Topic 7 + 8 - Equilibrium + Acid/Bases

(a) H2SO3 is a Brønsted-Lowry acid as it acts as a proton donor. (b) HSO3−(aq) and H2O(l) (c) nitric acid, HNO3 (d) the acidification of lakes and the leaching of minerals into lakes with the effect of killing young fish (e) S(s) + O2(g) → SO2(g); SO2(g) + H2O(l) → H2SO3(aq)

(a) Sulphurous acid, H2SO3, can be described as a Brønsted-Lowry acid. State what is meant by this description. (1) (b) The bisulphite ion, HSO3−, is said to be amphiprotic. In the reaction of HSO3− with the oxonium cation, H3O+, identify the two species acting as bases. (1) HSO3−(aq) + H3O+(aq) ⇋ H2SO3(aq) + H2O(l) (c) State one other acid present in acid rain that originates from car exhaust emissions. (1) (d) Describe two harmful effects of acid deposition on the natural environment. (3) (e) Starting with elemental sulphur, S(s), write a series of balanced equations showing how sulphurous acid, H2SO3(aq) is formed in the atmosphere. (3)

i. The forward reaction line is falling from the starting position where only reactants are present. Their concentration steadily decreases as does the rate of the forward reaction. ii. The reverse reaction line is increasing from zero as initially there were no products present. As product is produced the rate of the reverse reaction steadily increases. At equilibrium the rate of the forward reaction = rate of the reverse reaction and the two lines meet and stay level. (b) As the temperature is increased the value of Kc increases. This means there is a greater concentration of products and the equilibrium has shifted to the right. An increase in temperature always favours the endothermic process. The forward reaction is therefore endothermic.

(a) The following equilibrium is established at 1700°C. CO2(g) + H2(g) ⇋ H2O(g) + CO(g) If only carbon dioxide gas and hydrogen gas are present initially, sketch on a graph a line representing rate against time for: i. the forward reaction and ii. the reverse reaction until shortly after equilibrium is established. Explain the shape of each line. (7) (b) Kc for the equilibrium reaction is determined at two different temperatures. At 850°C, Kc = 1.1 whereas at 1700°C, Kc = 4.9. On the basis of these Kc values explain whether the reaction is exothermic or endothermic. (3)

1. [H+]=10−pH [H+]=10−2.75 [H+]=1.78×10−3mol/l 2. [H+]=10−pH [H+]=10−9.10 [H+]=7.94×10−10mol/l 3. Yes, the first solution has a pH below 7.0 and is therefore acidic.

1. Calculate the [H+] of a solution with a pH of 2.75. 2. Calculate the [H+] of a solution with a pH of 9.10. 3. Is either of the two solutions acidic?

1. [H+]=10−pH [H+]=10−4.8 [H+]=1.58×10−5mol/l [OH−]=Kw/[H+] [OH−] = 1.0×10−14/1.58×10−5 [OH−]=6.31×10−10mol/l 2. [OH−]=0.01mol/l [H+] = Kw/[OH−] [H+]=1.0×10−14/0.01 [H+]=1.0×10−12mol/l pH=−log[H+] pH=−log(1.0×10−12) pH=12

1. Calculate the [OH−] of a solution with a pH of 4.8. 2. Calculate the pH of a solution with [OH−] of 0.01 mol/l.

1. pH=−log[H+] pH=−log(0.38) pH=0.42 2. pH=−log[H+] pH=−log(3.8×10−9) pH=8.42 3. Yes, the second solution is basic because its pH is above 7

1. Calculate the pH of a solution with [H+] of 0.38 mol/l. 2. Calculate the pH of a solution with [H+] of 3.8×10−9mol/l. 3. Is either of the two solutions basic?

In this titration the acid is in the erlenmeyer (added by a pipette of 25.00 ml). Since the indicator is always added to the solution in the erlenmeyer, the indicator is initially in an acidic solution. When base is added to the point of equivalence, the solution is neutral. Using section 16 in the IB data booklet shows you the color will thus change from red to yellow.

25.00 ml of a solution of nitric acid is titrated with a NaOH solution. Predict the change of color when methyl red is used as indicator.

