Grade 11 Chemistry

30S Chemistry Exam Review

Multiple Choice

Review Unit

1. The law of conservation of mass specifically describes the behavior of substances that are undergoing

a. physical change. b. chemical change.

c. temperature change. d. phase change.

2. Which property of a sample of matter can be determined using a laboratory balance? a. density b. mass

c. homogeneity d. pressure

3. Copper is classified as an element because it a. occurs uncombined in nature. b. reacts with many nonmetallic elements. c. can be produced from its compounds. d. cannot be broken down into simpler components by ordinary chemical means.

4. Food can be cooked over a charcoal fire because the burning of charcoal is a. a chemical change that absorbs energy. b. a chemical change that gives off energy. c. a physical change that absorbs energy. d. a physical change that gives off energy.

5. Element A and element B combine chemically to form substance C. Substance C must be a. a solution. b. a compound.

c. an element. d. a mixture

6. Which fundamental particles are negatively charged?

a. alpha particles b. electrons

c. neutrons d. protons

7. A neutron and a proton have nearly the same

a. charge. b. electronegativity

c. mass. d. energy level.

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8. Two different isotopes of the same element must have the same number of

a. neutrons and electrons. b. neutrons and nucleons.

c. protons and nucleons. d. protons and electrons.

9. The atomic number of an atom is represented by Z and the mass number by A. Which expresses the number of protons in an atom?

a. A - Z b. A

c. Z - A d. Z

10. When a sodium atom becomes an ion, it

a. gains one proton. b. loses one proton.

c. gains one electron. d. loses one electron.

11. Which is an example of an ionic compound whose name ends in"-ide"?

a. H2O2 b. Cu(NO3)2

c. Al2(SO4)3

d. NH4Cl

12. The formula for the compound formed from SO42- and Al+3 is

a. (SO4)2Al3 b. Al3(SO4)2

c. Al2(SO4)3 d. (SO4)Al

13. Which name is given to the elements in Group 2?

a. alkaline earth metals b. alkali metals

c. semimetals d. halogens

14. Which element is most likely to form a positive ion?

a. O b. K

c. P d. Cl

15. Which element is most likely to form an ionic bond with chlorine?

a. potassium b. iodine

c. phosphorus d. hydrogen

16. Which kind of bond is formed when two atoms share electrons to form a molecule? a. ionic b. metallic

c. electrovalent d. covalent

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17. Multiple covalent bonds occur when more than two a. atoms are found in a molecule. b. electrons are shared by the same two atoms. c. atoms share the same pair of electrons. d. units of charge are transferred.

18. Which does NOT result in the formation of a chemical bond? a. transfer of a neutron from one atom to another b. transfer of an electron from one atom to another c. distribution of many electrons among many metal ions d. sharing of pairs of electrons between two atoms

19. A potassium ion, K+, is similar to a chloride ion, Cl-, in that both have the same a. formula mass. b. electronegativity

c. number of protons. d. number of electrons.

The Mole and Stoichiometry

1. Which is the best value of the mass of one molecule of CO2?

a. 22.01 g b. 22.01 u

c. 44.01 g d. 44.01 u

2. In the formula C3H4(OH)3 , the number of oxygen atoms is a. equal to the number hydrogen atoms

b. less than the number of hydrogen atoms. c. greater than the number of hydrogen atoms.

3. A mole of oxygen molecules consists of a. 6.02 x 1023 oxygen atoms. b. 12.04 x 1023 oxygen atoms. c. 1 oxygen atom. d. 2 oxygen atoms.

