Химия на английском. 2

2.27. Which element in each pair is more electronegative? a) K, As; b) N, Sb; c. Sr, Be.

2.28. For each of the following properties, indicate whether fluorine or bromine has a larger value.

2.29.Give the correct responses in the following questions. Explain why according to the location of the elements in the periodic table.

a)Which has the lower ionization energy? Li or K.

b)Which would be more polar? HF or HBr.

c)Which is more nonmetallic? F or I.

d)Which is more electronegative? K or Rb.

e)Which has more outer shell electrons? Ca or C.

f)Which would you expect to be more ionic? LiF or HCl.

2.30.Element A, with three electrons in its outer energy level, is in Period 4 of the periodic table. How does the number of its valence electrons compare with that of Element B, which is in Group IIIA and Period 6? Use Lewis structures to help you express your answer.

2.31.Predict the composition and ionic or molecular character for binary compounds of the following elements with hydrogen:

a) Li, Sn, Se; b) K, C, Se; c) Na, B, Sb.

2.32. Write formulas for the products of reactions of the following elements with the excess of oxygen. Are they of acidic or basic character?

a) Na, Mg, Al, Si, P, S, Cl; b) Rb, Sr, In, Sn, Sb, Te, I.

2.33. Describe an element, using only the periodic table. The following answers must be given fully.

a)What is the name and the symbol of the element?

b)What is its location in the periodic table (group and period)?

c)How many protons and neutrons are in the nucleus of its atom?

d)How many electrons are in its atom?

e)What is the electron configuration of the element?

f)How many valence electrons does it have (according to Lewis structure)?

g)What element block does it belong to?

h)Is it a metal, semi-metal, nonmetal?

i)What is its compound with hydrogen, is it of ionic or molecular character?

j)What is the formula of its oxide (in the higher oxidation state), is it of acidic or basic character?

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3.ENERGY AND CHEMICAL CHANGE

3.1.An exothermic reaction releases 86.5 kJ. How many kilocalories of energy are released?

3.2.If an endothermic process absorbs 256 J, how many kilocalories are absorbed?

3.3.What is the equivalent in joules of 126 calories?

3.4.Convert 455 kilojoules to kilocalories.

3.5.In the construction of bridges and skyscrapers, gaps must be left between adjoining steel beams to allow for the expansion and contraction of the metal due to heating and cooling. The temperature of a sample of iron with a mass of 10.0 g changed from 50.4°C to 25.0°C with the release of 114 J heat. What is the specific heat of iron?

3.6.If the temperature of 34.4 g of ethanol increases from 25.0°C to 78.8°C, how much heat has been absorbed by the ethanol?

3.7.A nugget of pure gold with the mass of 4.50 g absorbed 276 J of heat. What was the final temperature of the gold if the initial temperature was 25.0°C? The specific heat of gold is in the table (problem 3.8.).

3.8.A 155 g sample of an unknown substance was heated from 25.0°C to 40.0°C. In the process, the substance absorbed 5696 J of energy. What is the specific heat of the substance? Identify the substance among those listed in the table:

Specific Heats of Common Substances at 298 K (25 oC)

3.9.What is the specific heat of an unknown substance if a 2.50-g sample releases 12.0 cal as its temperature changes from 25.0°C to 20.0°C?

3.10.The temperature of 55.6 grams of a material decreases by 14.8°C when it loses 3080 J of heat. What is its specific heat?

3.11.What is the specific heat of a metal if the temperature of a 12.5 g sample increases from 19.5°C to 33.6°C when it absorbs 37.7 J of heat?

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3.12.A sample of ethylene glycol, used in car radiators, has a mass of 34.8 g. The sample liberates 783 J of heat. The initial temperature of the sample is 22.1°C. What is the final temperature?

3.13.A sample of ethanol, C2H5OH, absorbs 23.4 kJ of energy. The temperature of the sample increases from 5.6°C to 19.8°C. What is the mass of the ethanol sample? The

specific heat capacity of ethanol is 2.46 J/(g °C).