Balance the reaction equations: H2SO4 + 2NaOH → 2H2O + Na2SO4 Convert known amount to moles: 25ml × 4.5 mmol / ml = 112.5 mmol H2SO4 present Find the molar ratio of unknown/known:NaOH / H2SO4 = 2/1 Multiply the known amount in moles with the molar ratio to find the amount of unknown substance in moles: 112.5mmol × 2/1 = 225mmol NaOH needed. Convert to other units if necessary: 225mmol/1.42 mmol/ml = 158 ml of NaOH solution needed

Calculate the volume of a 1.42 M NaOH solution required to titrate 25.00 ml of a 4.50 M H2SO4 solution.

NaOH - strong base CH3COOH - weak acid NH3 - weak base HCl - strong acid H2CO3 - weak acid H2SO4 - strong acid

Classify the following chemicals as: strong acid or weak acid, strong base or weak base. NaOH CH3COOH NH3 HCl H2CO3 H2SO4

Amphiprotic substances can donate hydrogen ions (H+) as well as accept them. Examples of amphiprotic substances are H2O, NH3, HSO−4, H2PO−4, etc. Remember that amphiprotic substances can donate and accept protons. Amphoteric substances can react both as acid and as base. This is the more general definition of the two described here. The issue is that some amphoteric substances can react both as acid and as base without donating or receiving protons.* E.g. ZnO+2OH−+H2O→[Zn(OH)4]2−. Although ZnO does not have any protons to donate, it does react with base and is thus considered an acid. Likewise, ZnO will act as base without accepting protons when reacted with acid, e.g. ZnO+2H+→Zn2++H2O.

Compare the terms amphiprotic and amphoteric.

OH−+HSO−4⇌H2O+SO2−4 - Identify the acid and the base. In this example HSO−4 is the acid and OH− is the base. The acid will donate a proton (H+) to the base. This turns HSO−4 into SO2−4 and OH− into H2O. HSO−4, because it donates a proton. H2O, because the conjugate acid of a base is the species you get when one proton is added to the base. HSO−4, because it can donate a proton to a base as well as accept a proton from an acid. Note that H2O can also be accepted as an amphiprotic substance, because it too can donate a proton to a base as well as accept a proton from an acid.

Complete the following reaction equation: OH−+HSO−4⇌ Which species acts as acid in the reaction above? Identify the conjugate acid of OH−. Identify the substance in the reaction that is amphiprotic.

a. Kc=[PCl3][Cl2]/[PCl5] b. Kc=[NH3]2/[N2][H2]3 c. Kc=[NO2]2/[N2O4] The value of Kc can only be changed by changing the temperature of the system. Any other change will be followed by a response from the equilibrium just so that eventually Kc will get the same value as it had before.

Deduce the equilibrium constant expression Kc for the following reactions: a. PCl5⇌PCl3+Cl2 b. N2+3H2⇌2NH3 c.N2O4⇌2NO2 How can the value of Kc be changed?

Kc=[C]c×[D]d / [A]a×[B]b

Derive the equilibrium constant (Kc) expression from the balanced reaction equation below: aA+bB⇌cC+dD

Le Châtelier's principle states that when a change is applied to a system at equilibrium, the system will shift the position of the equilibrium in order to fight that change. In other words, the system will try to offset changes that are made to the system's conditions.

Describe Le Châtelier's principle.

⇌ You must use a double arrow only for reactions between a weak acid and a weak base. This is to indicate it is a reversible reaction in which the backward reaction also occurs. HCOOH(weak)+NH3(weak)⇌HCOO−+NH+4 → When either the acid or the base is strong - or both acid and base are strong - you must use a single arrow. This is to indicate the reaction goes to completion. The backward reaction does not take place. HCl(strong)+NH3(weak)→Cl−+NH+4

Describe when to use a double arrow or a single arrow in an acid-base reaction and give one example of each.

Position of an equilibrium mixture This concept indicates the relative amounts of reactants and products in a system. If the position of an equilibrium mixture lies to the right, most of the particles in the mixture are products (on the right-hand side of the reaction equation) - and relatively few of them are reactants. Shift of an equilibrium This concept indicates that essentially there has been a shift in position of the equilibrium mixture. Basically it means that there is a change in the relative amounts of reactants and products in the mixture.

Explain the difference between the position of an equilibrium mixture and the shift of an equilibrium.

