Ultimate MCAT General Chemistry 4 - Acids and Bases

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1

What are the 3 definitions for acids and bases?

-Arrhenius (H₃O⁺, OH⁻)
-Bronsted-Lowry (H⁺donor, H⁺ acceptor)
-Lewis (electron acceptor, electron donor)

2

Explain the properties of water

It is amphoteric therefore it can act as either acid or base

3

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⁺=
[H₃O⁺][A⁻]/[HA]

4

What does Ka and pKa rely on/not rely on?

It relies on the strength of the acid and NOT its concentration

5

Explain the relationship between strength of an acid to Ka and pKa.

As strength increases ↑, the Ka increases↑ and pKa↓

6

Explain the relationship between acid strength and conjugate base strength

As acid strength ↑ (Ka↑, pKa↓), conjugate base strength ↓ (Kb↓,pKb↑)

7

How is the strength of an acid assessed?

By how much it ionizes/dissociates in water

8

State the relationship between pKa and pKb

pKa + pKb =14
HA A⁻

only at 25°C

9

Strong acids have what tendency to dissolve?

Dissolve completely → large Ka and small pKa

Ka>>1 pKa<0

10

Describe the dissolving pattern of weak acids

They only dissolve partially
1>Ka>10⁻¹⁴ 0<pKa<14

11

Name some weak acids

Carboxylic acids
Alkyl ammoniums
Phenols

12

Name the product of a weak acid

A weak conjugate base

13

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

14

Which haloacids are weak? Strong?

HF = weak
The rest are strong

15

In a row/period, what dictates acid strength?

Electronegativity

16

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increase ↑ number of O = increase↑ acidity because electrons are withdrawn to the resonance effect

17

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

18

Which is weaker HClO₂ or HIO₂?

HIO₂ because HClO₂ = stronger because Cl is more electronegative than I

19

Non-metal oxides and non-metal hydroxides (oxyacids) act as

Lewis acid (electron acceptor) and Bronsted Lowry acid (H⁺ donor)

20

Metal oxides and metal hydroxides act as

Lewis base (electron donor) and Bronsted Lowry base (H⁺ acceptor)

21

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
pKa (AA)=9-10

22

Explain normality

moles of equivalent/L of solution
therefore 1 mol of H₂SO₄=2 normal since sulfuric acid is diprotic

23

Name the equation for calculating pH

pH=-log₁₀[H₃O⁺] = -log₁₀[H⁺]

24

In calculating pH, which matters concentration or volume?

Concentration

Volume does not matter

25

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

26

When numbers are multiplied or divided, how does this work for logs?

Add

Subtract

27

Explain the relationship between equilibrium constants for an acid its conjugate base

Ka * Kb = 10⁻¹⁴

28

Name the shortcut for determining pH in weak reagents

pH= .5pKa - .5log₁₀[HA]

as long as
1- [HA]>Ka
2- 2<pKa<12

29

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

30

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)

31

Describe the Henderson-Hasselbach equation

pH= pKa + log([A⁻]/[HA])

32

What is the conjugate base of a weak acid?

A weak base