Reactivity 2.3—How far? The extent of chemical change

Deduce the expression for the equilibrium constant, K_c, for this equation.

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for this reaction.

What is the equilibrium constant expression for the following equation?

2NO₂ (g) + F₂ (g) $\rightleftharpoons$ 2NO₂F (g)

A. $\frac{2[\text{N}{\text{O}}_2\text{F]}}{2[\text{N}{\text{O}}_2]+[{\text{F}}_2]}$

B. $\frac{2[\text{N}{\text{O}}_2\text{F]}}{2[\text{N}{\text{O}}_2][{\text{F}}_2]}$

C. $\frac{{[\text{N}{\text{O}}_2\text{]}}^2\text{[}{\text{F}}_2\text{]}}{{[\text{N}{\text{O}}_2\text{F}]}^2}$

D. $\frac{{[\text{N}{\text{O}}_2\text{F]}}^2}{{[\text{N}{\text{O}}_2]}^2[{\text{F}}_2]}$

K_c for 2N₂O (g) $\rightleftharpoons$ 2N₂ (g) + O₂ (g) is 7.3 × 10³⁴.

What is K_c for the following reaction, at the same temperature?

N₂ (g) + $\frac{1}{2}$O₂ (g) $\rightleftharpoons$ N₂O (g)

A. 7.3 × 10³⁴

B. $\frac{1}{\sqrt{7.3\times {10}^{34}}}$

C. $\frac{1}{7.3\times {10}^{34}}$

D. $\frac{1}{2\times 7.3\times {10}^{34}}$

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for the reaction above.

Deduce the expression for the equilibrium constant, K_c, for this equation.

Deduce the expression for the equilibrium constant, K_c, for this equation.

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for this reaction.

What is the equilibrium constant expression for the following equation?

2NO₂ (g) + F₂ (g) $\rightleftharpoons$ 2NO₂F (g)

A. $\frac{2[\text{N}{\text{O}}_2\text{F]}}{2[\text{N}{\text{O}}_2]+[{\text{F}}_2]}$

B. $\frac{2[\text{N}{\text{O}}_2\text{F]}}{2[\text{N}{\text{O}}_2][{\text{F}}_2]}$

C. $\frac{{[\text{N}{\text{O}}_2\text{]}}^2\text{[}{\text{F}}_2\text{]}}{{[\text{N}{\text{O}}_2\text{F}]}^2}$

D. $\frac{{[\text{N}{\text{O}}_2\text{F]}}^2}{{[\text{N}{\text{O}}_2]}^2[{\text{F}}_2]}$

K_c for 2N₂O (g) $\rightleftharpoons$ 2N₂ (g) + O₂ (g) is 7.3 × 10³⁴.

What is K_c for the following reaction, at the same temperature?

N₂ (g) + $\frac{1}{2}$O₂ (g) $\rightleftharpoons$ N₂O (g)

A. 7.3 × 10³⁴

B. $\frac{1}{\sqrt{7.3\times {10}^{34}}}$

C. $\frac{1}{7.3\times {10}^{34}}$

D. $\frac{1}{2\times 7.3\times {10}^{34}}$

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for this reaction.

State the equilibrium constant expression, K_c, for the reaction above.

State the equilibrium constant expression, K_c, for the reaction above.

Determine the enthalpy change, ΔH, for the Haber–Bosch process, in kJ. Use Section 11 of the data booklet.

Demonstrate that your answer to (b)(i) is consistent with the effect of an increase in temperature on the percentage yield, as shown in the graph.

What is the strongest acid in the equation below?

H₃AsO₄ + H₂O $\rightleftharpoons$ H₂AsO₄⁻ + H₃O⁺      K_c = 4.5 × 10⁻⁴

A.  H₃AsO₄

B.  H₂O

C.  H₂AsO₄⁻

D.  H₃O⁺

The equilibrium 2H₂ (g) + N₂(g) $\rightleftharpoons$ N₂H₄(g) has an equilibrium constant, K, at 150 °C. 

What is the equilibrium constant at 150 °C, for the reverse reaction?

