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Review of Chemistry 20 817NEL

Unit 1 Chemical Bonding(Chapter 3)

1. Distinguish between the two important types ofscientific knowledge.

2. Identify the characteristics of acceptable scientifictheories.

3. Explain the octet rule and how it relates to chemicalreactivity.

4. Copy and complete the following table.

5. (a) State the types of elements expected to react toform compounds containing covalent bonds.

(b) State the types of elements expected to react toform compounds containing ionic bonds.

(c) Explain your answers to (a) and (b) using theconcept of electronegativity.

(d) Why is it difficult to predict the type of bondingin some compounds using onlyelectronegativities?

6. The two major types of compounds are ionic andmolecular.

(a) Compare the naming of these compounds.

(b) Given the name of an example of eachcompound, outline how the chemical formula isobtained. Use specific examples in your answer.

7. Theories are created to explain observations. For eachof the following properties of ionic compounds, writea brief theoretical explanation.

(a) Ionic compounds are hard solids with highmelting and boiling points.

(b) Ionic compounds are electrical conductors inmolten and aqueous states.

8. Using Lewis symbols and formulas, write theformation equation for each of the followingcompounds.

(a) potassium bromide

(b) sodium oxide

(c) calcium fluoride

9. Why are chemical formulas for ionic compoundsalways based the simplest whole number ratio ofions? Is the simplest whole number ratio also used formolecular formulas? Why or why not?

10. Compare ionic and covalent bonds, including howthey are formed, according to theory and the natureof the bond.

11. For each of the following molecular formulas, drawthe Lewis, structural, and stereochemical formulas,and state the shape around the central atom.

(a) OCl2 (d) HCN

(b) SiH4 (e) CH2O

(c) NCl312. Classify each of the molecules represented in the

previous question as polar or nonpolar. Justify youranswer using the molecular shape and bond dipoles(charge distributions).

13. Methylisocyanate is a toxic pesticide that ismanufactured using the following chemical reaction.

CS2 � CH3NH2 → CH3NCS � H2SRewrite this chemical equation using structuralformulas for all reactants and products.

14. Define the three types of intermolecular forces. Foreach type of force, state how you would know if thistype of force is likely present among molecules of asubstance.

15. Each of the following four substances is either aliquid at SATP or converted to a liquid by changingthe conditions: C3H7F, C3H5(OH)3, C3H7NH2, C3H8(a) Construct a table to summarize the types of

intermolecular forces believed to be presentamong molecules of each of these substances.

(b) Predict the order of boiling points from lowest tohighest. Justify your answer.

Your review of Chemistry 20 for Chemistry 30 will be more successful if you study the highlighted Summaries, Sample Problems,and Communication Examples in each chapter. By answering the following questions you will find out where you need to check yourunderstanding before starting Chemistry 30.

Appendix G REVIEW OF CHEMISTRY 20

G

Table 1 Theoretical Descriptions of Selected Elements

Element Lewis Group Number Number Numbername symbol number of valence of lone of bonding

electrons pairs electrons

calcium

aluminium

arsenic

oxygen

bromine

neon

Appendix G-I_Chem20 12/21/06 10:06 AM Page 817

818 Appendix G NEL

16. Why are boiling points often used as an indirectmeasure of the strength of intermolecular forcesamong molecules of a substance?

17. Explain each of the following observations in termsof the characteristics of molecules and intermolecularforces.

(a) The boiling point of fluorine is significantly lessthan that of chlorine.

(b) Drops of ethanol are attracted to an electricallycharged strip.

(c) Ice has a regular hexagonal structure.

18. A simple, but useful, distinction that is often made isto classify the water on Earth as either fresh water (asin most lakes and streams) or salt water (as in theoceans).

(a) Contrast these two terms from a scientificperspective.

(b) How is this distinction useful from atechnological perspective?

19. Describe an example in which scientific research ledto the development of a new technology.

Unit 2 Gases (Chapter 4)1. List seven ways by which empirical knowledge is

communicated.

2. List the three characteristics of acceptable scientificlaws and generalizations.

3. Describe one natural phenomenon and onetechnological product that each depend on theproperties of gases.

4. Complete the following statements.

(a) At a constant temperature and chemical amountof gas, as the pressure increases, the volume________.

(b) At a constant pressure and chemical amount ofgas, as the temperature decreases, the volume________.

(c) At a constant volume and temperature, if thechemical amount of gas inside a container isincreased, the pressure ________.

5. Choose one of the statements in question 4 and writea general design for an experiment to test thestatement. Include the identification of all importantvariables.

6. For each statement in question 4, sketch a graph ofthe relationship between the manipulated andresponding variables.

7. Convert 95.8 kPa into units of millimetres of mercuryand atmospheres.

8. A 1.5 L volume of gas is compressed at a constanttemperature from 1.0 atm to 5.0 atm. Calculate thefinal volume.

9. A balloon can hold 800 mL of air before breaking. Aballoon at 4.0 °C containing 750 mL of air is allowedto warm up. Assuming a constant pressure inside theballoon, determine the minimum Celsiustemperature when the balloon breaks.

10. A sample of argon gas at 101 kPa and 22.0 °Coccupies a volume of 150 mL. If the volume doublesat a temperature of 150 °C, determine the newpressure.

11. Using the kinetic molecular theory, explain Boyle’sand Charles’ laws.

12. Illustrate the law of combining volumes using asimple example. Describe the theory used to explainthis law.

13. Many people use propane barbeques for outdoorcooking. Predict the volume of carbon dioxideproduced when 15 L of propane completely burns atSATP.

14. Describe and compare the behaviour of real and idealgases using the kinetic molecular theory.

15. Predict the volume that 25.0 g of oxygen gas wouldoccupy at 22.0 °C and 98.1 kPa.

16. Compare the volume that 0.278 mol of hydrogenwould occupy at STP and SATP.

17. An average bungalow requires about 400 kmol ofmethane per year for space heating.

(a) Determine the volume of methane at SATP usedin one year.

(b) Predict the volume of methane used if thepressure is 98.5 kPa and the temperature is 12.7 °C.



Unit 3 Solutions, Acidsand Bases (Chapters 5 & 6)

1. For each of the following perspectives write a briefstatement describing the focus or concern of thatpoint of view.

• scientific

• technological

• economic

• ecological

• political

2. List three topics that are current STS issues.

3. Classify each of the following statements using one ofthe issue perspectives listed in question 1. All of thestatements concern sulfur dioxide emissions.

(a) An industry spokesman reported that emissionsof sulfur dioxide were within the limits set byenvironmental legislation.

(b) Laboratory research has provided evidence thatsulfur dioxide from the combustion of fossil fuelsis converted to sulfur trioxide in the presence ofoxygen.

(c) The cost of ending sulfur dioxide pollution of theatmosphere will be high. The longer we delayfacing the problem, the greater will be the cost.

(d) Sulfur oxides and their related dissolved acids areparticularly damaging to soil microbes, water lifeforms, plants, building materials, and people.