4. What is the total number of atoms represented by the formula(NH4)2HPO4 ? a. 10 b. 12

c. 16 d. 22

5. Which sample has the greatest mass?(Atomic masses: K = 39.1; Ca = 40.1; Cu = 63.5; Zn = 65.4)

a. 1.0 mole of calcium b. 1.0 mole of copper

c. 1.0 mole of potassium d. 1.0 mole of zinc

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6. What is the formula mass of CH3COOH? (Atomic masses: C = 12.0; H = 1.01; O = 16.0) a. 41.1 u b. 52.1 u

c. 60.1 u d. 94.1 u

7. What is the number of oxygen atoms in 33.6 L of CO2 at STP?

a. 3.01 x 1023 atoms b. 6.02 x 1023 atoms

c. 9.05 x 1023 atoms d. 1.81 x 1023 atoms

8. In the equation for an exothermic reaction, the energy term is written on the

a. right side of the arrow, with the reactants. b. right side of the arrow, with the products. c. left side of the arrow, with the reactants. d. left side of the arrow, with the products.

9. The equation H2 + O2 ---> H2O for the reaction between hydrogen and oxygen to produce water is unbalanced because

a. no energy term is included. b. symbols for phase have been omitted. c. too few oxygen atoms are represented as produced. d. too many molecules are represented as reactants.

10. Chemical equations become balanced by a. adjustment of subscripts of chemical formulas. b. adjustment of coefficients of chemical formulas. c. rearrangement of parentheses in chemical formulas. d. elimination of spectator ions from chemical formulas.

11. The symbol "HCl(aq)" represents a

a. hydrated salt. b. water solution of hydrochloric acid. c. pair of spectator ions. d. precipitate.

12. In the chemical equation 2H2(g) + O2(g) ---> 2H2O(l) the term to the right of the arrow means a. 2 grams of water. b. 2 liters of water.

c. 2 atoms of water. d. 2 moles of water.

13. When the equation NH4OH + CaCl2 ---> Ca(OH)2 + NH4Cl is balanced using the smallest whole-number coefficients, what is the coefficient of NH4OH?

a. 1 b. 2

c. 3 d. 4

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14. Which is the balanced equation for the reaction of aluminum with hydrochloric acid? a. 2Al + 6HCl ---> 2AlCl3 + 3H2 b. Al + HCl ---> AlCl3 + H2 c. Al2 + 6HCl ---> 2AlCl3 + 3H2 d. Al + HCl3 ---> AlCl3 + H

15. To which category does the following reaction belong? Cu + 2AgNO3 ---> 2Ag + Cu(NO3)2 a. synthesis b. decomposition

c. single replacement d. double replacement

16. Copper(II) chloride and sodium sulfide are ionic solids that dissolve readily in water. When water solutions of these two compounds are mixed, a chemical reaction takes place. The products are copper(II) sulfide, a solid that is not soluble in water, and sodium chloride, which is soluble in water.

Which is the correct equation for this reaction? a. CuCl2 (l) + Na2 S(l) ---> CuS(s) + 2NaCl(l) b. CuCl2 (aq) + Na2 S(aq) ---> CuS(s) + 2NaCl(aq) c. CuCl2 (s) + Na2 S(s) ---> CuS(aq) + 2NaCl(s) d. CuCl2 (aq) + Na2 S(aq) ---> CuS(s) + 2NaCl(s)

17. Aluminum chloride and lead (II) nitrate are both solids. When water solutions of these two compounds are mixed, a precipitate of lead chloride is formed. Which is the correct molecular equation for this reaction?

a. 2AlCl3 (aq) + 3Pb(NO3)2 (aq) ---> 3PbCl2 (s) + 2Al(NO3)3 (aq) b. AlCl3 (aq) + Pb(NO3)2(aq) ---> PbCl3(s) + Al(NO3)3(aq) c. Pb+2(aq) + 2Cl-(aq) ---> PbCl2 (s) d. AlCl3(aq) + Pb(NO3)2(aq) ---> PbCl2(s)

18. According to Avogadro's hypothesis, all 1-mole samples of different gases at the same conditions

a. are made of diatomic molecules. b. have the same mass.

c. occupy the same volume. d. have the same density.

