Strength of Acids and Bases
The strength of an acid or base is determined by its extent of ionization in water. Strong acids and bases ionize completely, while weak ones only partially ionize.
Key Points
Acid Dissociation Constant (Ka): Measures acid strength – higher Ka = stronger acid
pKa = -log Ka: Lower pKa = stronger acid
Strong acids: HCl, H₂SO₄, HNO₃, HBr, HI, HClO₄
Weak acids: CH₃COOH, H₂CO₃, HF, HCN
Strong bases: NaOH, KOH, Ca(OH)₂, Ba(OH)₂
Weak bases: NH₃, CH₃NH₂, C₅H₅N
Conjugate base of strong acid = weak base
Conjugate acid of strong base = weak acid
Tips & Tricks
Remember the 7 strong acids: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₃, HClO₄
For bases: Group 1 hydroxides are strong, Group 2 hydroxides are somewhat soluble
pKa and Ka relationship: pKa decreases as acid strength increases
Water (H₂O) is both an acid and a base – it’s amphoteric
Use the formula: pKa = -log(Ka) and Ka = 10^(-pKa)
Memorization Strategy
Create flashcards with common acids/bases and their Ka/pKa values
Use mnemonic: “So I Brought No Clean Clothes” for strong acids: H₂SO₄, HI, HBr, HNO₃, HCl, HClO₃, HClO₄
Practice calculating pKa from Ka and vice versa
Group acids/bases by strength in a table for visual learning
HA + H₂O ⇌ H₃O⁺ + A⁻ Ka = [H₃O⁺][A⁻] / [HA]
| Property | Strong Acids | Weak Acids |
|---|---|---|
| Ionization | Complete (≈100%) | Partial (<5%) |
| Ka Value | Very high (Ka >> 1) | Small (Ka << 1) |
| Examples | HCl, HNO₃, H₂SO₄ | CH₃COOH, H₂CO₃, HF |
Buffer Solutions
Buffer solutions resist changes in pH when small amounts of acid or base are added. They consist of a weak acid and its conjugate base, or a weak base and its conjugate acid.
Key Points
Buffers maintain nearly constant pH
Composition: Weak acid + its salt OR weak base + its salt
Blood buffer: H₂CO₃ / HCO₃⁻ maintains pH 7.35-7.45
Buffer capacity depends on concentrations of buffer components
pH = pKa + log([A⁻]/[HA]) – Henderson-Hasselbalch equation
Effective buffer range: pKa ± 1
Tips & Tricks
To prepare a buffer: Mix weak acid with salt of its conjugate base
Maximum buffer capacity when [HA] = [A⁻] (pH = pKa)
Blood pH outside 6.8-7.8 is fatal
Bicarbonate buffer is the most important in blood
Carbonic anhydrase enzyme speeds up CO₂ + H₂O ⇌ H₂CO₃ reaction
Memorization Strategy
Remember the Henderson-Hasselbalch equation: pH = pKa + log([A⁻]/[HA])
Create a mental image of a buffer as a “pH sponge” that soaks up added H⁺ or OH⁻ ions
Make a table of common buffer systems and their pH ranges
Practice calculating buffer pH with different ratios of acid to conjugate base
H₂CO₃ ⇌ H⁺ + HCO₃⁻ (Blood buffer system)
Note: The bicarbonate buffer system is crucial for maintaining blood pH. Disruption can lead to acidosis or alkalosis, which are serious medical conditions.
Solubility Product and Precipitation
The solubility product constant (Ksp) describes the equilibrium between a solid and its ions in a saturated solution.
Key Points
Ksp applies only to slightly soluble salts
For AmBn(s) ⇌ mAn+ + nBm-: Ksp = [An+]m[Bm-]n
If ionic product > Ksp, precipitation occurs
If ionic product < Ksp, more solid dissolves
If ionic product = Ksp, solution is saturated
Lower Ksp = less soluble salt
Tips & Tricks
Pure solids don’t appear in Ksp expression
For 1:1 salts (like AgCl), Ksp = s² where s is solubility
For 1:2 salts (like CaF₂), Ksp = 4s³
Common ion effect reduces solubility
To compare solubilities, convert Ksp to molar solubility
Memorization Strategy
Practice writing Ksp expressions for different salt types
Create a chart of common Ksp values
Remember that stoichiometry affects the Ksp expression
Work through example problems calculating solubility from Ksp
AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq) Ksp = [Ag⁺][Cl⁻] = 1.8 × 10⁻¹⁰
Common Ion Effect
The common ion effect occurs when adding an ion that is already present in an equilibrium, suppressing ionization or reducing solubility.
Key Points
Based on Le Chatelier’s Principle
Adding common ion shifts equilibrium to reduce its concentration
Reduces dissociation of weak electrolytes
Decreases solubility of slightly soluble salts
Used in qualitative analysis and salt purification
Tips & Tricks
Adding NaF to HF solution reduces HF dissociation
Adding NaCl to saturated AgCl solution causes precipitation
Common ion effect is used in buffer solutions
In brine purification, HCl gas is passed to precipitate pure NaCl
NH₄Cl added to NH₄OH controls OH⁻ concentration
Memorization Strategy
Think of the common ion as “crowding out” the equilibrium
Remember that common ions always decrease solubility or dissociation
Practice calculating the effect of common ions on solubility
Create examples of common ion effect in everyday applications
HF ⇌ H⁺ + F⁻ (Adding NaF shifts equilibrium left)
Partition Coefficient
The partition coefficient (Kpc) describes how a solute distributes between two immiscible solvents.
Key Points
Kpc = [Solute]solvent1 / [Solute]solvent2
Based on Nernst’s distribution law
Unitless ratio of concentrations
Higher Kpc means solute prefers that solvent
“Like dissolves like” – polar solutes prefer polar solvents
Affected by temperature, solute size, and solute-solvent interactions
Tips & Tricks
Polar solutes (e.g., NaCl) prefer polar solvents (e.g., water)
Non-polar solutes (e.g., benzene) prefer non-polar solvents (e.g., octanol)
Temperature generally increases solubility
Smaller molecules have higher partition coefficients
Used in extraction processes and drug design
Memorization Strategy
Remember “like dissolves like” as the golden rule
Create a mental image of oil and water separating
Practice calculating partition coefficients from concentration data
Make a table of common solvents and their polarities
Kpc = [Solute]organic / [Solute]aqueous