3.14.A child’s swimming pool contains 1000 L of water. When the water is warmed by solar energy, its temperature increases from 15.3°C to 21.8°C. How much heat does the water absorb?

3.15.What temperature change results from the loss of 255 kJ from a 10.0 kg sample of water?

3.16.If 335 g water at 65.5°C loses 9750 J of heat, what is the final temperature of the water?

3.17.The temperature of a sample of water increases from 20.0°C to 46.6°C as it absorbs 5650 J of heat. What is the mass of the sample?

3.18.Explain how you could calculate the heat released in freezing 0.250 mol water.

3.19.How much heat is required to warm 122 g of water by 23.0°C?

3.20.Which of the following processes are exothermic? Endothermic?

a)C2H5OH(l) → C2H5OH(g).

b)NH3(g) → NH3(l).

c)Br2(l) → Br2(s).

d)NaCl(s) → NaCl(l).

e)C5H12(g) + 8 O2(g) → 5 CO2(g) + 6 H2O(l).

3.21.Write the correct sign of H for each of the following changes in physical state.

a)C2H5OH(s) → C2H5OH(l).

b)H2O(g) → H2O(l).

c)CH3OH(l) → CH3OH(g).

d)NH3(l) → NH3(s).

3.22.A reaction is characterized by H = –500 kJ/mol. Does the reaction mixture absorb heat from the surroundings or release heat to them?

3.23.A reaction is characterized by H = +280 kJ/mol. Does the mixture of reactants and products release heat to the surroundings or absorb heat from them?

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3.24. For each of the following reactions, determine: (a) does the enthalpy increase or decrease; (b) is Hreactant > Hproduct or is Hproduct > Hreactant; (c) is H positive or negative?

a)Al2O3(s) → 2Al(s) + 3/2 O2(g) (endothermic);

b)Sn(s) + Cl2(g) → SnCl2(s) (exothermic).

3.25. Consider the following reaction:

N2(g) + O2(g) → 2 NO(g); Ho = –180.6 kJ.

a)Draw an enthalpy diagram for the reaction.

b)What is the enthalpy change for the formation of one mole of nitrogen monoxide?

c)What is the enthalpy change for the reaction of 1.00 102 g of nitrogen with sufficient oxygen?

3.26.The reaction of iron with oxygen is very familiar. You can see the resulting rust on buildings, vehicles, and bridges. You may be surprised, however, at the large amount of heat that is produced by this reaction.

4 Fe(s) + 3 O2(g) → 2 Fe2O3(s) + 1.65 103 kJ.

a)What is the enthalpy change for this reaction?

b)Draw an enthalpy diagram that corresponds to the thermochemical equation.

c)What is the enthalpy change for the formation of 23.6 g of iron(III) oxide?

3.27. Tetraphosphorus decoxide, P4O10, is an acidic oxide. It reacts with water to produce phosphoric acid, H3PO4, in an exothermic reaction.

P4O10(s) + 6 H2O(l) → 4 H3PO4(aq); Ho = –257.2 kJ

a) How much energy is released when 5.00 mol of P4O10 reacts with excess water?

c)How much energy is released when 235 g of H3PO4 is formed?

3.28.Calcium oxide, CaO, reacts with carbon in the form of graphite. Calcium carbide, CaC2, and carbon monoxide, CO, are produced in an endothermic reaction.

CaO(s) + 3 C(s) + 462.3 kJ → CaC2(s) + CO(g)

a)246.7 kJ of energy is available to react. What mass of calcium carbide is produced, assuming sufficient reactants?

b)What is the enthalpy change for the reaction of 46.7 g of graphite with excess calcium oxide?

c)1.38 1024 molecules of calcium oxide react with excess graphite. How much energy is needed?