(e) One of the most promising scrubbers to removesulfur dioxide gas from a smoke stack is thelimestone−dolomite process.

4. Compare the goals of science and technology.

5. Describe a homogeneous mixture and provide severalexamples.

6. Define the two main parts of a solution. State anexample using a chemical formula and identity thetwo parts in words.

7. In the exploration of outer space, scientists usuallylook for the presence of water as a strong indicationof the existence of living things. Briefly explain thisstatement in terms of solutions and reactions.

8. List at least six examples of manufactured solutionsfound in the home and six examples of naturalsolutions found in the environment.

Appendix G

9. Distinguish between electrolytes and non-electrolytes,including examples of types of substances in eachcategory.

10. Explain, in terms of breaking and forming bonds,why the dissolving of substances in water can beeither exothermic or endothermic.

11. Compounds may be ionic or molecular and may alsobe acids, bases, or neutral compounds.

(a) Design an experiment to classify the solute ineach of a number of different solutions.

(b) Outline the expected results.

12. Write dissociation or ionization equations for thefollowing pure substances dissolving in water.

(a) lithium phosphate solid

(b) hydrogen chloride gas

(c) aluminium sulfate solid

13. For each of the following pure substances, write theformulas for the entities present when each substanceis placed in water.

(a) Sr(OH)2(s) (d) CH3COOH(l)

(b) HNO3(l) (e) AgCl(s)

(c) C3H8(g) (f) CH3OH(l)

14. List the three advantages of solutions fortechnological applications.

15. Suppose you are given four unlabelled beakers, eachcontaining a colourless aqueous solution of onesolute. The possible solutions are NaCl(aq), HCl(aq),BaCl2(aq), and CH3Cl(aq). Write a series ofdiagnostic tests to distinguish each solution from theothers.

16. Compare the ways in which solution concentrationsare expressed in chemistry labs, consumer products,and environmental studies.

17. A household cleaner contains 12.5 g of sodiumhypochlorite in 500 mL of solution. Determine thepercentage mass by volume concentration of thissolution.

18. A drain cleaner contains 2.75 mol/L sodiumhydroxide. Calculate the volume of solution thatcontains 0.375 mol of sodium hydroxide.

19. A windshield washer solution was prepared bydissolving 100 g of methanol in water to form 2.00 Lof solution. Calculate the amount concentration ofthe solution.

G


820 Appendix G NEL

20. A 0.251 mol/L calcium chloride solution is requiredfor an experiment.

(a) Calculate the mass of calcium chloride that needsto be measured.

(b) Write a specific procedure for an untrainedlaboratory technician to prepare this solution.

21. (a) Predict the volume of concentrated, 14.6 mol/Lphosphoric acid required to prepare 250 mL of a0.375 mol/L solution.

(b) Write a specific procedure to prepare thissolution.

22. Calculate the amount concentration of each ion in a2.1 mol/L solution of iron(III) chloride?

23. How does the solubility of solids and gases change asthe temperature increases?

24. Excess copper(II) sulfate is added to water in a closedsystem until no more solute dissolves at a constanttemperature.

(a) Describe some empirical properties of thismixture.

(b) Provide a brief theoretical explanation of theseproperties.

25. Write the acid formula for each of the followingsubstances.

(a) aqueous hydrogen bromide

(b) aqueous hydrogen sulfite

(c) hydrofluoric acid

(d) sulfuric acid

26. Copy and complete the following table.

27. The pH of pure water is 7, of carbonated water about5, and of a cola drink about 3. How many times moreacidic is a cola drink than carbonated water and purewater?

28. Use the modified Arrhenius theory to write chemicalequations explaining the following evidence.

(a) A vinegar solution is acidic.

(b) A baking soda (sodium hydrogen carbonate)solution has a pH of 8.

(c) Some muriatic (hydrochloric) acid is neutralizedwith a lye (sodium hydroxide) solution.

29. A simple window cleaning solution containing 0.25 mol/L ammonia has a pOH of 2.5.

(a) Convert the pOH into an amount concentrationof hydroxide ions.

(b) Write a balanced chemical equation to explainthis basic solution.

(c) Is ammonia a strong or weak base? Justify youranswer.

30. Write a design for an experiment to identify strongand weak acids. Include three different diagnostictests and identify important controlled variables.

31. Polyprotic acids and bases occur naturally and aremanufactured for a variety of purposes.

(a) Distinguish between monoprotic and polyproticacids and bases.

(b) Using boric acid (aqueous hydrogen borate) as anexample, write a series of chemical equationsshowing successive reactions with water.

32. Most scientists agree that the increasing emission ofcarbon dioxide into the atmosphere from the burningof fossil fuels is the prime cause of global warming.This problem might be even worse if it were not forthe fact that approximately half of the carbon dioxideproduced is absorbed by the world’s oceans. However,recent research has shown that this is making theoceans more acidic—about 30% more acidic over thepast two hundred years.

(a) Use the modified Arrhenius theory to write achemical equation explaining the increasedacidity of the world’s oceans.

(b) Scientists are not certain what effect the increasedacidity will have. If we assume there will be aproblem in the oceans, describe some solutions toreduce the addition of carbon dioxide to theoceans.

33. Using pesticides as an example, summarize theintended and unintended consequences of thischemical technology.

Table 2 Hydroxide Concentrations and pHs

[H3O+(aq)] (mol/L) pH Acidic/basic/neutral

1.0 � 10�7

8

3.7

6.23 � 10�9



Unit 4 QuantitativeRelationships (Chapters 7 & 8)

1. Compare the fields of chemistry and chemicaltechnology.

2. Describe two examples of chemical technologies,used by consumers, that are based on thestoichiometry of chemical reactions.

3. Distinguish between qualitative and quantitativechemical analysis and provide an example of eachtype of analysis.

4. For each of the following mixtures, write a balancednet ionic equation and identify all spectator ions. Allreactant solutions are assumed to be at least 0.10 mol/L in concentration.

(a) sodium hydroxide and cobalt(II) chloridesolutions

(b) silver nitrate and calcium iodide solutions

(c) silver nitrate solution and zinc metal

(d) hydrochloric acid and solid calcium hydroxide

(e) the precipitation of aluminium hydroxide inqualitative analysis

5. In your own words, describe the meaning ofstoichiometry.

6. List the three types of stoichiometry and describehow each type is recognized.

7. In general, how do chemical industries use theprinciples of stoichiometry to maximize yields andminimize waste?

8. In the steel industry, carbon reacts with iron(III)oxide (from iron ore) to produce molten iron andcarbon dioxide.

(a) Write a complete balanced chemical equation forthis reaction.

(b) Translate this chemical equation into an Englishsentence including all chemical amounts andstates of matter.

(c) Using the coefficients, calculate the mass of eachreactant and product in this balanced chemicalequation.

(d) How does the total mass of reactants comparewith the total mass of products? What principledoes this illustrate?