(a) Kc = [HI(g)]2 / [H2(g) ] [I2(g)] (b) The value of Kc indicates that the concentration of product is greater than that of the reactants. (c) An increase in pressure will have no effect on this equilibrium. (d) The forward reaction is endothermic and so an increase in temperature will favour this reaction, the proportion of products will increase as the equilibrium position shifts to the right.

For the reversible reaction: H2(g) + I2(g) ⇋ 2HI(g) ΔH > 0 the equilibrium constant Kc = 60 at a particular temperature. (a) Give the equilibrium expression. (1) (b) For this reaction, what information does the value of Kc provide about the relative concentrations of the product and reactants at equilibrium? (1) (c) What effect, if any, will an increase in pressure have on the equilibrium position? (1) (d) Explain why an increase in temperature increases the value of the equilibrium constant for the above reaction. (1)

Q is the equilibrium expression with non-equilibrium concentrations, so the expression of Q is deduced from reaction equations, just like expressions of Kc. For example: NaOH+CO2⇌NaHCO3 Q=[NaHCO3][NaOH][CO2] If we know the value of Q, we can estimate the position of the system, just like we can for Kc. When the value of Q equals the value of Kc the system is at equilibrium. When Q and Kc have different values, we can use these values to deduce whether the system approaches equilibrium from the left side or from the right side of the reaction equation.

Give an example of an expression of the reaction quotient Q and describe what it can be used for.

H2CO3 NO−3 CH3COOH NH−2

Give the conjugate acid of HCO−3. Give the conjugate base of HNO3. Give the conjugate acid of CH3COO−. Give the conjugate base of NH3.

1. Given: HI⇌1/2 H2+1/2 I2 with Kc=0.01 Rule: Doubling the coefficients in the reaction leads to squaring of Kc. For: 2HI⇌H2+I2 the value is: Kc=(0.01)2 Kc=0.0001 2. Given: 2HI⇌H2+I2 with Kc=0.0001 Rule:Inversing the reaction leads to inverting of Kc. For: H2+I2⇌2HI the value is: Kc=10.0001 Kc=10,000or1×104

Given the reaction: HI⇌1/2 H2+1/2 I2 with Kc = 0.01 Calculate Kc for the following reactions at the same temperature: 2HI⇌H2+I2 H2+I2⇌2HI

(a)A strong acid is fully ionized or dissociated in solution.HCl(aq) → H+( aq) + Cl−(aq); A weak acid is only partly ionized or dissociated in solution.H2CO3(aq) ⇌ 2H+(aq) + CO32−(aq) (b) Test the electrical conductivity of the solution; HCl will have higher conductivity, H2CO3 will have lower conductivity. React both with a reactive metal (such as magnesium); HCl will react faster than H2CO3. React with a base and measure the temperature change; HCl will have a higher temperature increase than H2CO3. (c) 1:104 as each increase of a pH unit is a factor of 10 × weaker. (d) i. acid, HCl; conjugate base, Cl- ;base, HCO3−; conjugate acid, H2CO3 ii. Brønsted-Lowry theory

Hydrochloric acid (HCl) is a strong acid, while carbonic acid (H2CO3) is a weak acid. (a) Explain, with the help of equations, what is meant by strong and weak acid using the above acids as examples. (4) (b) Outline two ways, other than using pH, in which you could distinguish between carbonic acid and hydrochloric acid of the same concentration. (2) (c) A solution of hydrochloric acid, HCl(aq), has a pH of 1 and a solution of carbonic acid, H2CO3(aq), has a pH of 5. Determine the ratio of the hydrogen ion concentrations in these solutions. (2) (d) The relative strengths of the two acids can be illustrated by the following equation: HCO3−(aq) + HCl(aq) ⇌ H2CO3(aq) + Cl−(aq) i. Identify the acid and its conjugate base in this equation; likewise identify the base and its conjugate acid. (2) ii. Name the theory of acids and bases that is illustrated in d-i. (1)

There is still so much alcohol in the bottle because the liquid alcohol and gaseous alcohol are in dynamic equilibrium with each other: Alcohol(l)⇌Alcohol(g) The liquid alcohol evaporates easily, but can't escape from the closed system. Gaseous alcohol will condense back to liquid alcohol and a dynamic equilibrium will establish itself.

If alcohol evaporates so quickly that you can feel and see it evaporate from your hands, how come there's still so much alcohol in the bottle? Explain your answer.

