CHM112 Exam 3

Chap 17 acids & bases. Chap 18 acid-base equilibria & solubility equilibria

Strong acids

H₂SO₄ - Sulfuric

HI - hydroiodic

HBr - hydrobromic

HNO₃ - nitric acid

HCl - hydrochloric

HClO₄ - perchloric

“So I Brought No Clean Clothes”

Strong bases

Hydroxides of group 1 and 2 metals

Most common:

LiOH - lithium hydroxide

NaOH - sodium hydroxide

KOH - potassium hydroxide

Ba(OH)₂ - barium hydroxide

Ca(OH)₂ - calcium hydroxide

Arrhenius definition of acids and bases

Acid: A substance that produces hydrogen ions (H+) (aka
protons) in aqueous solution

Base: A substance that produces hydroxide ions (OH-) in
aqueous solution.

What’s a neutralization reaction?

a reaction between an acid and a base. Usually produces water and a salt. The overall result is the cancellation of the acid's and base's acidic and basic properties

HCl(aq) + NaOH(aq) → H₂O(l) + NaCl(aq)

Brønsted-Lowry definition of acids and bases

Acid is a proton (H+ ion) donor.
Base is a proton (H+ ion) acceptor.

Why is water considered amphiprotic?

Because it can either donate or accept a proton.

H₂O(l) + H₂O(l) ⇆ H₃O⁺(aq) + OH^-(aq)
Acid Base

What is a conjugate acid-base pair?

If two species only differ by the presence or absence of
an H+ ion, they are a conjugate acid-base pair.

NH₃(aq) + H₂O(l) ⇆ NH₄⁺(aq) + OH^-(aq)

base acid conj acid conj base

Dissociation of strong acids vs weak acids

Strong acids fully dissociate (completely forms H⁺)

Goes to completion (→)

ex: HCl → H⁺ + Cl^–

Weak acids partially dissociate

Reaction is a chemical equilibrium (⇆)

ex: HF ⇆ H⁺ + F^–

What determines the strength of an acid HA in terms of its dissociation?

Stronger acids (HA) more easily dissociate into H⁺ and A⁻.

How does electronegativity of atom A in HA affect acid strength

Higher electronegativity = share e- less equally = weaker H–A bond = easier to remove H⁺ = stronger acid.

What is an oxoacid?

An acid containing H, O, and a central nonmetal atom (ex. H₂SO₄).

How do double bonds in an oxoacid affect its acid strength?

More double bonds = more resonance structures = more stable conjugate base = stronger acid.

ex. H₂SO₄ stronger than H₂SO₃. Sulfuric acid has more double bonds

What functional group defines a the organic acid, carboxylic acid?

The carboxyl group (–COOH)

R - COOH

Are carboxylic acids weak or strong and how does this affect their ability to dissociate?

They are weak acids, and only partially dissociate, existing in equilibrium

How does the R group of a carboxylic acid affect the acid strength

Electron-withdrawing groups (EWGs) pull electron density away from the COOH group (toward R group) → delocalize and spread out the negative charge = the negative charge on the conjugate base is more stable → stronger acid.

EWGs often have high electronegativity or resonance

Properties of strong bases vs weak

Strong: Tend to be ionic compounds (contains metal and hydroxide)

Fully dissociates (completely forms OH–)
Ex: NaOH → Na⁺ + OH^–

Reaction goes to completion

Weak: Tend to be molecular compounds (contains all nonmetal)

Partially dissociates (partially forms OH–)
Ex: NH₃ + H₂O ⇄ NH_4⁺ + OH^-
Reaction is a chemical equilibrium

What ions are always soluble

Ammonium (NH₄⁺)
Hydrogen (H⁺)
Alkali metals (group 1A)
Nitrate (NO₃^- )
Perchlorate (ClO₄^- ) & Chlorate (ClO₃^-)
Acetate (CH₃COO^-)

What ions are usually soluble and their exceptions

Halides/group 17 (F^-,Cl^-,Br^-, I^- )

Exceptions (insoluble if with):
Pb²⁺, Hg₂²⁺, Ag⁺

Sulfate (SO₄^2-)

Exceptions (insoluble if with):

Pb²⁺, Hg₂²⁺, Ag⁺, Ba²⁺, Ca²⁺, Sr²⁺

What ions are sparingly soluble (insoluble) and their exceptions?

