ACER 10/28 Lecture #18: Introduction to Acids, Bases, and Buffer Solutions

Describe what happens when acids and bases combine (What is this reaction called)?

Neutralization reaction: acid + base → salt + water

What is K_a?

How is it calculated? Include a reaction formula and a equation.

K_a is the acid dissociation constant (or acidity dissociation constant).

According to the formula where an acid reacts with water to form hydronium and a conjugate base:

• HA(aq) + H₂O(l) ⇌ A^-(aq) + H₃O⁺(aq)

• K_a = [A^-][H₃O⁺]/[HA]

What is K_b?

How is it calculated? Include a reaction formula and a equation.

K_b is the base dissociation constant.

According to the formula where a base reacts with water to form hydroxide and a conjugate acid:

• B:(aq) + H₂O(l) ⇌ H:B⁺(aq) + OH^-(aq)

• K_b = [H:B⁺][OH^-]/[B:]

What is pK_a? How is it calculated?

What is pK_b? How is it calculated?

pK_a = -log[K_a]

pK_b = -log[K_b]

How is reaction different for strong acids/bases and weak acids/bases?

Discuss K and pK.

Strong acids/bases completely dissociate in solution.

• K_a (for acids) or K_b (for bases) is very large, so pK_a and pK_b are difficult to calculate.

Weak acids/bases only partially dissociate.

• K_a and K_b can be calculated using the equilibrium concentrations of the acids/bases, conjugate bases/acids, and hydronium H₃O⁺/hydroxide OH^-.

When a strong acid and strong base are mixed, the pH of resulting solution…

is neutral, with pH = 7

What is pH? What is the value of pH for acids versus bases versus neutral solutions?

pH is the scale we use to measure acidity or basicity.

pH is below 7 for acids and above 7 for bases. A pH of 7 is neutral.

How do you calculate pH?

Explain the two versions of the equation.

pH = -log[H₃O⁺] or -log[H⁺]

• These are the same things; H⁺ doesn’t exist in aqueous solutions and will become H₃O⁺.

How do you determine sig figs for pH (for logs in general)?

the number to the left of the decimal point tells you the power of 10 and is NOT included in sig figs

Only the numbers to the right of the decimal are sig figs.

Describe/compare the three definitions of acids/bases.

Click hint after.

Do bases/acids have to fit all three definitions of acids/bases?

Include an example.

Brønsted-Lowry definition: An acid is a proton donor. A base is a proton acceptor.

Lewis definition: An acid is an electron pair acceptor. A base is an electron pair donor.

Arrhenius definition: An acid generates hydronium (H₃O⁺). A base generates hydroxide (OH^-).

• The strict definition states that an acid has to contain H⁺ and a base has to contain OH^-.

A base/acid does not have to fit all three definitions. For example, ammonia NH₃ can accept a proton and donate an electron pair, so it fits the Brønsted-Lowry and Lewis definition of a base. However, it is only an Arrhenius base based on the broad definition (it generates OH^-) but not by the strict definition (it does not contain OH^-).

After a neutralization reaction:

a strong acid becomes…

a strong base becomes…

a weak acid becomes…

a weak base becomes…

Strong acids will become weak conjugate bases.

Strong bases will become weak conjugate acids.

Weak acids will become strong conjugate bases.

Weak bases will become strong conjugate acids.

Water is amphiprotic. What does this mean?

Include a reaction formula and an equation for equilibrium constant. What is the constant called?

Water is both an acid and a base. This means that it can ionize itself. This process is called autoprotolysis (or autoionization).

H₂O (l) + H₂O (l) → H₃O⁺ (aq) + OH^- (aq)

At 298 K, K_w = [H₃O⁺][OH^-] = 1 x 10^-14.

• K_w is the autoionization constant, aka the ion-product constant.

What equation represents pK_w? What is pK_w at 25 C?

pK_w = -log(K_w) = pH + pOH = 14

Relate pK_w to pK_a and pK_b. (2 equations)

Under what conditions is this true?

K_w = K_aK_b

pK_w = pK_a + pK_b = 14

• This equation is only true at 298 K or 25°C where pK_w = 14 and K_w = 10^-14.

What is the relationship between the strength of an acid/base, the K_a or K_b value, and the pK_a or pK_b value?

The weaker the acid, the lower the K_a value, and the higher the pK_a value.

The weaker the base, the lower the K_b value, and the higher the pK_b value.

What is a buffer solution? What is an example of a buffer?

A buffer is a solution that can resist small changes to pH (small additions of H⁺/OH^- will not change pH).

Blood is a buffer.

What is needed to create a good buffer solution?

The best buffers…

A significant amount of a weak acid and some of its conjugate base (or weak base and conjugate acid).

The best buffers have an equal amount of the weak acid and conjugate base (or weak base and conjugate acid).

How do buffers work?

Buffers work by applying Le Châtelier’s Principle to weak acid/base equilibrium.

When acid (or H⁺) is added:

• In a buffer made of weak acid and conjugate base, the conjugate base reacts with the acid to neutralize it.

• In a buffer made of weak base and conjugate acid, the weak base reacts with the acid to neutralize it.

When base (or OH^-) is added:

• In a buffer made of weak acid and conjugate base, the weak acid reacts with the base to neutralize it.

• In a buffer made of weak base and conjugate acid, the conjugate acid reacts with the base to neutralize it.

Describe the common ion effect.

Optional: Include one example.

Equilibrium shifts when an ion in the solution (a "common ion") is added from another source, following Le Chatelier’s principle.

For HA(aq) + H₂O(l) ⇌ A⁻(aq) + H₃O⁺(aq):

• Adding more A^- ions will shift equilibrium to the left. This causes the pH to increase and lowers the H₃O⁺ (or H⁺) ion concentration.

• Adding the acid HA will shift equilibrium to the right. This causes the pH to decrease and raises the H₃O⁺ (or H⁺) ion concentration.

Describe the Henderson-Hasselbalch equations.

Before using these equations, check if…

pH = pK_a + log([A^-]/[HA])

pOH = pK_b + log([B:H⁺]/[B:])

The concentrations used are equilibrium concentrations. Check if “x is small” approximation applies (when K_a or K_b is extremely small or when the initial concentrations are relatively large). If not, the ICE table must be used to determine the concentrations.