Appendix G

9. Predict the mass of lead(II) iodide precipitate thatforms when 2.93 g of potassium iodide in solutionreacts with excess lead(II) nitrate.

10. In a hard water analysis, a calcium chloride solution isreacted with excess aqueous sodium oxalate toproduce 0.452 g of calcium oxalate precipitate.Determine the mass of calcium chloride present inthe original solution.

11. Analysis for sulfate ions is usually done by firstprecipitating barium sulfate from a sample. The filterpaper containing the barium sulfate precipitate isthen ignited. Carbon from the burnt filter paper thenreacts with the barium sulfate as shown in thebalanced chemical equation below.

BaSO4(s) + 2 C(s) → BaS(s) + 2 CO2(g)(a) Predict the mass of carbon required to react with

1.50 g of barium sulfate precipitate.

(b) List the assumptions you have made in thiscalculation.

12. In a test of the stoichiometric method, an excess ofsodium hydroxide solution is reacted with a solutioncontaining 1.50 g of aluminium sulfate.

(a) Predict the mass of precipitate expected in thisreaction.

(b) If the actual yield in this experiment was 0.96 g ofprecipitate, calculate the percent yield.

(c) Outline at least three possible reasons for thediscrepancy between the theoretical (predicted)yield and the actual yield.

13. Powdered aluminium metal is one of the fuels used inthe solid rocket boosters for the NASA Space Shuttle.What volume of oxygen at SATP is required to reactcompletely with 100 kg of aluminium?

14. A portable hydrogen generator uses the reaction ofsolid calcium hydride and water to form calciumhydroxide and hydrogen. Determine the volume ofhydrogen at 96.5 kPa and 22 °C that can be producedfrom a 50 g cartridge of CaH2(s).

15. A volumetric analysis shows that it takes 32.0 mL of2.12 mol/L NaOH(aq) to completely react with 10.0 mL of sulfuric acid from a car battery. Calculatethe amount concentration of sulfuric acid in thebattery solution.

G


822 Appendix G NEL

16. In a laboratory, silver metal can be recycled toproduce silver nitrate by the following reaction.

3 Ag(s) � 4 HNO3(aq) →3 AgNO3(aq) � NO(g) � 2 H2O(l)

Predict the volume of 15.4 mol/L nitric acid requiredto react with 1.68 kg of silver metal.

17. Distinguish between limiting and excess reagents.

18. Describe the purpose of using an excess reagent in aquantitative analysis?

19. Calcium carbonate is commonly used in simpleantacid products to counteract acidity in thestomach. Suppose you add a 750 mg tablet of calciumcarbonate to 200 mL of 0.10 mol/L hydrochloric acid(representing the stomach acid).

(a) Which reactant is in excess and by how much?Give your answer in moles.

(b) Predict the mass of calcium chloride formed inthis reaction.

20. Complete the Materials and Analysis of the followinglab report.

Problem What is the amount concentration of an unknownsodium carbonate solution?

Design Samples of sodium carbonate solution were titratedwith a standardized hydrochloric acid solution usingmethyl orange as the indicator.

Evidence

22. Titration curves are useful in studying the progress ofa reaction, such as an acid−base reaction.(a) Sketch a general curve for the titration of a strong

base with a strong acid. Label the axes andprovide a title for the graph. No numbers arerequired.

(b) Place an “X” on the curve where the reaction iscomplete. At what pH should this occur?

(c) Identify a suitable indicator for any strong base−strong acid titration and justify your answer.

(d) Would your answers to (a), (b), and (c) change ifa strong acid were titrated with a strong base?Note any differences.

Table 3 Titration of 25.0 mL Samples of Na2CO3(aq) with 0.352 mol/L HCl(aq)

Trial 1 2 3 4

Final burette reading (mL) 16.5 31.8 47.0 16.4

Initial burette reading (mL) 0.6 16.5 31.8 1.2


Appendix G Review of Chemistry 20 Answer Key Appendix G Review of Chemistry 20 Answer Key Unit 1: Chemical Bonding (Chapter 3) 1. The two types of scientific knowledge are Empirical and Theoretical. Empirical

knowledge is based on observations and can always be identified using our senses or instruments. These are observations made during investigation. Theoretical knowledge is not observable. It attempts to explain the observations in terms of ideas.

2. An acceptable scientific theory is one that logically describes, explains, and can

predict observations. If a theory is tested and the investigation results in data that agree with the prediction (within reasonable limits of experimental error) and can be replicated, the theory is acceptable.

3. The octet rules states that a given atom or ion (above energy level 1) may have a

maximum of eight (8) valence electrons (arranged in 4 pairs). This arrangement of electrons represents stability for the atom.

4. Table 1 Theoretical Descriptions of Selected Elements

Element Name

Lewis Symbol

Group number

Number of valence

electrons

Number of lone pairs

Number of bonding electrons

calcium 2 2 0 2 aluminium

13 3 0 3

arsenic

15 5 1 3

oxygen

16 6 2 2

bromine Br 17 7 3 1

neon

18 8 4 0

5. (a) Covalent bonds form between two nonmetallic atoms in a compound. (b) Ionic bonds form between metallic ions and nonmetallic ions in a compound. (c) Electronegativity measures the relative ability of an atom to attract a pair of

bonding electrons. When two nonmetallic atoms combine, each atom has roughly the same “pull” on these electrons. Therefore the electrons are shared between the bonded atoms. When a metal and nonmetal combine there tends to be a large difference in the strength of the attraction on these electrons. (The nonmetal attracts the electrons more strongly than the metal) This results in the electron(s) being transferred from one atom to the other, which results in the formation of oppositely charged ions.

(d) There are a number of exceptions to the rules above. Some metal/nonmetal combinations may have very similar electronegativity values, and yet form ionic bonds. It should be noted that real bonds have both ionic and covalent characteristics.

Nelson Chemistry Alberta 20—30 www.science.nelson.com © 2009 Nelson Education 1

Appendix G Review of Chemistry 20 Answer Key 6. (a) An ionic compound always consists of anions and cations. When naming these

compounds, the name of the cation is always provided first followed by the name of the anion. Example: sodium chloride If the cation is multi-valent, the ion charge being used is converted to a Roman numeral and inserted, in brackets, after the cation name, e.g., copper(I) chloride. When naming binary molecular compounds, Greek prefixes are used to indicate the number of each kind of atom present in the molecule, e.g., carbon monoxide.

(b) Naming ionic Compounds:

(i) Write the chemical symbol with the ion charge for each ion. (ii) Predict the simplest whole number ratio to obtain a net charge of zero. (iii) Write the cation first, followed by the anion. Indicate the number of each ion required with a subscript. Note: if more than one of a polyatomic ion is needed, put the ion in brackets with the subscript outside the bracket. Example A: calcium chloride (i) Ca2+Cl− (ii) We would need two of the Cl− ions to balance the positive charge of the Ca2+. (iii) Final formula: CaCl2 Example B: magnesium nitrate (i) Mg2+NO3− (ii) We need two of the polyatomic nitrate ions, NO3−, to balance the charge of the magnesium ion,Mg2+. (iii) Final formula: Mg(NO3)2 Naming Molecular Compounds: (i) Write out each element symbol.