No, an equilibrium shift to the right does not necessarily mean that the new position of the equilibrium lies to the right as well.

Imagine an equilibrium shift to the right. Will the new position of the equilibrium always lie to the right? Explain your answer.

C. HSO -(aq) + CH COOH(aq) → H SO (aq) + CH COO-(aq) 43243

In which reaction is HSO -(aq) acting as a Brønsted - Lowry base? 4 A. HSO -(aq) + NH (aq) → SO 2-(aq) + NH +(aq) 4344 B. HSO -(aq) + OH-(aq) → SO 2-(aq) + H O(l) 442 C. HSO -(aq) + CH COOH(aq) → H SO (aq) + CH COO-(aq) 43243 D. HSO-(aq) + CHNH(aq) → SO2-(aq) + CHNH+(aq) 432433

C. K-1c

Kc is the equilibrium constant for the reaction between hydrogen and iodine: H2(g) + I2(g) ⇌ 2HI(g) ∆H = - 10.4 kJ What will be the value of the equilibrium constant for the reverse reaction? A. Kc B. Kc1⁄2 C. K-1c D. Kc2

A. I and II only

One of the steps in the manufacture of nitric acid is the oxidation of ammonia according to the following equation. 4NH3(g) + 5O2(g) ⇌ 4NO2(g) + 6H2O(g) ∆H = - 901 kJ Which conditions will increase the yield of nitrogen dioxide in the equilibrium mixture? I. Removing the water vapour II. Decreasing the temperature III. Increasing the pressure A. I and II only B. I and III only C. II and III only D. I, II and III

D. Both solutions produced an equal amount of gas.

Separate 1.0 mol dm-3 solutions of hydrochloric acid and ethanoic acid reacted completely with 10.0 g of calcium carbonate to give carbon dioxide. Which is a correct statement? A. Both acids also gave hydrogen as a product. B. The hydrochloric acid solution reacted faster as it has a higher pH. C. The ethanoic acid solution gave a larger amount of gas. D. Both solutions produced an equal amount of gas.

Observation 1: The solution becomes warmer. Observation 2: If an indicator is present, the indicator changes color at certain point(s). Observation 3: Bubbles can be released. When carbonates or hydrogen carbonates react with acid, these bubbles are formed by gaseous CO2 molecules. When reactive metals (those above Cu in the reactivity series, see topic 9) react with acid, these bubbles are formed by gaseous H2 molecules

State three possible observations you can make when acids and bases react.

B. The reaction is exothermic.

The equilibrium constant Kc is 160 at 298 K for the following reaction. H2(g) + I2(g) ⇌ 2HI(g) At 700 K the value for the equilibrium constant is 54. What can be deduced from this information? A. The position of equilibrium lies on the right at all temperatures. B. The reaction is exothermic. C. The rate of the reaction is higher at higher temperatures. D. Hydrogen iodide will decompose at low temperatures

C. propanoic acid (Ka = 1.37 x 10-5)

The equilibrium constant for a weak acid dissociating into its ions is known as the acid dissociation constant and has the symbol Ka. Which is the weakest acid? A. butanoic acid (Ka = 1.51 x 10-5) B. 2-methylpropanoic acid (Ka = 1.41 x 10-5) C. propanoic acid (Ka = 1.37 x 10-5) D. ethanoic acid (Ka = 1.74 x 10-5)

D. [CH4(g)] x [H2Og)] / [H2(g)]3 x [CO(g)] will be a constant value.

The hydrogen needed to make ammonia can be obtained from natural gas by the following reaction. CH4(g) + H2O(g) ⇌ CO(g) + 3H2(g) ∆H = + 210 kJ Which statement is always correct when the reaction is at equilibrium? A. ∆H for the reverse reaction is + 210 kJ. B. The mixture will contain three times as much hydrogen as methane. C. The concentrations of carbon monoxide and hydrogen are equal. D. [CH4(g)] x [H2Og)] / [H2(g)]3 x [CO(g)] will be a constant value.

A. [H+(aq)] in A is ten times that in B.

The pH of solution A is 1 and the pH of solution B is 2. Which is a correct statement? A. [H+(aq)] in A is ten times that in B. B. [H+(aq)] in B is ten times that in A. C. [H+(aq)] in A is twice that in B. D. [H+(aq)] in A is half that in B.