19. In a mass-mass problem, which information is used to change from moles of substance sought to mass of substance sought?

a. formula mass of the substance given b. formula mass of the substance sought c. subscripts from the chemical formulas d. coefficients from the chemical equation

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20. What mass of oxygen at STP contains the same number of molecules as 14 grams of nitrogen? (Atomic masses: N = 14.0; O = 16.0) a. 14 b. 16

c. 28 d. 32

21. What is the number of atoms in a system that contains 1.0 mole of NH3?

a. 1 b. 4

c. 1 x 6.02 x 1023

d. 4 x 6.02 x 1023

22. For the reaction whose balanced equation is 4NH4(g) + 5O2(g) ---> 4NO2(g) + 6H2O(g)

what volume of O2 (g) at STP must react to produce 80.0 L of NO2 (g) at STP? a. 5.00 L

b. 20.0 L c. 80.0 L

d. 100 L

23. According to the reaction whose balanced equation is 2Na(s) + 2H2O(l) ---> 2NaOH(aq) + H2(g)

what volume of H2 (g) at STP is produced when 11.5 grams of Na(s) react? (Atomic masses: H = 1.0; O = 16.0; Na = 23.0) a. 5.60 L

b. 11.2 L c. 22.4 L

d. 44.8 L

24. When potassium chlorate, KClO3, is heated, it decomposes to produce potassium chloride and oxygen gas. How many grams of oxygen gas will be produced when 2.50 moles of potassium chlorate is decomposed? (Atomic masses: O = 16.0; Cl = 35.5; K = 39.1)

a. 120 g b. 60.0 g

c. 53.3 g d. 26.7 g

25. At STP, a sample of 35 mL of oxygen is mixed with 20 mL of hydrogen, and the mixture is ignited. What is the volume, at STP, of unreacted gas that remains?

a. 15 mL of oxygen b. 25mL of oxygen

c. 2.5 mL of hydrogen d. 10 mL of hydrogen

Grade 11 Chemistry

Matter & Gases

1. Which property of a sample of gas can be measured by use of a manometer? a. average kinetic energy b. pressure

c. density d. volume

2. On the average, the pressure of the atmosphere at sea level can support a column of mercury whose height is

a. 101.3 millimeters. b. 22.4 liters.

c. 760 millimeters. d. 14.7 meters.

3. Which causes the boiling temperature of a liquid to decrease?

a. increasing the pressure on the surface of the liquid b. decreasing the pressure on the surface of the liquid c. increasing the temperature of the liquid d. decreasing the temperature of the liquid

4. As the temperature of ethyl alcohol is raised from 30 oC to 40 oC, its vapor pressure a. increases. b. decreases.

c. remains the same.

5. The temperature at which the vapor pressure of a liquid reaches atmospheric pressure is the a. atmospheric temperature. b. condensation temperature.

c. vaporization temperature. d. boiling temperature.

6. A sample of water is boiling in an open vessel at 100 oC. As the atmospheric pressure decreases, the temperature at which the water will boil

a. increases. b. decreases.

c. remains the same.

7. Which term is applied to the spontaneous spreading of a substance from high concentration to low concentration?

a. evaporation b. diffusion

c. dissolving d. partial pressure

8. The volume of a confined gas can be decreased by the application of pressure at a constant temperature. Which property of gases accounts for this change in volume?

a. Gas molecules take up space b. Gas molecules are in constant motion c. Gas molecules are relatively far apart d. Gas molecules collide without loss of energy

Grade 11 Chemistry

9. Real gas behavior deviates from ideal gas behavior because the molecules of real gases a. never collide. b. attract one another. c. have no volume. d. are arranged in a regular geometric pattern.

10. Deviations from ideal gas behavior become most significant at a. high pressure and high temperature. b. high pressure and low temperature. c. low pressure and low temperature. d. low pressure and high temperature.

