18. Acids and Bases HL

State the expression for the ionic product constant of water

K_w is the ionic product constant of water, as expressed by the equation K_w = [H⁺] [OH^-]

            (Its the equilibrium constant for water)

At standard thermochemical conditions, K_w is equal to 1.00 x 10^-14 mol dm^-3 (meaning there are that many moles of ions dissociated per cubic dm of water)

Deduce [H⁺] and [OH^-] for water at different temperatures given K_w values

As temperature increases, the equilibrium shifts in favor of the products, meaning more ions will be dissociated into the water, decreasing pH (even though it will still be considered neutral at that level)

Solve problems involving [H⁺], [OH^-], pH and pOH

Example: What is the concentration of hydroxide for a solution with a [H⁺] of 1.00 x 10^-4

K_w = [H⁺] [OH^-]

1.00 x 10^-14 = (1.00 x 10^-4 x [OH^-])

1.00 x 10^-14 = (1.00 x 10^-4 x [OH^-])

1.00 x 10^-4 =    1.00 x 10^-4

1.0 x 10^-10 = [OH^-]

1.0 x 10^-10 mol dm^-3 [OH^-]

pOH stands for power of hydroxide. Similar to pH, this is a measure of the –log of the concentration of hydroxide in an aqueous solution.

pK_w stands for power of the ionic dissociation constant. At standard thermochemical conditions, this is equal to 14.

pK_w = -logK_w

State the equation for the reaction of any weak acid or weak base with water, and deduce the expressions for K_a and K_b

K_a =  [H⁺][A^-] 

[HA]

Ka is the acid dissociation constant. It shows how many ions in the equilibrium expression [HA] ßà [H⁺][A^-] are dissociated into the water

To find the K_a of a weak acid, use the equation x²

                                                                             a

a is the concentration of the acid

x is the concentration H⁺_. Since these are equal in the weak acid, we can use the concentration squared

pK_a is the power of the acid dissociation constant (-logK_a)

K_b is equal to [BH⁺][OH^-]

                             [B]

K_b is the base dissociation constant. It shows how many of the ions in the equilibrium expression [B]à [BH⁺][OH^-].

To find the K_b of a weak base, use the equation y²

                                                                             b

b is the concentration of the base

y is the concentration OH^-_. Since these are equal in the weak acid, we can use the concentration squared

pK_b is the power of the base dissociation constant (-logK_b)

Solve problems involving the solutions of weak acids and bases using these expressions:

K_a x K_b = K_w

pK_a + pK_b = pK_w

pH + pOH = pK_w

Example: a 0.280 moldm^-3 solution of a weak acid has a pH of 4.67. What is the pK_b?

[H⁺]= 10 ^-4.67 = 2.138 x 10^-5

K_a = x²/a

K_a = (2.138 x 10^-5)²  =  4.570 x 10^-10  =   1.632 x 10^-9

                0. 280                 0.280

pK_a = -logK_a = 8.787

pK_a + pK_b = pK_w

8.737 + pKb = 14.00

pKb = 5.213

Identify the relative strengths of acids and bases using the values of K_a, K_b, pK_a, and pK_b

The higher K_a is, the stronger the acid and the lower K_a is, the weaker the acid

The higher pK_a is, the weaker the acid. The lower pK_a is, the stronger the acid

The higher K­_b­ is, the stronger the base and the lower K_b is, the weaker the base

The higher pK_b is, the weaker the base. The higher pK_b is, the weaker the acid.

Deduce whether salts form acidic, alkaline, or neutral aqueous solutions

Salt- an ionic compound made of the cations from a base and the anions from an acid

Use this chart to help you identify the acidity of your salt

Anion from:		Cation from:		Resulting salt is…
Strong Acid	+	Strong Base	=	NEUTRAL (strong v. strong)
Strong Acid	+	Weak Base	=	SLIGHT ACIDIC (strong v. weak)
Weak Acid	+	Strong Base	=	SLIGHT BASIC (weak v. strong)
Weak Acid	+	Weak Base	=	NEUTRAL (weak v. weak)

Complex ions are very acidic in water. The charge of the metal ions helps dissociate water, then bind to OH^-, leaving excess H⁺ floating in solution. For example:

[Fe(H₂O)₆] ßà [Fe(OH)(H₂O)₅ + H⁺]

Describe the composition of a buffer solution and explain its action.

Buffer- Solution containing a weak acid + conjugate base (OR a weak base + conjugate acid) that resists changes in pH when small amounts of acid/base are added.

Effect of Buffers on equation: HA ßà H⁺ + A^-

            When small amounts of strong acid are added, the equilibrium shifts to the left

When small amounts of strong base are added, the equilibrium shifts to the right because the H⁺ reacts to form H₂O.

Concentrations of acid and base in a buffer must be higher than the strong acid/base added to increase the resistance to change (buffering action).

Examples of buffer solutions

- Ethanoic acid and sodium ethanoate

CH₃COOH ßà H⁺ + [Na]CH₃COO^-

-Ammonia and Ammonium chloride

NH₃ + H₂O ßà NH₄⁺[Cl]

Buffer pH = pKa – log [HA]/ [A^-]

Buffer pH depends on the K_a of the weak acid and the ratio of concentrations

Optimum buffer- Buffer that’s most effective when pH is equal to pKa.

Solve problems involving the composition and pH of a specified buffer system.

(You mostly have to use the K_a /K_b stuff for this)

Sketch the general shapes of graphs of pH against volume for titrations involving strong and weak acids and bases and explain their important features

Important terminology

Equivalence point- the point at which the amount of acid = the amount of base (i.e. when the pH = 7.)

***Any further addition of an acid or a base (whichever is NOT part of the buffer solution) will à quick jump to an extreme acidity or an extreme alkalinity

Intercept with pH axis- the initial pH of the solution

Buffer action- When something is being added to a weak acid/base, there is a gradual decline to the equivalence point. This is called the buffer action because the concentrated weak acid/base is preventing drastic changes in pH.

Buffer region- Where the buffer action takes place

Read about titration curves at this link:
http://www.ausetute.com.au/titrcurv.html

Describe the qualitative features of an acid-base indicator.

Indicator- substance that has a different color in acidic and basic solutions, so it can show the end point of an acid-base titration

Indicator is often a weak organic acid/base

Color changes because of a shift in the equilibrium HIn ßà H⁺ + In^-

(where In = indicator)

Acid additions shift the equilibrium left because the excess H⁺ bind to the In^-

Base additions shift the equilibrium right because the lack of H⁺ means more In^- is left in the ionic form.

With an Indicator, Ka = [H^-] [In^-]/ [HIn]

State and explain how the pH range of an acid-base indicator relates to its pK_a value

Each indicator changes color over a different range of pHs.

Middle point of an indicator depends on its pKa

When pH = pKa, the indicator will be in the middle of its color change (as the concentrations are equal on both sides of the equilibrium)

Most indicators change over a region of 2 pH units

REMEMBER* pH is logarithmic so it goes by a unit of 10

REMEMBER**[HIn] additions make it acidic

So, if [HIn] is 10x greater than [In], the pH = pKa – 1

So, if [Hin] is 100x greater than [In], the pH = pKa – 2

Etc etc…

REMEMBER***[In^-] additions make it basic

So, if [In^-] is 10x greater than [HIn], the pH = pKa + 1

So, if [In^-] is 100x greater than [HIn], the pH = pKa + 2

Etc etc…

Identify the appropriate indicator for a titration, given the equivalence point of the titration and the pH range of the indicator

There’s info in the data booklet to help with this. Search for something with a  pKa that matches the equivalence point