N₂H₄(g) $\rightleftharpoons$ 2H₂ (g) + N₂(g)

A.  K

B.  K⁻¹

C.  −K

D.  2K

Determine the enthalpy change, ΔH, for the Haber–Bosch process, in kJ. Use Section 11 of the data booklet.

Demonstrate that your answer to (b)(i) is consistent with the effect of an increase in temperature on the percentage yield, as shown in the graph.

Determine the enthalpy change, ΔH, for the Haber–Bosch process, in kJ. Use Section 11 of the data booklet.

Demonstrate that your answer to (b)(i) is consistent with the effect of an increase in temperature on the percentage yield, as shown in the graph.

What is the strongest acid in the equation below?

H₃AsO₄ + H₂O $\rightleftharpoons$ H₂AsO₄⁻ + H₃O⁺      K_c = 4.5 × 10⁻⁴

A.  H₃AsO₄

B.  H₂O

C.  H₂AsO₄⁻

D.  H₃O⁺

The equilibrium 2H₂ (g) + N₂(g) $\rightleftharpoons$ N₂H₄(g) has an equilibrium constant, K, at 150 °C. 

What is the equilibrium constant at 150 °C, for the reverse reaction?

N₂H₄(g) $\rightleftharpoons$ 2H₂ (g) + N₂(g)

A.  K

B.  K⁻¹

C.  −K

D.  2K

The dissociation of citric acid is an endothermic process. State the effect on the hydrogen ion concentration, [H⁺], and on the equilibrium constant, of increasing the temperature.

When a bottle of carbonated water is opened, these equilibria are disturbed.

State, giving a reason, how a decrease in pressure affects the position of Equilibrium (1).

Predict, referring to Equilibrium (2), how the added sodium hydrogencarbonate affects the pH.(Assume pressure and temperature remain constant.)

Comment on why peracetic acid, CH3COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) ⇌ CH₃COOOH (aq) + H₂O (l)

Calculate the concentration of H₃PO₄ if 25.00 cm³ is completely neutralised by the addition of 28.40 cm³ of 0.5000 mol dm⁻³ NaOH.

When a bottle of carbonated water is opened, these equilibria are disturbed.

State, giving a reason, how a decrease in pressure affects the position of Equilibrium (1).

Predict, referring to Equilibrium (2), how the added sodium hydrogencarbonate affects the pH.(Assume pressure and temperature remain constant.)

Comment on why peracetic acid, CH₃COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) $\rightleftharpoons$ CH₃COOOH (aq) + H₂O (l)

Consider the following equilibrium reaction.

2N₂O (g) + O₂ (g) $\rightleftharpoons$ 4NO (g)        ΔH = +16 kJ

Which change will move the equilibrium to the right?

A. Decrease in pressure

B. Decrease in temperature

C. Increase in [NO]

D. Decrease in [O₂]

A solution of bleach can be made by reacting chlorine gas with a sodium hydroxide solution.

Cl2 (g) + 2NaOH (aq) $\rightleftharpoons$ NaOCl (aq) + NaCl (aq) + H₂O (l)

Suggest, with reference to Le Châtelier’s principle, why it is dangerous to mix vinegar and bleach together as cleaners.

The dissociation of citric acid is an endothermic process. State the effect on the hydrogen ion concentration, [H⁺], and on K_a, of increasing the temperature.

Calculate the concentration of H₃PO₄ if 25.00 cm³ is completely neutralised by the addition of 28.40 cm³ of 0.5000 mol dm⁻³ NaOH.

Consider the following equilibrium reaction:

2SO₂(g) + O₂(g) $\rightleftharpoons$ 2SO₃(g)

State and explain how the equilibrium would be affected by increasing the volume of the reaction container at a constant temperature.

What effect does increasing both pressure and temperature have on the equilibrium constant, K_c?

N₂ (g) + 3H₂ (g) $\rightleftharpoons$ 2NH₃ (g)           ΔH = −45.9 kJ

A.  Decreases

B.  Increases

C.  Remains constant

D.  Cannot be predicted as effects are opposite

What is correct when temperature increases in this reaction at equilibrium?