11. A 1.00 L sample of a gas at a pressure of 1.00 atmospheres is compressed to a volume of 0.25 L at constant temperature. What is the new pressure of the gas?

a. 0.50 atm b. 2.0 atm

c. 0.25 atm d. 4.0 atm

12. Which sample of gas has molecules with the highest average kinetic energy?

a. helium at 0 oC b. carbon dioxide at 20 oC

c. hydrogen chloride at 40 oC d. nitrogen at 60 oC

13. Which statement best explains why a confined gas exerts pressure?

a. The molecules are in random motion. b. The molecules travel in straight lines. c. The molecules attract each other. d. The molecules collide with the container walls.

14. What is the resulting volume when 750 mL of an ideal gas is warmed from 15 oC to 40 oC and the pressure is changed from 96 kPa to 101 kPa?

a. 65.8 mL

b. 775 mL c. 85.5 mL

d.1.91 x 103 mL

15. Which of the following contains factors that will all cause the volume of gas to increase?

a. Increasing the temperature and pressure, and decreasing the mass of the sample b. Increasing the temperature and mass of the sample, and decreasing the pressure c. Decreasing the temperature, and increasing the pressure and mass of the sample d. Increasing the temperature, the pressure, and the mass of the sample

Grade 11 Chemistry

16. A sample of gas has a volume of 1.50L at 106.4 kPa. What is its volume when the pressure is decreased to 65.9 kPa, assuming there is no change in temperature?

a. 0.455L b. 0.029L

c. 2.42L d. 1.08L

17. A sample of gas measures 125mL at 20 oC. What will be its volume when the temperature is changed to 0 oC at constant pressure?

a. 8.59 x 10-3mL b. 1.56 x 10-3mL

c. 134mL d. 116mL

18. How many moles of oxygen must be placed in a rigid container whose volume is 50.0L in order to produce a pressure of 4.50 atm at standard temperature?

a. 10.0 moles b. 0.496 mole

c. 1.79 moles d. 0.0996 mole

Solutions

1. A solution of sodium chloride is a

a. heterogeneous compound. b. homogeneous compound.

c. heterogeneous mixture. d. homogeneous mixture.

2. Which two liquids are completely miscible?

a. ether and water b. ethyl alcohol and water

c. benzene and water d. gasoline and water

3. The dissolved substance in a solution is called the

a. solvent. b. solute.

c. tincture. d. hydrate.

4. Which increases the rate of dissolving of a solid in a liquid?

a. increasing the pressure on the liquid b. decreasing the surface area of the solvent c. lowering the temperature of the liquid d. breaking the solid into small pieces

5. Ionic solids will dissolve to the greatest extent in liquid solvents that are a. polar. b. nonpolar.

c. amalgams. d. hydrated.

6. An alloy is an example of a

a. gaseous solution. b. liquid solution.

c. solid solution. d. miscible solution.

Grade 11 Chemistry

7. When the rate of dissolving of a substance is greater than the rate at which the substance is crystallizing out of solution, the solution is said to be a. saturated. b. unsaturated.

c. supersaturated. d. concentrated.

8. Substances whose water solutions conduct electricity are

a. volatile. b. electrolytes.

c. electrovalent. d. colligative.

9. Which is NOT true of solutions? a. Solutions are homogeneous mixtures. b. Dissolved particles in solutions are too small to be seen. c. Dissolved particles in solutions can be removed with filters. d. The components of a solution form a single phase.

10. What is the number of phases present in a solution of sodium chloride in water? a. one b. two

c. three d. four

11. Breaking a solid into small pieces increases the rate of dissolving because a. the concentration of the solution is decreased. b. the molecules of the solid move at a faster rate. c. more surface area of the solid is exposed to the liquid. d. the volume of the solid is decreased.