(a) i acid X is the least acidic. (b) [H+] greater in acid Y by a factor of 10. (c) acid Y > acid Z > acid X;As the pH value decreases the concentration of hydrogen ions, and hence the conductivity, increases.

The pH values for three solutions of organic acids, all of the same concentration, were measured. acid X pH = 5acid Y pH = 2acid Z pH = 3 (a) Identify which of these solutions is the least acidic.(1) (b) Deduce how the [H+] values compare in solutions of acids Y and Z. (1) (c) Arrange the three acid solutions in decreasing order of electrical conductivity, starting with the greatest conductivity, giving a reason for your choice. (2)

C. The concentration of ethyl ethanoate will have decreased.

The value of Kc for the following reaction is equal to 4 at 298 K. C2H5OH(l) + CH3COOH(l) ⇌ CH3COOC2H5(l) + H2O(l) A. B. C. D. The reaction quotient, Q, for a mixture of ethanol, water, ethyl ethanoate and ethanoic acid is equal to 6 at 298 K at time t. Which statement will be true when the mixture reaches equilibrium at 298 K? A. The concentration of ethanol will have decreased. B. The concentration of ethanoic acid will have decreased. C. The concentration of ethyl ethanoate will have decreased. D. The concentration of water will be unchange

1. TRUE 2. FALSE (Each step in pH equals a factor ten in [H+].) 3. TRUE

True or false? 1. A soft drink with a pH of 3 has a [H+] that is ten thousand times higher than the [H+] of pure water at 298 K. 2. The [H+] of a solution with pH = 8 is three times lower than the [H+] of a solution with pH = 5. 3. The pH of water depends on the temperature.

1. TRUE 2. FALSE (The value of Q approaches the value of Kc when the system goes from a state of disequilibrium to a state of equilibrium. However, at the point of reaching perfect equilibrium, reaction quotient Q ceases to exist and the correct term to use is equilibrium constant Kc.) 3. TRUE

True or false? 1. The position of an equilibrium mixture can lie to the left while the equilibrium shifts to the left at the same time. 2. Q and Kc can have the same value for a given system at given atmospheric conditions. 3. Reactions can take place in a sophisticated, well-monitored system that shows no sign of any measured change.

According to Le Châtelier's principle a system will try to offset changes made to the system's conditions. Change done to the system: Removing SO3 Effect on the system: [SO3] decreases The system can fight against the dropping [SO3] by increasing the forward reaction that produces SO3 and by slowing down the backward reaction that uses SO3. The result of these two opposite changes in reaction rates is that the equilibrium shifts to the right. [SO3] will increase, [SO2] and [O2] will decrease.

Use Le Châtelier's principle to deduce what happens to the position of the equilibrium when SO3 is removed from the equilibrium. 2SO2(g)+O2(g)⇌2SO3(g) ΔH=−197kJ/mol

According to Le Châtelier's principle a system will try to offset changes made to the system's conditions. There are in total 2 + 1 = 3 gas molecules on the reactant side and only 2 gas molecules on the product side of the equation. The system can thus fight against the increase of pressure by increasing the forward reaction and by slowing down the backward reaction; this reduces the amount of gas molecules in the system. The result of these two opposite changes in reaction rates is that the equilibrium shifts to the right. [SO3] will increase, [SO2] and [O2] will decrease.

Use Le Châtelier's principle to deduce what happens to the position of the equilibrium when the equilibrium below is put under extra pressure. 2SO2(g)+O2(g)⇌2SO3(g) ΔH=−197kJ/mol

Adding a catalyst reduces the activation energy and thus speeds up the forward and the backward reactions equally. As a result, it makes the system reach equilibrium faster. However, adding a catalyst does not change the position of the equilibrium.

Use enthalpy diagrams to show how an equilibrium system is affected when a catalyst is added.

Since cars are moving in both directions, the flow of cars is an analogy of forward and backward reactions. Consider this the process: A⇌B.When there are more cars on island A than on island B, the position of the process lies to the side of A.When at any given time there are more cars driving from A to B than vice-versa, the process is shifting to the side of B.Only when the amount of cars driving from A to B equals the amount of cars driving from B to A, is the process in equilibrium and the net amount of cars on A and B remains the same.