$2\mathrm{NOCl} \left(\mathrm{g}\right)\rightleftharpoons 2\mathrm{NO} \left(\mathrm{g}\right)+{\mathrm{Cl}}_2 \left(\mathrm{g}\right)$    $∆H^{⦵}=+75.5 \mathrm{kJ}$

Which changes produce the greatest increase in the percentage conversion of methane?

CH₄(g) + H₂O (g) $\rightleftharpoons$ CO (g) + 3H₂(g)

The exothermic reaction $\text{I}$₂ (g) + 3Cl₂ (g) $\rightleftharpoons$ 2$\text{I}$Cl₃ (g) is at equilibrium in a fixed volume. What is correct about the reaction quotient, Q, and shift in position of equilibrium the instant temperature is raised?

A.   Q > K, equilibrium shifts right towards products.

B.   Q > K, equilibrium shifts left towards reactants.

C.   Q < K, equilibrium shifts right towards products.

D.   Q < K, equilibrium shifts left towards reactants.

Which of these changes would shift the equilibrium to the right?

[Co(H₂O)₆]²⁺ (aq) + 4Cl⁻ (aq) $\rightleftharpoons$ [CoCl₄]²⁻ (aq) + 6H₂O (l)

I.   Addition of 0.01 M HCl
II.  Addition of concentrated HCl
III. Evaporation of water

A.  I and II only

B.  I and III only

C.  II and III only

D.  I, II and III

The dissociation of citric acid is an endothermic process. State the effect on the hydrogen ion concentration, [H⁺], and on the equilibrium constant, of increasing the temperature.

The dissociation of citric acid is an endothermic process. State the effect on the hydrogen ion concentration, [H⁺], and on the equilibrium constant, of increasing the temperature.

When a bottle of carbonated water is opened, these equilibria are disturbed.

State, giving a reason, how a decrease in pressure affects the position of Equilibrium (1).

Predict, referring to Equilibrium (2), how the added sodium hydrogencarbonate affects the pH.(Assume pressure and temperature remain constant.)

When a bottle of carbonated water is opened, these equilibria are disturbed.

State, giving a reason, how a decrease in pressure affects the position of Equilibrium (1).

Predict, referring to Equilibrium (2), how the added sodium hydrogencarbonate affects the pH.(Assume pressure and temperature remain constant.)

Comment on why peracetic acid, CH3COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) ⇌ CH₃COOOH (aq) + H₂O (l)

Comment on why peracetic acid, CH3COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) ⇌ CH₃COOOH (aq) + H₂O (l)

Calculate the concentration of H₃PO₄ if 25.00 cm³ is completely neutralised by the addition of 28.40 cm³ of 0.5000 mol dm⁻³ NaOH.

Calculate the concentration of H₃PO₄ if 25.00 cm³ is completely neutralised by the addition of 28.40 cm³ of 0.5000 mol dm⁻³ NaOH.

When a bottle of carbonated water is opened, these equilibria are disturbed.

State, giving a reason, how a decrease in pressure affects the position of Equilibrium (1).

Predict, referring to Equilibrium (2), how the added sodium hydrogencarbonate affects the pH.(Assume pressure and temperature remain constant.)

When a bottle of carbonated water is opened, these equilibria are disturbed.

State, giving a reason, how a decrease in pressure affects the position of Equilibrium (1).

Predict, referring to Equilibrium (2), how the added sodium hydrogencarbonate affects the pH.(Assume pressure and temperature remain constant.)

Comment on why peracetic acid, CH₃COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) $\rightleftharpoons$ CH₃COOOH (aq) + H₂O (l)

Comment on why peracetic acid, CH₃COOOH, is always sold in solution with ethanoic acid and hydrogen peroxide.

H₂O₂ (aq) + CH₃COOH (aq) $\rightleftharpoons$ CH₃COOOH (aq) + H₂O (l)

Consider the following equilibrium reaction.

2N₂O (g) + O₂ (g) $\rightleftharpoons$ 4NO (g)        ΔH = +16 kJ

Which change will move the equilibrium to the right?

A. Decrease in pressure

B. Decrease in temperature

C. Increase in [NO]

D. Decrease in [O₂]

A solution of bleach can be made by reacting chlorine gas with a sodium hydroxide solution.