12. Approximately what volume of solution will be obtained when 40 g KNO is used to prepare a 2.0M solution?

a. 50 mL b. 200 mL

c. 400 mL d. 800 mL

13. Which of the following have universal solubility in water?

a. PO4-3

b. CO3-2

c. SO4-2

d. S-2

14. When magnesium acetate is dissolved in water how many acetate ions are in solution?

a. 1 b. 2

c. 3 d. 4

Grade 11 Chemistry

Organic Chemistry

1. Structural formulas have an advantage over molecular formulas because structural formulas show the

a. bonding capacity of each carbon atom b. geometric arrangement of each atom c. number of atoms of each element present d. percentage composition of the compound 2. Which series of hydrocarbons is characterized by a triple bond between a pair of adjacent

carbon atoms? a. the alkane series b. the alkene series

c. the alkyne series d. the alkadiene series

3. Which series of hydrocarbons has the general formula CnH2n ? a. the alkane series b. the alkene series

c. the alkyne series d. the alkadien4

4. The SECOND member of a particular homologous series has theformula C3H6. What is the formula for the FOURTH member of this series?

a. C4H8 b. C4H6

c. C5H10 d. C5H8

5. Which formula represents a member of the homologous series to which C8H14 belongs? a. C4H6 b. C4H7

c. C4H8 d. C4H10

6. Propyne is a member of the ___ family. a. alkane b. alkene

c. aromatic d. alkyne

Grade 11 Chemistry

Long Answer and Problem Solving Review Unit 1. a) What is the difference between an ionic bond and a covalent bond? b) Are the following pairs of atoms more likely to form ionic or covalent bonds? Give reasons for your answer.

i. chlorine and chlorine

ii. potassium and iodine

iii. carbon and oxygen

iv. magnesium and fluorine

2. Name the following compounds: a) CuCl b) MnO2 c) SO3 d) Ca(NO3)2

e) N2O4 f) NiF2 g) Na2S h) P2O5

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3. Draw a Lewis structure for each of the following: a) O2 b) CO2 c) CH4 d) NH3 e) HCN f) N2H4

a.

b. c.

d.

e. f.

4. Write the formula for each of the following compounds. a) aluminum chloride h) ammonium nitrate b) copper(II) sulphide i) sodium phosphide c) calcium hydroxide j) tin (II) bromide d) dichlorine heptaoxide k) iron(III) carbonate e) phosphorus pentachloride l) potassium oxide f) iron (II) iodide m) cobalt(III) sulphate g) carbon monoxide n) triphosphorus pentanitride

Grade 11 Chemistry

5. Classify each of the following reactions as synthesis, single displacement, double displacement, combustion or decomposition. a) iron + copper(I) nitrate iron(II) nitrate + copper b) phosphorus + oxygen diphosphorus pentoxide c) calcium carbonate calcium oxide + carbon dioxide d) propane + oxygen carbon dioxide + water e) lead(II) hydroxide lead(II) oxide + water f) ammonia + sulphuric acid ammonium sulphate g) potassium phosphate + magnesium chloride magnesium phosphate + potassium chloride 6. For each of the following, use a reactivity series to determine which single displacement reactions will proceed. For the reactions that do occur, predict the products. a) Cu (s) + HCl (aq)

b) Au (s) + ZnSO4 (aq)

c) Pb (s) + CuSO4 (aq)

d) Cl2 (s) + NaBr (aq)

e) Fe (s) + AgNO3 (aq)

Grade 11 Chemistry

7. Predict which is the more reactive in each pair of elements based on their position on the periodic table. Circle your answer in each pair. a) F & Cl b) Mg & Ca c) Na & Mg d) S & Cl e) Ne & F 8. Complete each equation. Indicate if no reaction occurs. a. C6H6 (g) + O2 (g)

b. Ni (s) + I2 (s)

c. CoBr2 (s)

d. Ca (s) + O2 (g)

e. Na (s) + H2O (l)

f. AlCl3 (aq) + Pb(NO3)2 (aq)

g. SO2 (g) + H2O (l)

h. MgO (s) + H2O (l)

i. Ca (s) + H2O (l)

j. H3PO4 (aq) + Al(OH)3 (aq)

k. BaCO3 (s)

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9. Define following terms: a) mass number b) atomic mass

The Mole and Stoichiometry Show all Givens, Conversion Factors and Work! 1. Define the term “mole”. 2. A sample of glucose (C6H12O6) has a mass of 36.2 g. a) How many moles of glucose molecules are present? b) How many molecules are there? c) How many atoms of oxygen are there?