Sulfide (S^2-)
Hydroxide (OH^- )
Oxide (O^2- )
Carbonate (CO₃^2- )
Phosphate (PO₄^3-)
Chromate (CrO₄^2- )

Exceptions:
soluble if with any of
the cations listed in the
always soluble box

What functional group defines the organic base amine?

A nitrogen atom with a lone pair (ex. –NH₂, –NH–)

Ionic product constant of water (K_w)

K_w = K_c = [H₃O⁺] [OH^–] = 1.0 x 10^-14 at 25°C

The pH scale and acidity/basicity

[H+] or [H3O+] pH Value Solution Type
> 1.0 x 10-7 pH < 7 Acidic
= 1.0 x 10-7 pH = 7 Neutral
< 1.0 x 10-7 pH > 7 Basic

pH/pOH and [H⁺]/ [OH^-] calculations

pH = -log[H⁺]

[H⁺] = 10^-pH

pOH = -log[OH^-]

[OH^-] = 10^-pOH

equation for pK_w

pK_w = 14 = pH + pOH

Acid dissociation constant (K_a) of weak acids

HA (aq) ⇄ H⁺ (aq) + A^– (aq)

K_a = [H⁺]_eq[A^-]_eq / [HA]_eq

Acid strength and acid dissociation constant

K_a↑ = pK_a ↓ = ↑ weak acid strength

The greater the value of K_a, the stronger the acid

Percent ionization equation and acid strength

Percent ionization = (concentration ionized / original concentration) x 100 = (Δ[HA] / [HA]_initial) x 100

The greater the percent ionization, the stronger the acid.

What is a polyprotic acid?

Acids that have more than one ionizable proton (ex. H₂SO₄, H₂CO₃, H₃PO₄)

How do weak polyprotic acids ionize?

In successive steps, each releasing one proton at a time.

H₂CO₃(aq) ⇄ H⁺(aq) + HCO₃^- (aq) K_a1
HCO₃^- (aq) ⇄ H⁺(aq) + CO₃^2- (aq) K_a2

K_a values for ionization steps of a polyprotic acid

Each step has its own Ka
Easier to remove the first proton than the second, etc.
K_a1 > K_a2 > K_a3....

(Because after losing a proton, the conjugate base is more negatively charged and holds onto the next proton more tightly)

Successive equilibrium constants have less and less impact on pH

Relationship between K_a, K_b, and K_w

K_a x K_b = K_w = 1.0 × 10^-14

Relationship between pK_a and pK_b

pK_a + pK_b = 14

Strength of conjugate acid/base if it’s formed from strong acids/base

extremely weak - cannot act as an acid/base

Strength of conjugate acid/base if it’s formed from weak acids/base

Conjugates of weak acids/bases are stronger than the original, but still weak and able to act as an acid/base

What is the composition of a salt?

An ionic compound composed of a metal cation and a nonmetal anion

Ex. NaCl, KCl, MgSO₄

How to determine acidic/basic properties of a salt

Separate it into its constituent cation and anion. Then, identify the parent acid and base of each ion

How to Determine if an Ion is Acidic, Basic, or Neutral

Break the salt into its positive cation and negative anion.

Identify the base that would combine with the acid to form the cation. 

Identify the acid that would combine with the base to form the anion.

Ex. NaCl

Na⁺ Cl^-

Base: Acid:

NaOH HCl

Combined Effect of Cations and Anions in Salt Acidity/basicity

1) neutral + neutral = neutral.

2) neutral + acidic = acidic

3) neutral + basic = basic

6) Strong acid + strong base = neutral

7) Strong acid + weak base = acidic

8) weak acid + strong base = basic

9) weak acid + weak base = ?

if Ka > Kb, acidic

if Ka < Kb, basic

If Ka ≈ Kb, neutral

What is a buffer and what are its components

It is a solution that resists changes in pH

It's composed of a weak acid (base) and it's conjugate base (acid)

What does buffer capacity depend on and when is it most effective

The capacity depends on the concentrations of the weak acid and its conjugate base. more of each = greater buffer capacity.

