Ultimate MCAT General Chemistry 4 - Acids and Bases
What are the 3 definitions for acids and bases?
-Arrhenius (H₃O⁺, OH⁻)
-Bronsted-Lowry (H⁺donor, H⁺ acceptor)
-Lewis (electron acceptor, electron donor)
Explain the properties of water
It is amphoteric therefore it can act as either acid or base
Describe acid dissociation and its constant Ka
When an acid is added to water, it dissociates into hydronium ion (H₃O⁺) and its conjugate base
Ka = its equilibrium constant for dissociation of the reaction HA + H₂O=A⁻+ H₃O⁺=
What does Ka and pKa rely on/not rely on?
It relies on the strength of the acid and NOT its concentration
Explain the relationship between strength of an acid to Ka and pKa.
As strength increases ↑, the Ka increases↑ and pKa↓
Explain the relationship between acid strength and conjugate base strength
As acid strength ↑ (Ka↑, pKa↓), conjugate base strength ↓ (Kb↓,pKb↑)
How is the strength of an acid assessed?
By how much it ionizes/dissociates in water
State the relationship between pKa and pKb
pKa + pKb =14
only at 25°C
Strong acids have what tendency to dissolve?
Dissolve completely → large Ka and small pKa
Describe the dissolving pattern of weak acids
They only dissolve partially
Name some weak acids
Name the product of a weak acid
A weak conjugate base
How does size play a role in the acidity of haloacids? Why?
- The larger the halide = more acidic, since H dissociates easily with larger bonds
- The conjugate base is stabler since it can diffuse the negative charge over the larger anion
Which haloacids are weak? Strong?
HF = weak
The rest are strong
In a row/period, what dictates acid strength?
For oxyacids how does acidity work through what effect?
Acidity is based on how the H dissociates from an O.
increase ↑ number of O = increase↑ acidity because electrons are withdrawn to the resonance effect
Name another factor that governs acidity in oxyacids
Electronegativity if more EN = ↑increase withdraw electron density through the inductive effect
Therefore the oxyacid becomes more acidic
Which is weaker HClO₂ or HIO₂?
HIO₂ because HClO₂ = stronger because Cl is more electronegative than I
Non-metal oxides and non-metal hydroxides (oxyacids) act as
Lewis acid (electron acceptor) and Bronsted Lowry acid (H⁺ donor)
Metal oxides and metal hydroxides act as
Lewis base (electron donor) and Bronsted Lowry base (H⁺ acceptor)
What are 3 organic acids and their pKa ranges?
- Carboxylic acids (COOH) pKa (R)=3-5
pKa (AA) = 2-3
-Phenol pKa= 9.5-10.5
-Alkyl ammonium pKa (R)=9-11
moles of equivalent/L of solution
therefore 1 mol of H₂SO₄=2 normal since sulfuric acid is diprotic
Name the equation for calculating pH
pH=-log₁₀[H₃O⁺] = -log₁₀[H⁺]
In calculating pH, which matters concentration or volume?
Volume does not matter
What is the value of log 2 and log 3 and log 7?
log 2 = 0.3
log 3 = 0.48 = .5
log 7 = .845 = .85
When numbers are multiplied or divided, how does this work for logs?
Explain the relationship between equilibrium constants for an acid its conjugate base
Ka * Kb = 10⁻¹⁴
Name the shortcut for determining pH in weak reagents
pH= .5pKa - .5log₁₀[HA]
as long as
If the pH>pKa (or pH<pKa) determine the basicity/acidity of the solution
If pH>pKa, then the solution is basic and is deprotonated
If pH<pKa, then the solution is acidic and is protonated
For carbonic acid (H₂CO₃), what is the correct relationships for its pKa, pKb in conjugate pairs?
H₂CO₃ ↔ HCO₃⁻ + H⁺ (pKa₁ forward, pKb₂ reverse)
HCO₃⁻ ↔ CO₃²⁻ + H⁺ (pKa₂ forward, pKb₁ reverse)
Describe the Henderson-Hasselbach equation
pH= pKa + log([A⁻]/[HA])
What is the conjugate base of a weak acid?
A weak base