Use the idea of two islands A and B connected by a busy bridge as an analogy to explain dynamic equilibrium. Use four statements.

1. Derive the expression of Q from the reaction equation: Q=[CH3OH]/[CO][H2]2 2. Substitute the concentrations in the expression of Q: Q=1.35/(0.25)(0.45)2=26.7 3. Check how the values of Q and Kc compare: Q < Kc. This means that the system is not at equilibrium, otherwise Kcand Q would have (approximately) the same value. Since Q < Kc, the system lies more to the left than the equilibrium that eventually will form later. Thus, the system approaches equilibrium from the left.

Using the data below, determine whether the system is at equilibrium or not. If not, deduce from which side of the equation the system approaches equilibrium. CO+2H2⇌CH3OH [CO]=0.25M [H2]=0.45M [CH3OH]=1.35M Kc=520

A. HCl < CH3COOH < NH3 < KOH

What is the correct order of increasing pH value for 1.0 mol dm-3 solutions of the four substances? A. HCl < CH3COOH < NH3 < KOH B. CH3COOH<HCl<KOH<NH3 C. KOH < NH3 < CH3COOH < HCl D. NH3 < KOH < HCl < CH3COOH

A. It will decrease by 3

What will happen to the value of the pH of 1 cm3 of a solution of a strong base when 999 cm3 of water are added to it? A. It will decrease by 3 B. It will increase by 3 C. It will decrease by a factor of 3 D. It will increase by a factor of 3

B. It will increase

When carbon dioxide dissolves in water the following reaction occurs. CO (g) + H O(l) ⇌ H+(aq) + HCO -(aq) 223What will happen to the pH of the solution if the carbon dioxide escapes as a gas? A. It will decrease B. It will increase C. It will stay the same D. It will become negative

B. I and III only

Which are correct statements about the following reaction? SO2(g) + Cl2(g) ⇌ SO2Cl2(g) ∆H = - 85 kJ I. Increasing the pressure shifts the position of equilibrium to the right II. Increasing the temperature shifts the position of equilibrium to the right III. Increasing the pressure has no effect on the value of the equilibrium constant. A. I and II only B. I and III only C. II and III only D. I, II and III

D. I, II and III

Which are correct statements concerning a chemical reaction in a state of equilibrium? I. The rate of the forward reaction is equal to the rate of the reverse reaction II. The concentrations of reactants and products do not change III. The forward and reverse reactions are still continuing A. I and II only B. I and III only C. II and III only D. I, II and III

D. I, II and III

Which can be used to distinguish between a solution of a strong monoprotic acid and a solution of weak monoprotic acid with the same concentration? I. Add zinc to each solution and measure the rate of formation of gas produced. II. Add sodium hydroxide solution to each solution and measure the temperature change. III. Measure the pH A. I and II only B. I and III only C. II and III only D. I, II and III

D. Increasing the pressure.

Which change will increase the amount of sulfur trioxide in the equilibrium mixture? 2SO2(g) + O2(g) ⇌ 2SO3(g) ∆H = - 197 kJ A. Increasing the temperature. B. Adding vanadium pentoxide, V2O5(s), as a catalyst. C. Doubling the volume of the container. D. Increasing the pressure.

C. Na2O

Which compound will produce a solution with a pH greater than 7 when dissolved in water? A. NaCl B. NaHSO4 C. Na2O D. NaNO3

A. CH3COOH

Which gives an acidic solution when dissolved in water? A. CH3COOH B. CH3CHO C. CH3CH2OH D. CH3COCH3

B. OH- / O2-

Which is a Brønsted - Lowry conjugate acid/ base pair? A. H3O+ / OH- B. OH- / O2- C. H2SO4 / SO42- D. NH+/NH-

B. Kc = [NO(g)]2 x [Cl2(g)] / [NOCl(g)]2

Which is the correct expression for the equilibrium constant for the following reaction? 2NOCl(g) ⇌ 2NO(g) + Cl2(g) A. Kc = [2NOCl(g)] / [2NO(g)] x [Cl2(g)] B. Kc = [NO(g)]2 x [Cl2(g)] / [NOCl(g)]2 C. Kc = [NOCl(g)]2 / [NO(g)]2 x [Cl2(g)] D. Kc = [2NO(g)] x [Cl2(g)] / [2NOCl(g)]