Cl2 (g) + 2NaOH (aq) $\rightleftharpoons$ NaOCl (aq) + NaCl (aq) + H₂O (l)

Suggest, with reference to Le Châtelier’s principle, why it is dangerous to mix vinegar and bleach together as cleaners.

A solution of bleach can be made by reacting chlorine gas with a sodium hydroxide solution.

Cl2 (g) + 2NaOH (aq) $\rightleftharpoons$ NaOCl (aq) + NaCl (aq) + H₂O (l)

Suggest, with reference to Le Châtelier’s principle, why it is dangerous to mix vinegar and bleach together as cleaners.

The dissociation of citric acid is an endothermic process. State the effect on the hydrogen ion concentration, [H⁺], and on K_a, of increasing the temperature.

The dissociation of citric acid is an endothermic process. State the effect on the hydrogen ion concentration, [H⁺], and on K_a, of increasing the temperature.

Calculate the concentration of H₃PO₄ if 25.00 cm³ is completely neutralised by the addition of 28.40 cm³ of 0.5000 mol dm⁻³ NaOH.

Calculate the concentration of H₃PO₄ if 25.00 cm³ is completely neutralised by the addition of 28.40 cm³ of 0.5000 mol dm⁻³ NaOH.

Consider the following equilibrium reaction:

2SO₂(g) + O₂(g) $\rightleftharpoons$ 2SO₃(g)

State and explain how the equilibrium would be affected by increasing the volume of the reaction container at a constant temperature.

Consider the following equilibrium reaction:

2SO₂(g) + O₂(g) $\rightleftharpoons$ 2SO₃(g)

State and explain how the equilibrium would be affected by increasing the volume of the reaction container at a constant temperature.

What effect does increasing both pressure and temperature have on the equilibrium constant, K_c?

N₂ (g) + 3H₂ (g) $\rightleftharpoons$ 2NH₃ (g)           ΔH = −45.9 kJ

A.  Decreases

B.  Increases

C.  Remains constant

D.  Cannot be predicted as effects are opposite

What is correct when temperature increases in this reaction at equilibrium?

$2\mathrm{NOCl} \left(\mathrm{g}\right)\rightleftharpoons 2\mathrm{NO} \left(\mathrm{g}\right)+{\mathrm{Cl}}_2 \left(\mathrm{g}\right)$    $∆H^{⦵}=+75.5 \mathrm{kJ}$

Which changes produce the greatest increase in the percentage conversion of methane?

CH₄(g) + H₂O (g) $\rightleftharpoons$ CO (g) + 3H₂(g)

The exothermic reaction $\text{I}$₂ (g) + 3Cl₂ (g) $\rightleftharpoons$ 2$\text{I}$Cl₃ (g) is at equilibrium in a fixed volume. What is correct about the reaction quotient, Q, and shift in position of equilibrium the instant temperature is raised?

A.   Q > K, equilibrium shifts right towards products.

B.   Q > K, equilibrium shifts left towards reactants.

C.   Q < K, equilibrium shifts right towards products.

D.   Q < K, equilibrium shifts left towards reactants.

Which of these changes would shift the equilibrium to the right?

[Co(H₂O)₆]²⁺ (aq) + 4Cl⁻ (aq) $\rightleftharpoons$ [CoCl₄]²⁻ (aq) + 6H₂O (l)

I.   Addition of 0.01 M HCl
II.  Addition of concentrated HCl
III. Evaporation of water

A.  I and II only

B.  I and III only

C.  II and III only

D.  I, II and III

With reference to the reaction quotient, Q, explain why the percentage yield increases as the pressure is increased at constant temperature.

A reversible reaction has a reaction quotient, Q, of 4.5 and equilibrium constant, K_c, of 6.2.

2A (g) $\rightleftharpoons$ A₂(g)

Which statement describes the reaction at this time?

A.  The system has reached equilibrium.

B.  The rate of the forward reaction is greater than the rate of the reverse reaction.

C.  The concentration of reactant is greater than the concentration of product.

D.  At equilibrium, the concentration of reactant is greater than the concentration of product.

With reference to the reaction quotient, Q, explain why the percentage yield increases as the pressure is increased at constant temperature.