3.29.Acetylene, C2H2, undergoes complete combustion in oxygen. Carbon dioxide and water are formed. When one mole of acetylene reacts, 1.3 103 kJ of energy is released.

a)Draw a diagram to represent the thermochemical equation.

b)How much energy is released when the complete combustion of acetylene produces 1.50 g of water?

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3.30. Ethene, C2H4, reacts with water to form ethanol, CH3CH2OH: C2H4(g) + H2O(l) → CH3CH2OH(l).

Determine the enthalpy change of this reaction, given the following thermochemical equations.

1. CH3CH2OH(l) + 3 O2(g) → 3 H2O(l) + 2 CO2(g); Ho = –1367 kJ. 2. C2H4(g) + 3 O2(g) → 2 H2O(l) + 2 CO2(g); Ho = –1411 kJ.

3.31. A typical automobile engine uses a lead-acid battery. During discharge, the following chemical reaction takes place:

Pb(s) + PbO2(s) + 2 H2SO4(l) → 2 PbSO4(aq) + 2 H2O(l). Determine the enthalpy change of this reaction, given the following equations.

3.32. Mixing household cleansers can result in the production of hydrogen chloride gas, HCl(g). Not only is this gas dangerous in its own right, but it also reacts with oxygen to form chlorine gas and water vapour.

4 HCl(g) + O2(g) → 2 Cl2(g) + 2 H2O(g). Determine the enthalpy change of this reaction, given the following equations.

3.33. Calculate the enthalpy change of the following reaction between nitrogen gas and oxygen gas, given thermochemical equations (1), (2), and (3).

3.34. Calculate Ho for the reaction 2 C(s) + 2 H2(g) → C2H4(g) given the following thermochemical equations:

3.35. Calculate Ho for the reaction HCl(g) + NH3(g) → NH4Cl(s) given the following thermochemical equations:

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3.36. Calculate the enthalpy change of the following reaction, given equations (1), (2), and (3):

3.37. Hydrogen can be added to ethene, C2H4, to obtain ethane, C2H6.

C2H4(g) + H2(g) → C2H6(g).

Show that the equations for the formation of ethene and ethane from their elements can be algebraically combined to obtain the equation for the addition of hydrogen to ethene.

3.38. Zinc sulfide reacts with oxygen gas to produce zinc oxide and sulfur dioxide. 2 ZnS(s) + 3 O2(g) → 2 ZnO(s) + 2 SO2(g).

Write the chemical equation for the formation of the indicated number of moles of each compound from its elements. Algebraically combine these equations to obtain the given equation.

3.39. The standard molar enthalpy of formation of calcium carbonate is –1207.6 kJ/mol. Calculate the enthalpy of formation of calcium oxide, given the following equation:

CaO(s) + CO2(g) → CaCO3; Ho = –178.1 kJ.

3.40. Small amounts of oxygen gas can be produced in a laboratory by heating potassium chlorate, KClO3.

2 KClO3(s) → 2 KCl(s) + 3 O2(g).

Calculate the enthalpy change of this reaction, using enthalpies of formation, namely: Hof (KClO3) = –397.7 kJ/mol; Hof (KCl) = –436.5 kJ/mol.

3.41. Use the following equation to answer the questions below. CH3OH(l) + 1.5 O2(g) → CO2(g) + 2 H2O(g).

a) Calculate the enthalpy change of the complete combustion of one mole of methanol, using enthalpies of formation: Hof (CH3OH) = –239.2 kJ/mol; Hof (CO2) = –393.5 kJ/mol; Hof (H2O(g)) = –241.8 kJ/mol.

b) How much energy is released when 125 g of methanol undergoes complete combustion?

3.42. In the early 1960s, Neil Bartlett, at the University of British Columbia, was the first person to synthesize compounds of the noble gas xenon. A number of noble gas compounds (such as XeF2, XeF4, XeF6, and XeO3) have since been synthesized. Consider the reaction of xenon difluoride with fluorine gas to produce xenon tetrafluoride: XeF2(g) + F2(g) → XeF4(s). Use the following standard molar enthalpies of formation to calculate the enthalpy change for this reaction: Hof (XeF2) =

– 108 kJ/mol; Hof (XeF4) = –251 kJ/mol.