B. HCl, HNO3, H2SO4

Which list contains only strong acids? A. HCl, HNO3, H2CO3 B. HCl, HNO3, H2SO4 C. CH3COOH, H2CO3, H2SO4 D. H2SO4, CH3COOH, H2CO3

C. Ba(OH)2, KOH, NaOH

Which list contains only strong bases? A. Ba(OH)2, CH3NH2, KOH B. NH3, KOH, NaOH C. Ba(OH)2, KOH, NaOH D. NaOH, CH3NH2, NH3

D. Kc<<1

Which must be a correct statement about the value of the equlibrium constant, Kc, for a reversible reaction which hardly proceeds in the forward direction? A. Kc = 0 B. Kc = 1 C. Kc >>1 D. Kc<<1

B. CO(g) + H2O(g) ⇌ CO2(g) + H2(g)

Which of the following equilibria will not be affected if only the pressure is altered? A. 2SO2(g) + O2(g) ⇌ 2SO3(g) B. CO(g) + H2O(g) ⇌ CO2(g) + H2(g) C. PCl5(g) ⇌ PCl3(g) + Cl2(g) D. C2H4(g) + H2O(g) ⇌ C2H5OH(g)

D. I, II and III

Which species act as Brønsted - Lowry acids in the following reactions? NH -(aq) + H O(l) ⇌ NH (aq) + OH-(aq) CH3NH2(aq) + H2O(l) ⇌ CH3NH3+ (aq) + OH-(aq) I. NH3(aq) II. CH3NH3+ (aq) III. H2O(l) A. I and II only B. I and III only C. II and III only D. I, II and III

C. HNO3 + HSO4 -

Which species act as Brønsted - Lowry bases in the reaction between nitric acid and sulfuric acid? HNO + H SO ⇌ H NO + + HSO - A. HNO3 + H2SO4 B. H2SO4 + H2NO3+ C. HNO3 + HSO4 - D. HNO+ + HSO-

A. I and II only

Which statements about an isolated system at equilibrium are always correct? I. Energy cannot enter or leave the system II. None of the products can escape III. The concentrations of the reactants and products are equal A. I and II only B. I and III only C. II and III only D. I, II and III

B. I and III only

Which statements are correct for equal volumes of solutions of carbonic acid, H2CO3(aq), and hydrochloric acid, HCl(aq) with the same concentration? I. HCl is more dissociated than H2CO3. II. HCl will have a higher pH than H2CO3. III. HCl will be a better electrical conductor than H2CO3. A. I and II only B. I and III only C. II and III only D. I, II and III

B. NaCl

Which substance is not a base? A. CuO B. NaCl C. NaHCO3 D. CuCO3

D. Fe

Which substance reacts with dilute nitric acid to produce hydrogen gas? A. FeCO3 B. Fe(OH)3 C. Fe2O3 D. Fe

A. Ag

Which substance will not react with hydrochloric acid to form a salt? A. Ag B. Zn C. CuCO3 D. Ca(OH)2

A. Raising the temperature.

Which will affect the value of Kc in the following reaction? 2NO2(g) ⇌ N2O4(g) ∆H = - 24 kJ A. Raising the temperature. B. Lowering the pressure. C. Adding a catalyst. D. Adding some NO2(g) to the equilibrium mixture.

C. Decreasing the concentration of H+(aq) by a factor of 100.

Which will change the pH of a solution from a value of 2 to a value of 4? A. Doubling the concentration of H+(aq). B. Doubling the concentration of OH-(aq). C. Decreasing the concentration of H+(aq) by a factor of 100. D. Decreasing the concentration of OH-(aq) by a factor of 100.

D. I, II and III

Which will give an increased yield of sulfur trioxide in the equilibrium reaction for the Contact Process? 2SO2(g) + O2(g) ⇌ 2SO3(g) ∆H = - 197 kJ A. Adding more oxygen B. Decreasing the temperature C. Increasing the pressure A. I and II only B. I and III only C. II and III only D. I, II and III

B. Using a suitable catalyst

Which will increase the rate of the following reaction but not affect the value of Kc? 2SO2(g) + O2(g) ⇌ 2SO3(g) ∆H = - 197 kJ A. Decreasing the pressure B. Using a suitable catalyst D. Decreasing the temperature C. Increasing the temperature