With reference to the reaction quotient, Q, explain why the percentage yield increases as the pressure is increased at constant temperature.

A reversible reaction has a reaction quotient, Q, of 4.5 and equilibrium constant, K_c, of 6.2.

2A (g) $\rightleftharpoons$ A₂(g)

Which statement describes the reaction at this time?

A.  The system has reached equilibrium.

B.  The rate of the forward reaction is greater than the rate of the reverse reaction.

C.  The concentration of reactant is greater than the concentration of product.

D.  At equilibrium, the concentration of reactant is greater than the concentration of product.

0.50 mol of ${\text{I}}_2$ (g) and 0.50 mol of Br₂(g) are placed in a closed flask. The following equilibrium is established.

${\text{I}}_2$(g) + Br₂(g) $\rightleftharpoons$ 2$\text{I}$Br (g)

The equilibrium mixture contains 0.80 mol of $\text{I}$Br (g). What is the value of K_c?

A.  0.64

B.  1.3

C.  2.6

D.  64

Sulfur dioxide reacts with oxygen to form sulfur trioxide.

2SO₂ (g) + O₂ (g) $\rightleftharpoons$ 2SO₃ (g)           ΔH = −197 kJ

Which change increases the value of K_c?

A.  increasing the temperature

B.  decreasing the temperature

C.  decreasing [SO₂ (g)]

D.  decreasing [SO₃ (g)]

SO₂ (g), O₂(g) and SO₃(g) are mixed and allowed to reach equilibrium at 600 °C.

Determine the value of K_c at 600 °C.

At equilibrium, the concentrations of chlorine and iodine are both 0.02 mol dm^–3.

$\frac{1}{2}$Cl₂ (g) + $\frac{1}{2}$${\text{I}}_2$ (g) $\rightleftharpoons$ $\text{I}$Cl (g)       K_c = 454

What is the concentration of iodine monochloride, $\text{I}$Cl?

A.  $\frac{454}{0.02}$

B.  $454\times 0.02$

C.  $\frac{454}{0.04}$

D.  $454\times 0.04$

$\frac{1}{2}$Cl₂(g) + $\frac{1}{2}{\text{I}}_2$(g) $\rightleftharpoons$ $\text{I}$Cl (g)    K_c = 454

What is the K_c value for the reaction below?

2 $\text{I}$Cl (g) $\rightleftharpoons$ Cl₂(g) + ${\text{I}}_{\text{2}}$(g)

A.  $2\times 454$

B.  $\frac{1}{2\times 454}$

C.  ${454}^2$

D.  $\frac{1}{{454}^2}$

Which substance is the reducing agent in the given reaction?

H⁺ (aq) + 2H₂O (l) + 2MnO₄⁻ (aq) + 5SO₂ (g) → 2Mn²⁺ (aq) + 5HSO₄⁻ (aq)

A.  H⁺

B.  H₂O

C.  MnO₄⁻

D.  SO₂

For the reaction $\text{I}$₂(g) + 3Cl₂ (g) $\rightleftharpoons$ 2$\text{I}$Cl₃ (g) at a certain temperature, the equilibrium concentrations are (in mol dm⁻³):

[$\text{I}$₂] = 0.20, [Cl₂] = 0.20, [$\text{I}$Cl₃] = 2.0

What is the value of K_c?

A.  0.25

B.  50

C.  2500

D.  5000

0.50 mol of ${\text{I}}_2$ (g) and 0.50 mol of Br₂(g) are placed in a closed flask. The following equilibrium is established.

${\text{I}}_2$(g) + Br₂(g) $\rightleftharpoons$ 2$\text{I}$Br (g)

The equilibrium mixture contains 0.80 mol of $\text{I}$Br (g). What is the value of K_c?

A.  0.64

B.  1.3

C.  2.6

D.  64

Sulfur dioxide reacts with oxygen to form sulfur trioxide.

2SO₂ (g) + O₂ (g) $\rightleftharpoons$ 2SO₃ (g)           ΔH = −197 kJ

Which change increases the value of K_c?