Grade 11 Chemistry

3. Balance the following equations: a) NH3(g) + O2 (g) NO (g) + H2O(l) b) NO2(g) + H2O(l) HNO3(aq) + NO(g)

c) C12H22O11(s) + O2(g) CO2(g) + H2O(g)

d) KClO3(s) KCl(s) + O2(g)

e) MnO2(s) + HCl(aq) MnCl2(aq) + Cl2(g) + H2O(g)

f) Al2O3(s) Al(s) + O2(g)

g) KOH(aq) + H3PO4(aq) K3PO4(aq) + H2O(aq)

4. Sodium metal reacts with chlorine gas to produce sodium chloride. 2 Na(s) + Cl2(g) 2 NaCl a) What mass of sodium is needed to completely react with 15.00 g of chlorine gas?

Grade 11 Chemistry

b) What mass of sodium is required to produce, in excess chlorine, 8.00 g of magnesium chloride? 5. When a solution containing 15.0 g of aluminum chloride is mixed with a solution containing 15.0g of sodium hydroxide the following reaction occurs. AlCl3(aq) + 3 NaOH(aq) 3 NaCl(aq) + Al(OH)3(aq)

a) What is the limiting reactant? b) Calculate the mass of aluminum hydroxide produced.

Grade 11 Chemistry

6. When 8.40 g of zinc metal is placed in a solution in which 11.6 g of HCl is dissolved, hydrogen gas and zinc chloride are produced. Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g)

a) Identify the limiting reactant. b) If 0.19 g of hydrogen gas is produced, what is the percentage yield?

Grade 11 Chemistry

7. Given 6.0 mol of H2 gas and 1.5 mol of N2 gas according to the following reaction, calculate the volume of NH3 gas produced at S.T.P. 3H2 + N2 → 2NH3

8. Given 5.0 mol of sulphur and 8.4 mol of oxygen gas, calculate the mass of SO3 gas produced. 2S + 3O2 → 2SO3 9. Assuming the following reaction is 74.2 % efficient, how much nitrogen trioxide will be formed when 50 kg of nitrogen dioxide reacts with an excess of oxygen? 2NO2 + O2 ⟶ 2NO3 10. Consider the decomposition of boron mononitride into elemental boron and nitrogen gas. How much boron mononitride is needed to produce 5 g of boron if the reaction’s percent yield is 9.5% ?

Grade 11 Chemistry

SOLUTIONS AND SOLUBILITY Show all Givens, Equations and Work! 1. Water is known as a universal solvent. a) What is meant by this term? b) Describe the structure and polarity of water. c) Explain why water is able to dissolve a large number of substances. d) Explain why water is not able to dissolve oil

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2. As temperature increases, solubility in water will increase/decrease for a) solids b) liquids c) gases 3. Refer to the solubility curve below when answering the following questions.

a) Which salt is least soluble at 40 C? b) Which three compounds are gases? c) How much potassium nitrate can you dissolve in

100g of water at 70C? d) How much sodium nitrate can you dissolve in

200g of water at 55C? e) A solution containing 50g of NaCl in 100g of

water at 100C is (circle one) saturated/supersaturated/unsaturated f) A solution containing 40g of NaCl in 100g of

water at 100C is (circle one) saturated/supersaturated/unsaturated g) A solution containing 10g of NaCl in 100g of

water at 100C is (circle one) saturated/supersaturated/unsaturated

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4. Which pair of solutions will produce a precipitate when mixed? Circle your answers.