(ii) After each element symbol write a subscript indicating the number of atoms of each present (unless only one atom is required, in which case no subscript is written). The subscripts are derived from each prefix in the name. Note: the “mono” prefix for the first element is normally omitted therefore if no prefix is given, it is assumed that there is only a single atom of that element. Example C: dinitrogen tetroxide (i) N and O

(ii) “di” indicates two nitrogen atoms: N2. “tetr” indicates four oxygen atoms: O4. Final formula: N2O4 Example D: sulfur trioxide (i) S and O

(ii) No prefix with sulfur indicates one sulfur atom: S. ”tri” indicates three oxygen atoms: O3. Final formula: SO3


Appendix G Review of Chemistry 20 Answer Key 7. (a) Ionic compounds are hard solids because oppositely charged ions are strongly

attracted to each other and form a rigid crystal structure. Ionic compounds have high melting and boiling points because the bonds among the ions are very strong, so a lot of energy must be provided to overcome them.

(b) Molten and aqueous ionic compounds conduct electricity because the individual

ions have separated. This results in the presence of charged particles that can move freely and conduct a current.

+ Br K+ Br8. (a) K

(b) (c) Ca + F2 F Ca2+ F

Ca2+ F2

or 9. In any sample of matter containing ionic bonds, the ratio of cations to anions is

always the same. In sodium chloride for example, there is always one sodium ion for every chloride ion. They do not form unique pairs, however. Rather, each ion bonds with several other oppositely charged ions to form a crystal lattice arrangement. We do not know exactly how many of each ion are present in a crystal, so we simply state the ratio of ions. In molecular compounds, the ratio of one atom to another may vary from compound to compound. Methane (a carbon–hydrogen compound) contains four hydrogen atoms for each carbon atom present. They bond in discrete groups (molecules) of one carbon atom and four hydrogen atoms. Ethane (another carbon–hydrogen compound) is made up of discrete molecules of three hydrogen atoms and one carbon atom. Because we know exactly how many of each atom are present in a molecule, we can specify the actual atom numbers.

10. When ionic bonds form, one or more electrons are completely transferred from one

atom to another resulting in the formation of oppositely charged ions. These ions are attracted to each other and to other nearby ions. It is this attraction that forms the basis of the ionic bond. When covalent bonds are formed, one or more electrons are shared between two bonded atoms. These electrons are attracted to both nuclei at the same time.


Appendix G Review of Chemistry 20 Answer Key 11.

Molecular formula

Lewis formula Structural formula

Stereochemical formula

Molecular shape

(a) OCl2

angular

(b) SiH4

tetrahedral

(c) NCl3

trigonal pyramidal

(d) HCN linear (e) CH2O

trigonal planar

12. (a) OCL2 polar (b) SiH4 nonpolar (c) NCl3 polar (d) HCN polar (e) CH2O polar 13.

+ C

H

H

NH

H

H

CH

H

N

H

C S +S

H HS C S

14. There are three types of intermolecular forces: London forces, dipole−dipole forces,

and hydrogen bonding. London forces are found in all samples of molecular matter. They result from the attraction between the electrons of one molecule to the nuclei of another. Generally, the higher the total number of electrons, the stronger the London forces (and therefore the higher the melting point and boiling point). All compounds experience London forces, but the forces are largest in molecules with large numbers of electrons. Dipole−dipole forces exist in substances that have polar molecules. They result from the attraction between the positive end of one polar molecule to the negative end of another. Hydrogen bonding is a special, stronger kind of dipole-dipole force. It occurs between molecules that contain a hydrogen atom bonded to a fluorine, an oxygen, or


Appendix G Review of Chemistry 20 Answer Key

a nitrogen atom. The hydrogen end of the molecules is slightly positively charged and the other end is slightly negatively charged. The positive and negative ends of nearby molecules are attracted to each other and form bonds.

15.

Compound

(a) Intermolecular forces present

(b) Relative melting point

(lowest to highest)

Rationale

C3H7F

London dipole−dipole

2 In addition to London forces this is a polar molecule so dipole−dipole forces are present.

C3H5(OH)3

London hydrogen bonding

4 (highest)

Has multiple hydrogen bonding sites on the molecule and has more electrons so greater London forces are present.

C3H7NH2

London hydrogen bonding

3 Has hydrogen bonding but only on one site within the molecule.

C3H8

London 1 (lowest)

Has London forces only (not a polar molecule).

16. In order to evaporate a substance we must overcome the intermolecular forces that are

holding the molecules together (resulting in the liquid state). The quantity of energy required to overcome the forces and boil a substance must therefore indicate the relative strength of these forces of attraction.

17. (a) Fluorine and chloride are both made up of nonpolar molecules. Therefore,

London forces are the only intermolecular forces present. Since chlorine is a somewhat larger molecule containing more electrons (34, compared to 18), chlorine has more of these attractions to break. This results in chlorine’s higher melting point.

(b) Ethanol is a very polar molecule. When ethanol is brought close to a charged object, the molecules orient themselves so that the ends of the molecules closest to the strip have the charge opposite to the charge on the strip. There is then an overall force of attraction between the strip and the ethanol molecules..

(c) The water molecules in ice are held together because of their hydrogen bondingand an angular shape. This shape causes the molecules to line up in specific positions as the hydrogen bonds form between the molecules. The relative positions of the water molecules place them in a hexagonal structure.

18. (a) From a scientific perspective, salt water is different from fresh water because it

contains dissolved sodium chloride. This gives salt water quite different properties from fresh water: it conducts electricity; it has a different melting point; it has a greater density; and it can potentially undergo more chemical reactions. Salt water and fresh water in the environment also support different forms of life.


Appendix G Review of Chemistry 20 Answer Key (b) From a technological perspective, the properties of salt water and fresh water

have to be considered whenever they are part of a system. For example, ships float differently in salt water and in fresh water because of their different densities. Also, salt water may not be useful for some technological applications (such as cooling heavy equipment) because of the chance that the salt water may react with the machinery.

19. Answers will vary greatly. For example:

Scientific research into space and space travel has led to the development of many technologies like telescopes, Teflon, freeze-dried foods, and pens that write in the absence of gravity.



Appendix G Review of Chemistry 20 Answer Key Unit 2 Gases (Chapter 4) 1. Seven ways in which empirical knowledge may be communicated: descriptions,

tables of evidence, graphs, empirical hypotheses, empirical definitions, generalizations, and scientific laws.

2. (Scientific law and generalizations should

• be based on evidence collected from many examples and replicated by many scientists;

• be able to describe, explain and predict nature; and • describe and explain current observations and correctly predict future events.

