve Like

2. Graph (Know trend): Amount of Solid solute that dissolves with temperature

3. Graph (Know trend): Amount of gases solute that dissolves with temperature

4. Concentration Units

a. Molarity (M): definition?

5. Sample Problems

21. What are electrolytes? How do they differ from nonelectrolytes?

22. Which of the following are electrolytes: NaCl, HCl, C12H22O11, CaCl2,

P2O5¸Ca3(PO4)3?

23. How can you increase the solubility of a solute?

24. What is the difference between a saturated solution, unsaturated and super-

saturated solution?

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25. What happens to the solubility of a solid as temperature increases?

26. What happens to the solubility of a gas as temperature increases?

X. Chapter 11: Thermochemistry- Heat and Chemical Change

a. The Flow of Energy-Heat

i. Thermochemistry

ii. Energy

iii. Chemical potential energy vs kinetic energy

iv. Heat (q)

v. Law of conservation of energy

vi. Exothermic process

1. Energy diagram

2. Heat change = positive or negative

3. Real life examples?

vii. Endothermic process

1. Energy diagram

2. Heat change = positive or negative

3. Real life examples?

viii. Units of heat:

1. calorie

2. Calorie & Kilocalorie

3. Joule?

ix. Specific heat

1. Independent or dependent of the mass of the object?

2. Equation

3. Units?

4. Table 11.2 p. 296

b. Measuring and Expressing Heat Changes

i. Calorimetry

ii. Calorimeter

1. Example? Think of the food lab…

iii. Enthalpy (H)

1. Equation? Similar to what?

iv. Thermochemical equations

1. Exothermic

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a. Produces heat

2. Endothermic

a. Requires heat

3. Can be used as conversion factors in stoichiometry problems

4. Types

a. Heat of Combustion

i. Table 11.4 p. 305

2. Sample Problems:

15. If 1 calorie = 4.184 J, how many calories is 1000 J?

16. When 425 J of heat is added to 3.4 g of olive oil at 21oC, the temperature

increases to 85oC. What is the specific heat of olive oil?

17. How much heat is required to raise the temperature of 250.0 g of mercury to

52oC from 25oC if the specific heat of mercury is 0.14 J/goC?

18. Write the thermochemical equation: when carbon disulfide is formed from its

elements, the H = 89.3 KJ.

19. Write the thermochemical equation: the production of iron and carbon

dioxide from iron (III) oxide and carbon monoxide has a H = - 26.3 KJ.

20. For # 19, how many kilojoules of heat are produced when 3.40 mol of iron (III)

oxide reacts with an excess of CO?

IX. Chapter 20-21: Acids & Bases

1. Properties of Acids

a. Acid definitions of Arrhenius, Bronsted-Lowry, Lewis

i. Donate H+ which actually forms hydronium ions

ii. What is a hydronium ion?

2. Properties of Bases

a. Base definitions of Arrhenius, Bronsted-Lowry, Lewis

3. pH Scale & Calculations

a. Scale range?

i. What part is an acid?

ii. Base?

iii. Neutral?

b. pH = - log [H3O+], pOH = - log [OH-], pH + pOH = 14

4. Neutralization Reactions

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5. Sample Problems

27. What is the difference between acids and bases?

28. What is the pH range for acids and bases?

29. What is the pH of a 0.235M HCl solution? Is it an acid, base or neutral?

30. What is the pH of a 5.2x10-4 M NaOH solution? Is it an acid, base or

neutral?

31. What is the hydronium ion concentration of a solution with a pH of 2.3?

What is the hydroxide ion concentration?

32. What is the hydronium ion concentration of a solution with a pH of

11.8? What is the hydroxide ion concentration?