a) NaOH(aq) + H2S (aq)

b) CaCl2 (aq) + K2CO3 (aq)

c) Al(NO3)3 (aq) + Na2SO4 (aq)

d) CuSO4 (aq) + NH4Cl (aq)

5. Write the net ionic equation for each of the following: a) BaCl2(aq) + Na2SO4(aq) BaSO4(s) + 2 NaCl(aq)

b) CuSO4(aq) + 2 AgNO3(aq) Ag2SO4(s) + Cu(NO3)2(aq)

c) Pb(NO3)2(aq) + 2 KI(aq) PbI2(s) + 2 KNO3(aq)

Grade 11 Chemistry

6. Calculate the molar concentration (M) of each of the following solutions: a) 0.174 mol of sodium hydroxide dissolved in water to a final volume of 0.250 L of solution. b) 60.0 g of NaOH dissolved in water to a final volume of 750.0 mL of solution. 7. What mass of sodium carbonate is required to make 0.500 L of a 0.12 M solution? 8. Calculate the mass of solute in 24.9 mL of a 0.200 M solution of NaOH(aq).

Grade 11 Chemistry

9. What volume of concentrated 17.8 M stock solution of sulphuric acid would you need in order to prepare 2.0 L of 0.215 M sulphuric acid (H2SO4)? 10. A 15.0 mL sample of 11.6 M HCl(aq) is added to water to make a final volume of 500.0 mL. Calculate the concentration of the final dilute HCl(aq). 11. Explain the difference between an unsaturated, a saturated and a supersaturated solution. 12. What three things can generally be done to increase the amount of solid that will dissolve in a solvent?

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13. Write the equation which represents the following substances being dissolved in water. Determine the concentration of all aqueous species if 2 mol of solute were dissolved in 1 L of water. a) sodium hydroxide b) magnesium nitrate c) ethanol d) ammonium sulfide e) calcium carbonate f) potassium sulfide GASES 1. Describe the quantitative relationships which exist among the following variables for an ideal

gas.

a. Pressure & Volume

b. Temperature & Volume

c. Pressure & Temperature

d. Pressure & moles

e. Volume & moles

2. Explain Charles’ Law at the particle level.

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3. Explain Boyle’s Law at the particle level.

4. Explain Gay-Lussac’s Law at the particle level.

5. What is an ideal gas? 6. A balloon filled to 2.00 L at 98.0 kPa is taken to an altitude at which the pressure is 82.0 kPa,

the temperature remaining the same. What is the new volume of the balloon? 7. A sample of gas in a metal cylinder has a pressure of 135.0 kPa at 250C. What is the pressure in

the cylinder if the gas is heated to a temperature of 1250C?

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8. A sample of gas occupies 1.00 L at 220C and has a pressure of 700.0 kPa. What volume would this gas occupy at STP?

9. Calculate the volume occupied by 2.50 moles of nitrogen gas at 58.6 kPa and -40.00C.

10. What mass of chlorine gas is present in a sample that has a volume of 500.0 mL at 200C and exerts a pressure of 450.0 kPa?

11. 1.00 L of an unknown gas has a mass of 1.25 g at STP. Calculate the molar mass of the gas.

Grade 11 Chemistry

12. In the following reaction, what mass of zinc is necessary to produce 250.0 mL of hydrogen gas at STP?

Zn(s) + 2 HCl(aq) H2(g) + ZnCl2(aq)

13. When ammonium nitrite undergoes decomposition, only gases are produced according to the equation:

NH4NO2(s) N2(g) + 2H2O(g)

What is the volume of water vapour produced at 819K and 1.00 atm pressure when 128 g of ammonium nitrite undergoes the above decomposition reaction?

14. A 15L container of mixed gas has a pressure of 2.5atm. If you add 36.74g of N2O2 at a constant temperature of 20.0oC, what will the final pressure of the combined gases in the container be?