3. Natural phenomena: Answers will vary. wind, cloud formation

Technological products: Answers will vary. hot air balloons, neon lighting, car air bags

4. (a) At a constant temperature and chemical amount of gas, as the pressure increases,

the volume decreases. (b) At a constant pressure and chemical amount of gas, as the temperature decreases,

the volume decreases. (c) At a constant volume and temperature, if the chemical amount of gas inside a

container is increased, the pressure increases. 5. Answers will vary .

Test the effect of changing the amount of gas on the pressure of the gas (option (c)). Place a known chemical amount of gas in rigid sealed container with a pressure gauge attached. Record the pressure. Double the amount of gas. Record the pressure. Repeat for several trials.

Variables: manipulated – the chemical amount of gas in the container responding – the pressure of the gas controlled – the volume of the container, the temperature of the gas, gauge used, etc.

6. (a)



(b) (c)

7. 95.8 kPa × kPa101.325Hg mm760 = 718.55 mm Hg = 719 mm Hg

95.8 kPa × kPa101.325

atm1.00 = 0.9454 atm = 0.945 atm

95.8 kPa can be expressed as 710 mm Hg or 0.945 atm. 8. P1V1 = P2V2

V2 = 2

11

PVP

= atm 5.0

L 1.5 atm 1.0 ×

= 0.30 L According to Boyle’s law, the final volume of gas is 0.30 L.



9. 1

1

TV =

2

2

TV

T1 = (4.0 + 273) K = 277 K

T2 = 1

21

VVT

= L 0.750

L 0.800 K 277 ×

= 295.46 K T2 = (295.4 – 273) °C = 22 °C According to Charles’ law, the balloon is likely to burst at a temperature of 2 °C.

10. 1

11

TVP =

2

22

TVP

T1 = (22.0 + 273 K) = 295 K T2 = (150 + 273 K) = 423 K

P2 = 12

211

TVTVP

= K 295 L 0.300

K 423 L 0.150 kPa101×

××

= 72.41 kPa = 72.4 kPa According to the combined gas law, the pressure of the argon gas will be 72.4 kPa. 11. Boyle’s law states that as the volume of a gas decreases the pressure increases. This is

due to the compression of the gas molecules into a smaller space. As a result the molecules collide with each other and with the walls of the container more often. This increased number of collisions with the walls of the container is detected as an increase in pressure. Charles’ law states that as the temperature of a gas increases the volume also increases. This is due to the increased kinetic energy of the molecules. As the molecules speed up, they collide more often and with greater force. This forces them further apart, so they take up more volume.

12. The law of combining volumes states: in any chemical reaction, when measured at the

same temperature and pressure, volumes of gaseous reactants and products are always in simple ratios of whole numbers. Example: 1 L H2(g) + 1 L Cl2(g) → 2 L HCl(g) This is a 1:1:2 ratio. This is explained with Avogadro’s theory: equal volumes of gases contain equal numbers of molecules. As a result, gases combine in the same ratio as shown in the molar coefficients in a balanced chemical equation.



13. First calculate chemical amount of propane burned.

nC3H8 = 15.0 L × L24.8mol1

= 0.604 mol Chemical reaction equation: C3H8(g) + 5 O2(g) → 3 CO2(g) + 4 H2O(g) Ratio of propane used to carbon dioxide produced: 1:3 Calculate chemical amount of CO2 produced.

nCO2 = 0.604 mol × 13

= 1.81 mol

VCO2 = 1.81 mol × mol1L24.8

= 45.0 L According to the law of combining volumes and the molar volume of gases, when

15 L of propane burns, 45.0 L of carbon dioxide is produced. 14. Comparison of Ideal and Real Gases According to the KMT

Ideal gases Real gases The volume of individual molecules is negligible, even under high pressures.

The volume of individual molecules under pressure is significant, relative to the distance between them.

There are no forces of attraction between the molecules.

As gases are cooled, the forces of attraction between the molecules becomes significant, resulting in the conversion to a liquid.

No energy is lost when the molecules collide with each other.

Some energy is lost in a molecular collision resulting in a pressure lower than ideal.

15. nO2 = 25.0 g × g32.00mol1

= 0.781 mol T = (22.0 + 273) K = 295 K PV = nRT

V = P

nRT

= kPa98.1

K295KmolLkPa8.314mol0.781 ×⋅⋅

×

= 19.53 L = 19.5 L According to the ideal gas law, the oxygen would occupy 19.5 L.



16. STP: 0 °C and 101.325 kPa T = (0 + 273) K = 273 K

VH2 = PnRT

= kPa101.325


×

= 6.227 L = 6.23 L According to the ideal gas law, 0.278 mol of hydrogen would occupy 06.23 L at STP. Or use molar volume:

VH2 = 0.278 mol × mol1L22.4

= 6.23 L Using molar volume, 0.278 mol of hydrogen would occupy 06.23 L at STP. SATP: 25 °C and 1000 kPa T = (25.0 + 273) K = 298 K

V = P

nRT

= kPa100


×

= 6.887 L = 6.89 L

According to the ideal gas law, 0.278 mol of hydrogen would occupy 06.89 L at SATP.

or

VH2 = 0.278 mol × mol1L24.8

= 6.89 L Using molar volume, 0.278 mol of hydrogen would occupy 06.89 L at SATP. The difference in volume, between STP and SATP, is 0.66 L.

17. a) VCH4 = 400 kmol × mol1L24.8

kmol1mol1000

×

= 9 920 000 L = 9.92 × 106 L = 9920 kL = 9.92 × 103 kL

The volume of methane required for a year would be 9.92 × 103 kL.



b) T2 = (12.7 + 273) K = 285.7 K

V2 = 12

211

TPTVP

= K298kPa98.5

K285.7kL109.92kPa100 3

××××

= 9655 kL = 9.66 × 103 kL

According to the combined gas law, the volume of methane under the new conditions will be 9.66 × 103 kL.



Appendix G Review of Chemistry 20 Answer Key Unit 3: Solutions, Acids and Bases (Chapters 5 & 6) 1. A scientific perspective leads to research and explaining natural phenomena.

A technological perspective is concerned with the development and use of machines, instruments, and processes that have a social purpose. An ecological perspective deals with the relationships between living organisms and their environments. An economic perspective deals with the production, distribution, and consumption of wealth. A political perspective considers vote getting and re-election measures.

2. Answers will vary. The effect of burning fossil fuels on the global temperature; the

development of new medical diagnostic equipment; the use of water in oil sands development

3. (a) An industry spokesman… political (b) Laboratory research… scientific (c) The cost of ending sulfur… economic (d) Sulfur oxides and their… ecological (e) One of the most promising… technological 4. Answers will vary.

Patterns to look for: The goals for science are to make observations to develop an understanding of how things work, how they interact, what causes problems etc., and to develop logical explanations for the observations. The goals of technology are to develop substances, materials, and processes that solve problems, meet needs, save money, etc.