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3.43. Hydrogen is a very appealing fuel, in part because burning it produces only nonpolluting water. One of the challenges that researchers face in making hydrogen fuel a reality is how to produce hydrogen economically. Researchers are investigating methods of producing hydrogen indirectly. The following series of equations represent one such method.

3 FeCl2(s) + 4 H2O(g) → Fe3O4(s) + 6 HCl(g) + H2(g); Ho = 318 kJ; Fe3O4(s) + 3/2 Cl2(g) + 6 HCl(g) → 3 FeCl3(s) + 3 H2O(g) + 1/2 O2(g); Ho = –249 kJ;

3 FeCl3(s) → 3 FeCl2(s) + 3/2 Cl2(g); Ho = 173 kJ.

a)Show that the net result of the three reactions is the decomposition of water to produce hydrogen and oxygen.

b)Use Hess’s law and the enthalpy changes for the reactions to determine the enthalpy change for the decomposition of one mole of water. Check your answer, using

the enthalpy of formation of water, Hof (H2O(g)) = –241.8 kJ/mol.

3.44. Predict the sign of entropy change S for the following reaction. Explain the basis for your prediction: 2 H2(g) + O2(g) → 2 H2O (g).

3.45. Predict the sign of entropy change S for each reaction or process.

a)FeS(s) → Fe2+(aq) + S2–(aq)

b)SO2(g) + H2O(l) → H2SO3(aq)

3.46. If Ho = 285.4 kJ/mol and So = 137.55 J/mol K, calculate Gibbs free energy change Go at 25 oC for the supposed process. Is the reaction spontaneous? Is either or both of the driving forces ( Ho and So) for the reaction favorable?

3.47. Calculate Go at 25°C for the reaction 2 NO2(g) → N2O4(g) given functions Ho = –57.20 kJ/mol and So = –175.83 J/mol K. Is this reaction spontaneous? What is the driving force for spontaneity?

3.48. The standard Gibbs free energy of formation has the values: –286.06 kJ/mol for NaI(s), –261.90 kJ/mol for Na+(aq), and –51.57 kJ/mol for I–(aq) at 25°C. Calculate Go

for the reaction in water: NaI(s) Na+(aq) + I–(aq).

3.49. Calculate Go at 700 K for the reduction of the oxides of iron and copper by carbon, represented by the equations:

Values of Gof at 700 K are –92 kJ/mol for CuO(s), –632 kJ/mol for Fe2O3(s), and –395 kJ/mol for CO2(g). Which oxide can be reduced using carbon in a wood fire (which has a temperature of about 700 K), assuming standard state conditions?

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4.RATES OF CHEMICAL REACTIONS

4.1.The following reaction is second order in A and first order in B: 2 A + B → 3 C. What is the rate law equation for the reaction below? (Assume that A, B, and C are the same compounds for each reaction.)

4 A + 2 B → 6 C.

4.2. Consider the general reaction below:

a A + b B → c C + d D.

Based on this equation, is it correct to write the following rate law equation by inspection? Explain your answer.

Rate = k [A]a [B]b.

4.3. Consider the following rate law equation.

Rate = k [A]2 [B].

a)How does the reaction rate change if [A] decreases by a factor of 2 and [B] increases by a factor of 4?

b)How does the reaction rate change if [A] and [B] are doubled?

4.4. Consider the following rate law equation:

Rate = k [HCrO4–] [HSO3–]2 [H+].

a)What is the order with respect to each reactant?

b)What is the overall reaction order?

c)What are the units for the rate constant?

4.5.Cyclopropane, C3H6, is used in the synthesis of organic compounds and as a fastacting anesthetic. It undergoes rearrangement to form propene C3H6. If cyclopropane disappears at a rate of 0.25 mol/s, at what rate is propene being produced?