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15. The synthesis of ammonia gas is as follows: 3 H2 + N2 → 2 NH3

What is the total volume of hydrogen and nitrogen gas that is reacted if 76.28g of NH3 is produced at a temperature of 563K and 1.00atm of pressure? 16. A 16 L container of nitrogen gas is stored in a location where its pressure has been reduced to a third of its original value and its Kelvin temperature has been halved. What is the new volume of the nitrogen gas? MATTER Use the following graph to answer questions 1 – 7.

-100

-50

0

50

100

150

200

Tem

pe

ratu

re (°C)

Energy

A

B

C

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1. What are the boiling points of all three substances? 2. What are the melting points of all three substances? 3. Which of the three substances most likely has the strongest intermolecular forces? Why? 4. In what temperature range is substance A in the liquid state? 5. In what temperature range is substance B in the solid state? 6. Why are there plateau in the graphs? What do these sections represent, and why does the temperature not change during these sections? 7. How would the heating curves change if the mass of the samples were increased?

Grade 11 Chemistry

Use the following graph for questions 8 - 10

8. What is the approximate boiling point of A under standard conditions? 9. What must the approximate pressure be for C to boil at 80°C? 10. Which of the three substances most likely has the weakest intermolecular forces? 11. At what point does a substance boil? How does elevation affect this point? 12. Why does a balloon remain inflated over time?

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13. How are intermolecular forces related to phase changes? 14. What are the difference between the covalent bonds and the intermolecular forces in a molecular compound? Organic Chemistry

1. Compare and Contrast Alkanes, Alkenes and Alkynes in terms of:

Alkanes Alkenes Alkynes

Bonding

General Formula

Structural Formula

Reactivity

Uses

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2. Name the following:

A

CH3

C

H2

CH

2

CH3

B

CH3

C

H2

CH

2

CH CH

3

CH3

C

CH3

C

H2

CH

2

CH CH

3

CH3

D

CH3

CH CH

3

CH3

E

CH2

CH

CH

2

CH3

F

CH3

CH

CH

CH3

G

CH3

CH C

H CH C

H

CH2

CH3

CH3

CH3

H

CH

CH CH2

CHC

HCH

3

CH3

CH3CH

3

3. Draw and name all the structural isomers for the formula C5H10

Grade 11 Chemistry

4. Name the following:

A

CH

C

CH

2

CH3

B

CH3

CH C

H C

C

CH3

CH3

CH3

C

CH

C

C

H2

CH C

H2

CH

2

C

H2

CH3

CH2

CH3

D

CH

C

CH C

H CH C

H CH3

CH2

CH3

CH3

CH3

CH3

E

CH3

C

H2

CH

2

C

H2

CH CH

3

OH

F

CH3

CH C

HOH

CH3

CH3

G

CH3

CH C

H C

H2

CH3

OH

CH3

H

CH3

CH C

H OH

CHC

H2

CH3

CH3

CH3

I

CH3

C

H2

CH C

H2

CH

2

C

H2

CH C

H CH CH

3

CH3

CH3

OH

CH3

J

CH

2

C

H2

CH

OH

CH3

C

H2

CH3

Grade 11 Chemistry

K

CH3

C

H2

C

O

OH

L

CH

2

C

H2

C

O

OH

C

H2

CH

2

C

H2

CH

2

C

H2

CH3

M

CH

2

C

H2

C

O

OH

C

H2

CH

2

CH3

N

CH

2

C

H2

C

O

OH

C

H2

CH3

5. Draw each of the following structures and give the name of the group to which each belongs.

a. 3-ethyl -4-methyloctane b. 3,4-diethyl-2 hexene

c. 3-ethylhexane d. pentyl-butanoate

e. butanoic acid f. 2-propanol g. 2,4-dimethyl-3 heptanol h. ethyl propanoate

i. ethyl ethanoate

Grade 11 Chemistry

Grade 11 Chemistry

Grade 11 Chemistry

Grade 11 Chemistry