5. A homogeneous mixture is a combination of two or more substances in which all of

the particles are distributed uniformly throughout. It is a single phase. Examples: a glass of pop (with no bubbles), a crystal of sugar, air, salt dissolved in water

6. A solution is a homogeneous mixture consisting of a solute and a solvent. The solvent

is the substance in which the solute is dissolved. Usually there is a greater mass of solvent present than solute. The solvent is often, but not always, a liquid. Example: Table sugar dissolved in water: C6H22O11(s) in H2O(l), or C6H22O11(aq)

Sugar is the solute and water is the solvent. 7. Reactions between solids are rare. When solids dissolve in water, they break up into

individual molecules, ions, or atoms. Once in solution, the entities are mobile, and the change of two different entities colliding and reacting is increased. If the particles are ions, they may actually be attracted to one another via charge attraction. This will considerably increase the likelihood of reactions occurring. Life depends on many reactions taking place, in very specific ways. The majority of reactions in living



things take place in solution. For this reason, scientists believe that there is not likely to be life (as we know it) where there is no water.

8. Answers will vary.

Manufactured solutions at home: chlorine bleach, soft drinks, liquid weed killer, nail polish remover, liquid dish detergent, vinegar, window cleaner, etc. Natural solutions in the environment: air, ocean water, pond water, urine, blood serum, plasma, nectar, maple sap, etc.

9. An electrolyte is a substance that will conduct a current in aqueous solution or in the

molten state. Generally, ionic compounds and acids are electrolytes. Example: sodium chloride, NaCl(aq) A nonelectrolyte is a compound that will not conduct a current in aqueous solution or in the molten state. Generally, molecular compounds are nonelectrolytes. Example: sucrose, C12H22O11(aq)

10. Energy must be absorbed (supplied) to break bonds and energy is released whenever a new bond is formed. Both of these activities occur as a solute dissolves to become a solution. Energy is required to separate ions from an ionic crystal or to separate the molecules in a molecular solid. Some energy must also be supplied to break the bond between the water molecules, to allow the solute entities to come between them. New bonds form between the solute entities and the water molecules. If the amount of energy released is greater than the amount of energy supplied, the solution process is exothermic (i.e., releases heat, making the surroundings warmer). If the amount of energy supplied is greater than the amount of energy released, the solution process is endothermic (i.e., absorbs heat, making the surroundings cooler).

11. (a) Experimental Design

Each of the compounds is dissolved in a standard amount of water and tested with a conducting apparatus. Each solution is then tested with either pH paper or litmus paper.

(b) Expected Results Conductivity: If the solution conducts electricity, the compound is either ionic or

an acidic or basic molecular compound. If the solution does not conduct a current, the compound is neutral molecular. pH: If the solution has no effect on litmus, it is neutral. If the solution changes blue litmus to red, it is an acid. If the solution changes red litmus to blue, it is a base.

Analysis Solutions that conduct electricity but have neutral pH are likely to be ionic

compounds. Solutions that conduct electricity but do not have neutral pH are acids and bases (molecular). Solutions that do not conduct electricity are likely to have neutral pH, and are neutral molecular compounds.



12. (a) Li3PO4(s) → 3 Li+(aq) + PO43−(aq) (dissociation) (b) HCl(g) → H+(aq) + Cl− (aq) (ionization) (c) Al2(SO4)3(s) → 2 Al

3+(aq) + 3 SO42−(aq) (dissiciation) 13. (a) Sr2+(aq), OH−(aq), H2O(l) (b) H+(aq), NO3−(aq), H2O(l) (c) C3H8(aq), H2O(l) (d) CH3COO−(aq), H+(aq), H2O(l) (e) Ag+(aq), Cl−(aq), H2O(l) (f) CH3OH(l), H2O(l) 14. Solutions

• make handling and transport of chemicals easier, • makes it easier to run reactions to completion, and • makes it easier to control reaction type, extent of reaction, and rate of reaction.


Test each solution with blue litmus paper. The one that changes to red is HCl(aq). (HCl is the only acid present). Test the remaining solutions with a conducting apparatus. The one that does NOT conduct a current is CH3Cl(aq). (This is the only molecular compound remaining.) Add a small amount of H2SO4(aq) to each of the remaining two solutions. The one that forms a precipitate is the BaCl2(aq). (Barium ions react with sulfate ions to form barium sulfate, which is only slightly soluble. Sodium ions are soluble with sulfate ions.)

16. In chemistry labs, solution concentrations are usually expressed in moles per litre,

(mol/L). Most consumer products express concentration as a “mass per volume” percent, for example 5% by volume. Environmental studies usually use parts per million, ppm.

17. cNaCl = %100mL500

g12.5×

= 2.50 % W/V The solution has a concentration of 2.50 % W/V.

18. V = nc

= mol/L2.75

mol0.375

= 0.136 L or 136 mL 0.136 L of solution contains 0.375 mol of sodium hydroxide.



19. nCH3OH = g/mol32.05g100

= 3.12 mol

cCH3OH = nV

= L2.00

mol3.12

= 1.56 mol/L The windshield washer solution has an amount concentration of 1.56 mol/L. 20. (a) Calculate the chemical amount of calcium chloride required to make 1.00 L of

solution. n = c ·V = 0.251 mol/L × 1.00 L = 0.251 mol Determine the mass of solute required. m = nCaCl2 c = 0.251 mol × 110.98 g/mol = 27.855 g = 27.9 g The mass of calcium chloride required is 27.9 g.

(b) Procedure 1. Use a graduated cylinder to measure 300 mL of distilled water. Place the distilled water in a 1.00 L volumetric flask. 2. Weigh out 27.9 g of calcium chloride on a weighing paper. 3. Add the calcium chloride to the flask and swirl until completely dissolved. 4. Top the flask up with distilled water to exactly 1.00 L. (Fill up to the line) 5. Stopper the flask and invert several times to completely mix the contents.

21. (a) Determine the chemical amount of H3PO4(aq) required.

n = c ·V = 0.375 mol/L × 0.250 L = 0.09375 mol = 0.0938 mol Determine the volume of concentrated H3PO4(aq) required.

V = nc

= mol/L14.6

mol0.0938

= 0.006421 L = 0.00642 L = 6.42 mL 6.42 mL of concentrated phosphoric acid solution is required.



(b) Procedure 1. Use a 100 mL graduated cylinder to measure 100 mL of distilled water. Place the distilled water in a 250 mL volumetric flask. 2. Measure out 6.42 mL of concentrated phosphoric acid in a 10.0 mL graduated cylinder. 3. Pour the concentrated acid into the flask. 4. Top up the flask to 250 mL (to the line) with distilled water. 5. Stopper the flask and invert several times to mix contents.

22. FeCl3(aq) → Fe3+(aq) + 3 Cl−(aq)

2.1 mol/L 2.1 mol/L 3 × 2.1 mol/L

= 6.3 mol/L 23. For solids: (usually) As temperature increases, solubility increases.

