5B4 Acid-Base Equilibrium

Question 1

Define:

autoionization

Answer 1

Self-ionization, where water molecules react with each other to form ions.

Autoionization of water results in hydronium (H₃O⁺) and hydroxide (OH⁻) ions in a 1:1 molar ratio.

These ions form when water molecules act as both acids and bases, donating and accepting protons.

Question 2

Fill in the blank:

The chemical equation for the autoionization of water is _______.

Answer 2

H₂O + H₂O ⇌ H₃O⁺ + OH⁻

This equation represents the reversible reaction of water molecules forming hydronium and hydroxide ions.

Question 3

True or False:

The concentration of H₃O⁺ and OH⁻ is equal in pure water.

Answer 3

True

In pure water, both ions are present in equal concentrations of 10⁻⁷ M at 25°C.

Question 4

Define:

Kw

Answer 4

It is the autoionization constant of water, equal to [H₃O⁺][OH⁻].

It is calculated by multiplying the concentrations of H₃O⁺ and OH⁻ ions.

For pure water at 25°C: [H₃O⁺][OH⁻] = (10⁻⁷)(10⁻⁷) = 10⁻¹⁴.

Question 5

How does the concentration of H₃O⁺ affect the pH of a solution?

Answer 5

Higher H₃O⁺ concentration decreases pH, making the solution acidic.

pH is the negative logarithm of H₃O⁺ concentration: pH = -log[H₃O⁺].

Question 6

Fill in the blank:

When the concentration of OH⁻ increases, the solution becomes more _____.

Answer 6

basic

Basic solutions have higher OH⁻ concentrations and lower H₃O⁺ concentrations.

Question 7

What is the relationship between pKw and Kw?

Answer 7

pKw = -log(Kw)

At 25°C, pKw = 14.

Question 8

How is pKw calculated?

Answer 8

pKw = pH + pOH

At 25°C, pH + pOH = 14.

Question 9

Give a factor that affect the value of Kw.

Answer 9

Temperature

Kw is temperature-dependent and increases with rising temperature.

Question 10

How is the pH scale range determined?

Answer 10

By taking the negative logarithm of Kw.

Kw = 10⁻¹⁴; pKw = 14, so the pH scale generally ranges from 0 to 14.

Question 11

What happens to the concentrations of H₃O⁺ and OH⁻ when an acid is added to water?

Answer 11

H₃O⁺ concentration increases, OH⁻ concentration decreases.

Le Chatelier’s principle explains the shift in equilibrium to restore the product [H₃O⁺][OH⁻] = 10⁻¹⁴.

Question 12

What is the pH of a solution where [H₃O⁺] = 1 × 10⁻⁵ M?

Answer 12

pH = 5

pH = -log[H₃O⁺], so pH = -log(1 × 10⁻⁵) = 5.

Question 13

True or False:

The product [H₃O⁺][OH⁻] is always 10⁻¹⁴, regardless of solution conditions.

Answer 13

False

While Kw = 10⁻¹⁴ at 25°C, Kw varies with temperature.

Question 14

List the ions present in a basic solution.

Answer 14

• Hydroxide ions (OH⁻)
• Some hydronium ions (H₃O⁺)

In a basic solution, OH⁻ ions outnumber H₃O⁺ ions, but both are present.

Question 15

What is a dissociation constant?

Answer 15

It is the equilibrium constant for the dissociation of a substance in a solution.

The dissociation constant indicates how much of the substance dissociates into its ions at equilibrium.

Question 16

What is Ka in chemistry?

Answer 16

Ka is the acid dissociation constant.

It is a measure of how much an acid dissociates in solution.

Question 17

How is Ka used in acid-base equilibria?

Answer 17

Ka helps determine the strength of an acid and its dissociation in a solution.

Ka is a critical factor in determining how much an acid dissociates in water and its behavior in acid-base reactions.

Question 18

True or False:

Strong acids always have Ka > 1.

Answer 18

True

A Ka > 1 indicates complete dissociation, characteristic of strong acids.

Question 19

How is Ka calculated?

Answer 19

Ka = [H⁺][A⁻] / [HA]

This formula uses the concentrations of hydrogen ions (H⁺), conjugate base (A⁻), and the undissociated acid (HA).

Question 20

What is the difference between strong and weak acids based on Ka?

Answer 20

• Strong acids: They have high Ka values.
• Weak acids: They have low Ka values.

Strong acids dissociate completely, while weak acids dissociate partially, leading to different Ka values.

Question 21

What is the relationship between Ka and pKa?

Answer 21

pKa = -log(Ka)

pKa is the negative logarithm of Ka, providing a different scale to express the strength of acids.

Question 22

What is the Ka expression for acetic acid?