A.  increasing the temperature

B.  decreasing the temperature

C.  decreasing [SO₂ (g)]

D.  decreasing [SO₃ (g)]

SO₂ (g), O₂(g) and SO₃(g) are mixed and allowed to reach equilibrium at 600 °C.

Determine the value of K_c at 600 °C.

SO₂ (g), O₂(g) and SO₃(g) are mixed and allowed to reach equilibrium at 600 °C.

Determine the value of K_c at 600 °C.

At equilibrium, the concentrations of chlorine and iodine are both 0.02 mol dm^–3.

$\frac{1}{2}$Cl₂ (g) + $\frac{1}{2}$${\text{I}}_2$ (g) $\rightleftharpoons$ $\text{I}$Cl (g)       K_c = 454

What is the concentration of iodine monochloride, $\text{I}$Cl?

A.  $\frac{454}{0.02}$

B.  $454\times 0.02$

C.  $\frac{454}{0.04}$

D.  $454\times 0.04$

$\frac{1}{2}$Cl₂(g) + $\frac{1}{2}{\text{I}}_2$(g) $\rightleftharpoons$ $\text{I}$Cl (g)    K_c = 454

What is the K_c value for the reaction below?

2 $\text{I}$Cl (g) $\rightleftharpoons$ Cl₂(g) + ${\text{I}}_{\text{2}}$(g)

A.  $2\times 454$

B.  $\frac{1}{2\times 454}$

C.  ${454}^2$

D.  $\frac{1}{{454}^2}$

Which substance is the reducing agent in the given reaction?

H⁺ (aq) + 2H₂O (l) + 2MnO₄⁻ (aq) + 5SO₂ (g) → 2Mn²⁺ (aq) + 5HSO₄⁻ (aq)

A.  H⁺

B.  H₂O

C.  MnO₄⁻

D.  SO₂

For the reaction $\text{I}$₂(g) + 3Cl₂ (g) $\rightleftharpoons$ 2$\text{I}$Cl₃ (g) at a certain temperature, the equilibrium concentrations are (in mol dm⁻³):

[$\text{I}$₂] = 0.20, [Cl₂] = 0.20, [$\text{I}$Cl₃] = 2.0

What is the value of K_c?

A.  0.25

B.  50

C.  2500

D.  5000

Calculate the value of the equilibrium constant, K, at 298 K. Use sections 1 and 2 of the data booklet.

Phenylethene is manufactured from benzene and ethene in a two-stage process. The overall reaction can be represented as follows with ΔG^θ = +10.0 kJ mol⁻¹ at 298 K.

Calculate the equilibrium constant for the overall conversion at 298 K, using section 1 of the data booklet.

Calculate the standard Gibbs free energy change, $\mathrm{\Delta }{G}^{\text{θ}}$, in kJ mol⁻¹, for the first dissociation of citric acid at 298 K, using section 1 of the data booklet.

Calculate the value of the equilibrium constant, K, at 298 K. Use sections 1 and 2 of the data booklet.

Calculate the value of the equilibrium constant, K, at 298 K. Use sections 1 and 2 of the data booklet.

Phenylethene is manufactured from benzene and ethene in a two-stage process. The overall reaction can be represented as follows with ΔG^θ = +10.0 kJ mol⁻¹ at 298 K.

Calculate the equilibrium constant for the overall conversion at 298 K, using section 1 of the data booklet.

Phenylethene is manufactured from benzene and ethene in a two-stage process. The overall reaction can be represented as follows with ΔG^θ = +10.0 kJ mol⁻¹ at 298 K.

Calculate the equilibrium constant for the overall conversion at 298 K, using section 1 of the data booklet.

Calculate the standard Gibbs free energy change, $\mathrm{\Delta }{G}^{\text{θ}}$, in kJ mol⁻¹, for the first dissociation of citric acid at 298 K, using section 1 of the data booklet.

Calculate the standard Gibbs free energy change, $\mathrm{\Delta }{G}^{\text{θ}}$, in kJ mol⁻¹, for the first dissociation of citric acid at 298 K, using section 1 of the data booklet.