For gases: As temperature increases, solubility decreases. 24. (a) Observations: Initially the copper(II) sulfate dissolves as it is added, forming a

blue solution. After a certain quantity of solid has been added, the solid no longer dissolves but begins to collect at the bottom of the container. Once the container is sealed, no further changes are observed to take place in the solution.

(b) Initially, copper(II) ions and sulfate ions separate from the crystals and disperse throughout the solution. A small number of ions are also rejoining the crystals. Eventually, as the concentration of ions in solutions reaches saturations, as many ions rejoin the crystal as leave the crystal, so no change is visible. At this point the system has reached dynamic equilibrium.

25. (a) HBr(aq) (b) H2SO3(aq) (c) HF(aq) (d) H2SO4(aq) 26. Hydroxide Concentration and pH

[H3O+(aq)] (mol/L) pH Acidic / basic /neutral 1.0 × 10−7 7.00 neutral

1 ×10−8 8.0 basic

2 × 10−4 3.7 acidic

6.23 × 10−9 8.205 basic

27. Cola is 100 times more acidic than carbonated water and 10 000 times more acidic

than pure water. 28. (a) CH3COOH(aq) + H2O(l) → H3O+(aq) + CH3COO−(aq) (b) HCO3−(aq) + H2O(l) → H2CO3(aq) + OH−(aq) (c) H3O+(aq) + OH−(aq) → 2 H2O(l)



29. (a) [OH−(aq)] = 10(−pOH)

= 10(−2.5)

= 3 × 10−3

The hydroxide ions in the window cleaning solution are in a concentration of 3 × 10−3.

(b) NH3(aq) + H2O(l) → NH4+(aq) + OH− (aq) (c) Ammonia is a weak base. It does not react completely with water. 30. Answers will vary.

Experimental Design Test 1: Samples of the unknown acids can be tested with a pH meter. Convert the

measured pH value into a hydronium ion concentration. If the concentration of the ion is the same as the original concentration of the acid, the acid is strong. If the hydronium ion concentration is lower than the original concentration, the acid is weak. OR, to distinguish between two samples: Measure the pH of the samples. The lower pH is the strong acid and the higher pH is the weak acid. Test 2: Measure the conductivity of the samples. The one with the higher conductivity is the strong acid and the one with the lower conductivity is the weak acid. Test 3: React each sample with a strip of magnesium ribbon. The one in which the reaction is faster is the strong acid and the one with the slower reaction is the weak acid. Controlled variables: concentration of the acids (should be identical) and mass of magnesium used.

31. (a) A molecule of a monoprotic acid can only react once with water (having only one

ionizable proton). A molecule of a polyprotic acid can react more than once with water (having more than one ionizable proton). A molecule of a monoprotic base can react only one time with water (as it can accept only one proton). A molecule of a polyprotic base can react more than one time with water (accepting more than one proton).

(b) H3BO3(aq) + H2O(l) → H3O+(aq) + H2BO3−(aq) H2BO3−(aq) + H2O(l) → H3O+(aq) + HBO32−(aq) HBO32−(aq) + H2O(l) → H3O+(aq) + BO3 3−(aq)

32. (a) H2O(l) + CO2(g) → H2CO3(aq) (carbonic acid)

H2CO3(aq) + H2O(l) → H3O+(aq) + HCO3−(aq) (b) Answers will vary.

Reduce the amount of carbon dioxide produced by replacing fossil fuels with alternative sources of energy, such as batteries. Trap the carbon dioxide before it enters the atmosphere and store it in underground caverns.



33. Answers will vary. Intended consequences pesticides will kill unwanted species when applied as a spray. Unintended consequences: • Removal of a species of insect in an area may cause another undesirable species

to move into the area. • The food source of a desirable species, such as a particular bird, could be

reduced, so the bird population would drop. • The pesticide could get into groundwater and be carried to other area. • The pesticide may be taken up by animals higher in the food chain, be

biomagnified, and cause health problems. • Some pesticides may be carcinogenic. • Species may become resistant to the effects of the pesticide. • Some acidic pesticides could contaminate local ponds and streams, causing

them to become acidic and thereby affecting the aquatic species.



Appendix G Review of Chemistry 20 Answer Key Unit 4: Quantitative Relationships (Chapters 7 & 8) 1. Chemistry is a science that deals with matter and its changes. It is international in

scope and is involved with natural products and processes. It is largely theoretical and emphasizes ideas and concepts over practical applications. The focus of chemical technology is to provide solutions to practical problems. It is more localized and is involved with humanly developed processes and products. It is empirical in its approach and emphasizes methods and materials over understanding.


Batteries – must contain enough reactants to supply sufficient energy. Car engines – based on chemical reactions of combustion. In particular, the balance of gasoline to oxygen is important.

3. Qualitative analysis is used to identify whether a certain substance is present in a

sample. Example: a sample of an unknown gas could be tested with limewater to determine if it is carbon dioxide. Quantitative analysis is used to determine the amount of substance that is present in a sample. Example: A titration procedure could be used to determine the concentration of an acid and stoichiometry used to determine the chemical amount and therefore the mass of acid in a given sample.

4. (a) Co2+(aq) + 2 OH−(aq) → Co(OH)2(s) (spectator ions: Na+, Cl−) (b) Ag+(aq) + I−(aq) → AgI(s) (spectator ions: Ca2+, NO3−) (c) 2 Ag+(aq) + Zn(s) → 2 Ag(s) + Zn2+(aq) (spectator ions: NO3−) (d) H+(aq) + OH−(aq) → H2O(l) (spectator ions: Ca2+, Cl−) (e) Al3+(aq) + 3 OH−(aq) → Al(OH)3(s) (no spectator ions) 5. Stoichiometry refers to relative amounts (ratios) of reactants and products in a given

reaction. The ratios referred to come from the coefficients in the balanced chemical equation.

6. Gravimetric stoichiometry: This is a procedure used for calculating masses of

reactants and products. The measured substances must be either solids or liquids. Gas stoichiometry: This is a procedure is used to calculate the quantities of

gaseous reactants and products. This will involve gas volumes, temperatures, pressures, molar volumes, and the ideal gas law. Solution stoichiometry: This is a procedure (general method remains the same) for calculations involving dissolved substances (solutions). It involves the use of concentration and volume as conversion factors in order to determine the chemical amounts and masses of substances used or produced.



7. Stoichiometry allows industry to accurately determine the exact chemical amounts of each required reactant for a particular chemical reaction. By combining just the right masses (or volumes), all the reactants can be used up with none left over. This maximizes the desired yield and minimizes waste.