Answer 22

Ka = [H⁺][C₂H₃O₂⁻] / [HC₂H₃O₂]

The Ka measures the extent of acid dissociation in water.

Question 23

Calculate the Ka for acetic acid if the concentration of H⁺ is 0.0006 M and that of the acid is 1.2 M.

Answer 23

Ka = 3.0 x 10⁻⁷

This calculation uses the concentrations of H⁺ and acetic acid at equilibrium to find the Ka value.

Ka = (0.0006)(0.0006) / 1.2 = 3.0 x 10⁻⁷

Question 24

What is the relationship between pH and Ka for weak acids?

Answer 24

A lower pH indicates a stronger acid with a higher Ka.

pH measures the concentration of hydrogen ions, which is related to the dissociation constant of an acid.

Question 25

What does **Kb represent** in chemistry?

Answer 25

Kb is the **base dissociation constant**, a measure of how much a base dissociates. ## Footnote Strong bases have high Kb values, while weak bases have low Kb values.

Question 26

How is **Kb calculated**?

Answer 26

Kb = [BH⁺][OH⁻] / [B] ## Footnote Kb is calculated similarly to Ka, using the concentrations of the conjugate acid (BH⁺), hydroxide ions (OH⁻), and the base (B).

Question 27

How is **Kb related to Ka** for conjugate acid-base pairs?

Answer 27

Kb = Kw / Ka ## Footnote The product of Ka and Kb for a conjugate acid-base pair is equal to the ionization constant of water (Kw), which is 1.0 x 10⁻¹⁴ at 25°C.

Question 28

What is the pKa of a weak acid with a **Ka of 1.0 x 10⁻⁵**?

Answer 28

pKa = 5 ## Footnote pKa is the negative logarithm of Ka, so pKa = -log(1.0 x 10⁻⁵) = 5.

Question 29

# Which acid **dissociates completely** in water: HCl or CH₃COOH

Answer 29

HCl ## Footnote HCl is a strong acid, while acetic acid (CH₃COOH) is a weak acid that dissociates partially.

Question 30

# Define: buffer solution

Answer 30

It is a solution that **resists changes in pH**. ## Footnote A buffer is made of a weak acid and its conjugate base or vice versa. Buffers help maintain stable pH levels in chemical and biological systems.

Question 31

What type of acids and bases **cannot form buffers**?

Answer 31

**Strong** acids and bases. ## Footnote Strong acids/bases completely dissociate in water, leaving no conjugate acid/base to establish equilibrium.

Question 32

# True or False: A buffer is **most effective** when the acid-to-base ratio is 1:1.

Answer 32

True ## Footnote At a 1:1 ratio, the buffer has maximum capacity to neutralize added acid or base.

Question 33

What does the **Henderson-Hasselbalch equation** calculate?

Answer 33

The **pH of a buffer** solution. ## Footnote It uses pKa and the ratio of the concentrations of conjugate acid and base.

Question 34

What is the **Henderson-Hasselbalch equation**?

Answer 34

pH=pKa+log⁡([base]/[acid]) ## Footnote This equation relates pH, pKa, and the ratio of base to acid concentrations in a buffer solution.

Question 35

What is the pH of a buffer with **equal concentrations of acid and base**?

Answer 35

The pH **is equal** to the pKa. ## Footnote At a 1:1 ratio, log⁡(1)=0, so pH=pKa.

Question 36

Which principle explains **why buffers resist pH changes**?

Answer 36

Le Chatelier's principle ## Footnote Adding acid or base shifts the equilibrium to minimize changes in hydrogen ion concentration.

Question 37

# True or False: A strong base can be used to create a buffer **by reacting with a weak acid**.

Answer 37

True ## Footnote This reaction forms the conjugate base of the weak acid, which helps maintain buffer equilibrium.

Question 38

How does the pH change when there is **more base than acid** in a buffer?

Answer 38

The pH becomes **higher than the pKa**. ## Footnote The logarithmic term in the Henderson-Hasselbalch equation increases the pH.

Question 39

# True or False: A buffer solution can have a **neutral pH**.

Answer 39

False ## Footnote Buffers never have a neutral pH because they consist of a weak acid/base pair.

Question 40

What happens when a **strong acid is added to a buffer**?

Answer 40

The **buffer neutralizes the acid** by reacting with the conjugate base. ## Footnote This minimizes the change in pH.

Question 41

# Define: buffer capacity

Answer 41

The **amount of acid or base a buffer can neutralize** before its pH changes significantly. ## Footnote Buffer capacity depends on the concentrations of the buffer components.

Question 42

Which **graph feature** indicates the buffer region during titration?

Answer 42

A **flat slope** on the titration curve. ## Footnote The flat slope corresponds to the buffer's resistance to pH change.