It is most effective when the weak acid and conjugate base are in comparable amounts

What is the Henderson–Hasselbalch equation and what is it used for

Used for calculating the pH and pOH of buffer

pH = pKa + log [B]/[A] B = conj base, A = weak acid

pOH = pKb + log [A]/[B] A = conj acid, B = weak base

What types of substances can be mixed to make a buffer

1) A Weak acid and its conjugate base
- Ex. CH3COOH and CH3COONa

2) A weak base and its conjugate acid
- Ex. NH3 and NH4Cl

3) A strong acid and a weak base (with extra amount)
- Ex. HCl and NH3 H+(aq) + NH3(aq) → NH4+(aq)

4) A strong base and a weak acid(with extra amount)
- Ex. KOH, HF OH-(aq) + HF (aq) → F- (aq) + H2O(aq)

5) A weak acid and a weak base(either one in extra amount)
-Ex. CH3COOH, NH3
CH3COOH(aq) + NH3 (aq) → CH3COO- (aq) + NH3+(aq)

How to choose a buffer to maintain a specific pH

Choose a weak acid with pKa close to the desired pH

To reach the desired pH, adjust [A^-]/[HA]

To make the buffer work best, [A^-]/[HA] should be between 0.1 and 10

what is the equivalence point of a titration

the point where moles of acid = moles of base (reaction is stoichiometrically complete).

• xFound on a pH curve as the steepest part of the curve

What is the endpoint of a titration and how is it different from the equivalence point

Endpoint: The point in a titration where the indicator changes color, signaling that the titration should stop. It is a visual cue. It is usually close to equivalence point.

The equivalence point is the theoretical point in a titration where the moles of the titrant exactly equal the moles of the analyte

pH of equivalence points between a

• strong acid + strong base

• weak acid + strong base

• strong acid + weak base

• strong acid + strong base

pH = 7

• weak acid + strong base

pH > 7

• strong acid + weak base

pH < 7

Why do we usually avoid using a weak acid or weak base as the titrant?

Because weak acids/bases do not fully ionize, they produce a smaller pH change near the equivalence point. This makes the titration less accurate and the equivalence point harder to detect with an indicator.

pH changes during titration of a weak base (NH₃) with a strong acid (HCl)

• Beginning

• Before equivalence point

• At equivalence

• After equivalence

• Beginning: pH is set by weak base NH₃.

• Before equivalence point: NH₃ + NH₄⁺ solution forms a buffer, controlling pH.

• At equivalence point: All NH₃ is converted to NH₄⁺; NH₄⁺ hydrolysis lowers pH → pH < 7.

• After equivalence point: Excess HCl controls pH; effect of NH₄⁺ is negligible.

*Buffer region is the region between first addition of HCl and equivalence pt

What is a polyprotic acid/base and characteristics of its titration

Polyprotic acids are acids that can donate more than one proton

Theres an equivalence point for each acidic proton.

What is the solubility product constant

Salt (s) ⇌ Cation⁺ (aq) + Anion^- (aq)

K_sp = [Cation⁺][Anion^-]

K_sp and solubility

Solubility tells you how much of a salt can dissolve in a
certain amount of water

Larger K_sp = more soluble
Amount of salt does not alter K_sp
The value of Ksp depends only on the temperature.

What is molar solubility? How is it different from solubility in g/L?

• Molar solubility (M): Maximum moles of a solute that dissolve in 1 L of water.

• Solubility (g/L): Maximum grams of solute that dissolve in 1 L of water.

Example: PbCl₂ (s) ⇌ Pb²⁺ + 2Cl⁻

Step	Pb²⁺	Cl⁻
Initial (I)	0	0
Change (C)	+x	+2x
Equilibrium (E)	x	2x

• Molar solubility = x

• Solubility (g/L) = x · molar mass

What is the difference between Q and K, and how do they predict reaction direction?

• K = equilibrium constant (describes the system at equilibrium).

• Q = reaction quotient (describes the system at any moment).

• Q = K → system at equilibrium (no net change).

• Q < K → reaction shifts right (forms products).

• Q > K → reaction shifts left (forms reactants).

How do you predict whether a precipitate will form using Qsp and Ksp?

• Qsp = Ksp → solution is saturated → no precipitate

• Qsp < Ksp → solution is unsaturated → no precipitate forms

• Qsp > Ksp → solution is supersaturated → precipitate WILL form