8. (a) 3 C(s) + 2 Fe2O3(s) → 4 Fe(l) + 3 CO2(g) (b) Three moles of solid carbon are reacted with two moles of solid iron(III) oxide to

produce four moles of liquid iron and three moles of carbon dioxide gas. (c) masses:

36.03 g C + 319.40 g Fe2O3 → 223.40 g Fe + 132.03 g CO2 (d) The total mass of the reactants (355.43 g) is equal to the total mass of the

products. This illustrates the law of conservation of mass (matter). 9. 2 KI(aq) + Pb(NO3)2(aq) → 2 KNO3(aq) + PbI2(s)

nKI = 2.93 g × g166.00

mol1

= 0.0177 mol

nPbI2 = 0.0177 mol × 21

= 0.00882 mol

mPbI2 = 0.00882 mol × mol1g461.01

= 4.068 g = 4.07 g According to gravimetric stoichiometry, 4.07 g of lead(II) iodide is produced when

2.93 g of potassium iodide reacts with excess lead(II) nitrate. 10. CaCl2(aq) + Na2OOCCOO(aq) → CaOOCCOO(s) + 2 NaCl(aq) nCaOOCCOO(s) = 0.452g/128.10 g/mol = 0.003528 mol nCaCl2 = 0.003528 × 1

1

= 0.003528 mol mCaCl2 = 0.003528 mol × mol1

g110.98

= 0.3915 g = 0.392 g According to gravimetric stoichiometry the mass of calcium chloride present in the

original solution is 0.392 g.



11. (a) nBa(SO4)2 = 1.50 g × g233.40mol1

= 0.00643 mol

nC = 0.00643 mol × 12

= 0.0129 mol

mC = 0.0129 mol × mol1

g12.01

= 0.1543 g = 0.154 g According to gravimetric stoichiometry, the mass of carbon required to react with

1.50 g of barium sulfate is 0.154 g. (b) Assume that: all of the barium sulfate has reacted (so the reaction is quantitative);

sufficient amounts of carbon were present to react all the barium sulfate. 12. (a) 6 NaOH(aq) + Al2(SO4)3(aq) → 3 Na2SO4(aq) + 2 Al(OH)2(s)

nAl2(SO4)3 = 1.50 g × g342.17mol1

= 0.00438 mol

nAl(OH)3 = 0.00438 mol × 12

= 0.00877 mol

mAl(OH)3 = 0.00877 mol × mol1g78.01

= 0.6839 g = 0.684 g

According to the stoichiometric method, the mass of aluminium hydroxide precipitating out of solution is predicted to be 0.684 g.

(b) % yield = 100g0.684

g0.96×

= 140 %

(c) The actual yield could be higher than the predicted yield because of experimental error (such as inaccurate measuring of the reactants or products); because some other contaminant also precipitated out of solution; or because the theory (gravimetric stoichiometry) used to predict the result was incorrect.



13. 4 Al(s) + 3 O2(g) → 2 Al2O3(s)

nAl = g/mol26.98

g/kg1000kg100 ×

= 3.71 × 103 mol

nO2 = 3.71 × 103 mol ×

43

= 2.78 × 103 mol

VO2 = 2.78 × 103 mol ×

mol1L24.8

= 6.89 × 104 L According to the stoichiometric method, 1000 kg of aluminium will react with 6.89 ×

104 L of oxygen at SATP. 14. CaH2(s) + 2 H2O(l) → Ca(OH)2(s) + 2 H2(g)

nCaH2 = 50.0 g × g42.10mol1

= 1.19 mol

nH2 = 1.19 mol × 12

= 2.38 mol

VH2 = PnRT T = (22+273) K

= 295 K

= 2.38mol 8.314 kPa L/mol K 295K96.5kPa

× ⋅ ⋅ ×

= 60.37 L = 60.4 L According to the stoichiometric method, 50 g cartridge of calcium hydride can

produce 60.4 L of hydrogen gas under the conditions specified. 15. 2 NaOH(aq) + H2SO4(aq) → Na2SO4(aq) + 2 H2O(l) nNaOH = 2.12 mol/L × 0.032 L = 0.0678 mol

nH2SO4 = 0.0678 mol × 21

= 0.0339 mol

cH2SO4 = L0.0100mol0.0339

= 3.39 mol/L According to volumetric stoichiometry, the concentration of the sulfuric acid in the

car battery was 3.39 mol/L.



16. nAg = 1.68 kg × g107.87

mol1kg 1

g1000×

= 15.57 mol = 15.6 mol

nHNO3 = 15.6 mol × 34

= 20.8 mol

VHNO3 = mol/L15.4mol20.8

= 1.348 L = 1.35 L According to the stoichiometric method, at least 1.35 L of nitric acid will be required

to react completely with the solid silver. 17. The limiting reagent is completely used up during the reaction. There will be some of

the excess reagent remaining after the reaction is complete. The limiting reagent determines the quantity of products.

18. An excess of one of the reagents is used in order to guarantee that the other reagent is

completely consumed during the reaction. 19. (a) CaCO3(s) + 2 HCl(aq) → CaCl2(aq) + H2CO3(aq)

nCaCO3 = 750 mg × g100.09mol1

mg1000g1

×

= 0.00749 mol nHCl = 0.10 mol/L × 0.200 L = 0.020 mol Chemical amount of HCl required to use up all of the CaCO3:

nHCl = 0.00749 ml × 12

= 0.0150 mol There is an excess of HCl. All of the CaCO3 will be used up. nHCl in excess = 0.020 mol – 0.015 mol = 0.0050 mol According to the stoichiometric method, there is 0.0050 mol of HCl(aq) in excess.



(b) nCaCl2 = 0.00749 mol × 11

= 0.00749 mol

mCaCl2 = 0.00749 mol × mol1g110.98

= 0.8316 g = 0.832 g According to the stoichiometric method, there is 0.832 g of calcium chloride will

be formed. 20. Materials unknown sodium carbonate solution; standardized 0.352 mol/L hydrochloric acid);

burette; stand with ring and clamp; graduated cylinder; titration flask; methyl orange indicator Analysis

Na2CO3(aq) + 2 HCl(aq) → 2 NaCl(aq) + H2CO3(aq)

Average volume HCl(aq) = 3

mL15.2)15.2(15.3 ++

= 15.23 mL = 15.2 mL nHCl = 0.352 mol/L × 0.0152 L = 0.00536 mol

nNa2CO3 = 0.00536 mol × 21

= 0.00268 mol

cNa2CO3 = L0.025mol0.00268

= 0.107 mol/L According to the stoichiometric method, the concentration of the sodium

carbonate solution is 0.107 mol/L. 22. (a)



(b) (c) A suitable indicator for this titration would be bromothymol blue. This indicator

changes colour (blue to yellow) between a pH 6.0 and 7.6.This range corresponds to the predicted equivalence point of 7.0 and will correctly indicate the volume of acid required to neutralize this base.

(d) The graph would look similar but would start in the lower left of the graph and end in the upper right. The vertical portion would be in the same place. The answers to (b) and (c) would remain the same. The equivalence point remains at 7.0 and the indicator would work in either direction. It would now change from yellow to blue instead of blue to yellow.


AppendixG_AK.pdfReview Unit II final.pdfAppendix G Review of Chemistry 20 Answer KeyAccording to the law of combining volumes and the molar volu

Review unit III final.pdfAppendix G Review of Chemistry 20 Answer Key

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