4.6.Ammonia, NH3, reacts with oxygen to produce nitric oxide, NO, and water vapour.

4 NH3(g) + 5 O2(g) → 4 NO(g) + 6 H2O(g).

At a specific time in the reaction, ammonia is disappearing at rate of 0.068 mol/(L s). What is the corresponding rate of production of water?

4.7. Hydrogen bromide reacts with oxygen to produce bromine and water vapour. 4 HBr(g) + O2(g) → 2 Br2(g) + 2 H2O(g).

How does the rate of decomposition of HBr (in mol/(L s) ) compare with the rate of formation of Br2 (also in mol/(L s) )? Express your answer as an equation.

4.8. Magnesium metal reacts with hydrochloric acid to produce magnesium chloride and hydrogen gas:

Mg(s) + 2 HCl(aq) → MgCl2(aq) + H2(g). Over an interval of 1.00 s, the mass of Mg(s) changes by −0.011 g.

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a)What is the corresponding rate of consumption of HCl(aq) (in mol/s)?

b)Calculate the corresponding rate of production of H2(g) (in L/s) at 20°C and

101 kPa.

4.9. In the following reaction, the rate of production of sulfate ions is calculated to be 1.25 10–3 mol/(L s).

2 HCrO4– + 3 HSO3– + 5 H+ → 2 Cr3+ + 3 SO42–

a)What is the corresponding rate at which [HSO3–] decreases over the same time

interval?

b)What is the corresponding rate at which [HCrO4–] decreases over the same time interval?

4.10.Use the data in the following table to calculate the average reaction rates.

a) Calculate the average reaction rate expressed in moles H2 consumed per liter per second.

b) Calculate the average reaction rate expressed in moles Cl2 consumed per liter per second.

c) Calculate the average reaction rate expressed in moles HCl produced per liter per second.

4.11. When heated, ethylene oxide decomposes to produce methane and carbon

4.12. Iodine chloride reacts with hydrogen to produce iodine and hydrogen chloride:

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4.13. Sulfuryl chloride (also known as chlorosulfuric acid and thionyl chloride), SO2Cl2, is used in a variety of applications. At a certain temperature, the rate of decomposition

a)Write the rate law equation for the decomposition of sulfuryl chloride.

b)Determine the rate constant, k, for the reaction, with the appropriate units.

4.14.A first-order decomposition reaction has a rate constant of 2.34 10–2 year–1. What is the half-life of the reaction? Express your answer in years and in seconds.

4.15.When cyclopropane, C3H6, undergoes rearrangement to propene at 1000°C, the first-order rate constant for the decomposition of cyclopropane is 9.2 s−1.

a)Determine the half-life of the reaction.

b)What percent of the original concentration of cyclopropane will remain after 4 half-lives?

4.16.The following reaction is exothermic.

2 ClO(g) → Cl2(g) + O2(g).

Draw and label a potential energy diagram for the reaction. Propose a reasonable activated complex.

4.17. Consider the following reaction.

AB + C → AC + B; H = 65 kJ; Ea(rev) = 34 kJ.

a)Draw and label a potential energy diagram for this reaction.

b)Calculate and label Ea(fwd). Include a possible structure for the activated

complex.

4.18. Consider the reaction: C + D → CD; H = –132 kJ; Ea(fwd) = 61 kJ.

a)Draw and label a potential energy diagram for this reaction.

b)Calculate and label Ea(rev). Include a possible structure for the activated

complex.

4.19.In the upper atmosphere, oxygen exists in forms other than O2(g). For example, it exists as ozone, O3(g), and as single oxygen atoms, O(g). Ozone and atomic oxygen react to form two molecules of oxygen. For this reaction, the enthalpy change is −392 kJ and the activation energy is 19 kJ. Draw and label a potential energy diagram. Include a value for Ea(rev). Propose a structure for the activated complex.

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