MCAT Study Guide Chemistry Ch. 11 – Acid and Base 2017-08-15T06:45:06+00:00

I.          11.1:  DEFINITIONS

A.     ARRHENIUS ACIDS AND BASES – oldest definition

1.     Acids ionize in water to produce H+ ions

2.     Bases ionize in water to produce OH ions

B.     BRONSTED-LOWRY ACIDS AND BASES – more complex definition of base

1.     Acids are proton donors

2.     Bases are proton acceptors

C.    LEWIS ACIDS AND BASES – broadest definition of acids and bases

1.     Lewis acids are electron pair acceptors

2.     Lewis bases are electron pair donors



A.     Conjugate base

When a Bronsted-Lowry acid donates a proton, the remaining structure is the conjugate base

B.     Conjugate acid

When a Bronsted-Lowry base bonds with a proton, this is the conjugate acid of the original base



A.     Strong acid

Dissociates nearly completely in water

1.     Ka =

[H3O+][A]/[HA]  → equilibrium constant for acid; ↑ Ka = = stronger acid

2.     Ka is known as acid-ionization (acid-dissociation) constant

a)     Ka < 1 → weak acid (reactants favored)

b)     K> 1 → strong acid (products favored)

3.     Larger anions are better able to spread their negative charge, making them more stable; this makes them stronger acids

4.     Strong acids – HI, HBr, HCl, HClO4, H2SO4, HNO3

B.     Strong base

Same as stong acid; dissociates nearly completely in water

1.     Kb = [HB+][OH]/[B]

2.     Kb is known as base-ionization (base-dissociation) constant

3.     The larger the Kb, the stronger the base

4.     Strong bases – group I hydroxides, group I oxides, some group II hydroxides, metal amides

5.     Weak bases – ammonia, amines, conjugate bases of many weak acids


1.     The conjugate base of a strong acid has no basic properties

2.     The conjugate base of a weak acid is a weak base (the weaker the acid, the stronger the conjugate base)

3.     The conjugate acid of a strong base has no acidic properties

4.     The conjugate acid of a weak base is a weak acid (the weaker the base, the stronger the conjugate acid)



A.     Water auto-ionizes:

1.     H2O(l) + H2O(l) ⇋ H3O+(aq) + OH(aq)

B.     Kw = [H3O+][OH] = 1.0 x 10-14 (this is at 25° C; Kw ↑ with ↑ T)

C.    Kis constant regardless of the H3O+ concentration! This means that if H+ or OH are added to water, the ⇋ is disturbed, and reverse reaction will occur until Kw is back to 1.0 x 10-14


V.          11.5:  pH


1.    pH = -log[H+]

a)     Since [H+] = 10-7 in water, the pH of water = 7

b)     pH is usually between 0-14, but very strong acids and bases can go beyond this

2.     pOH = -log[OH]  → same idea of pH

3.     pH + pOH = 14


1.     pKa = -logKa → large Ka; pKb = -logKb

a)     Side note – relate this to pH:

(1)   Ka  = [H3O+][A]/[HA] → [H3O+] = Ka[HA]/[A]

(2)   pH = pKa + log([A]/[HA])  → Hasselbach-Henderson!

2.     Ka and Kb are related →

a)     KaKb = Kw = 1 x 10-14  ∴  pKa + pKb = 14

(1)   This applies to all acid base conjugate pairs!!

(2)   NH4+  ⇋  NH3 + H3O+ ⇒ Ka = [H3O+][NH3]/[NH4+]

(3)   NH3  ⇋ NH4+ + OH  ⇒  Kb = [NH4+][OH]/[NH3]

(4)   KaKb = [H3O+][OH] = Kw


1.     Strong acids

a)     Since strong acids dissociate completely, the molarity of the products will be the same as the molarity of the reactant

(1)   EX:  0.01 M HCl → [H+] = 0.01 M = 10-2 M → pH = 2

2.     Weak acids

a)     pH of weak acids is related to the amount of dissociation there is (more complicated)

(1)   EX:  0.2 mol HCN is added to water to create 1 L

(a)   Initially, before dissociation, solution is 0.2M HCN

(b)   At equilibrium, the concentration of HCN is now 0.2-x (where x = [H+] and x = [CN]

(c)   Ka = [H+][CN]/[HCN] = x*x/(0.02-x)

(d)   if if Ka = 4.9×10-10, then 4.9×10-10 = x2/(0.2-x)

(i)          Since Ka is so small, that means hardly any HCN dissociates, and we can assume that (0.2-x) is essentially the same as 0.2

(e)   ∴ 4.9×10-10 = x2/0.2  ⇒ x = 1×10-5

(f)     [H+] = 1×10-5  ∴ pH = 5



A.     Neutralization reactions

Reaction when an acid and base combine, often forming salt and H2O

1.     Strong acid + strong base = neutral pH

2.     If reaction involves a weak acid or weak base, the resulting solution will generally not be neutral

3.     No matter how weak an acid or base is, when it is mixed with equimolar amount of strong acid or base, neutralization is complete

4.     All  neutralizations are exothermic, and the energy released from the reaction is the same for all neutralizations

B.     Formula:

1.     a *[A]*VA = b*[B]*VB

a)     a = number of acidic hydrogens per formula unit

b)     b = number of H3O+ ions the base can accept

VII.          11.7:  HYDROLYSIS OF SALT

A.     Salt

Ionic compound that dissociates in water into ions; will either be acidic, basic, or neutral

B.     EX:

1.     NH4Cl → NH4+ + Cl will be slightly acidic because of NH4+

2.     Na(CH3COO) → Na+ + CH3COO will be slightly basic because of CH3COO



VIII.          11.8:  BUFFER SOLUTIONS

A.     Buffer solution

A solution that resists changing pH when a small amount of acid or base is added, comes from the presence of a weak acid and its conjugate base (or weak base and conjugate acid) in roughly equal concentrations

1.     EX:  acetic acid and sodium acetate (remember, acetic acid will only partially dissociate, but sodium acetate, a salt, will fully dissociate, leaving an excess of acetate ions)

a)     Ka of acetic acid = 1.75×10-5

b)     Ka = [H3O+][CH3COO]/[CH3COOH]

c)     [H3O+] = Ka[CH3COOH]/[CH3COO]

d)     Since the equilibrium concentrations of both the acid and acetate ion are 0.1M, then:

e)     [H3O+] = Ka[CH3COOH]/[CH3COO] = Ka[0.1M]/[0.1M]

f)       ∴ [H3O+] = 1.75×10-5  → pH = 4.76

2.     What if we add a little bit of strong acid? 0.005 mol HCl will dissociate into 0.005 mol H+ ions; find the new pH:

a)     [H3O+] = Ka[CH3COOH]/[CH3COO] = Ka[0.105M]/[0.095M]

b)     ∴ [H3O+] = 1.75×10-5(1.105) = 1.93×10-5 → pH = 4.71

B.    Henderson-Hasselbalch equation for acid:

1.     pH = pKa + log([conj. base]/[weak acid])

C.    Henderson-Hasselbalch equation for base:

1.     pH = pKb + log([conj. acid]/[weak base])


IX.          11.9:  INDICATORS

A.     Indicator

Weak acid that undergoes a color change when it’s converted to its conjugate base

1.     HA ⇋ H+ + A

2.     HA in nonionized form is one color

3.     A- (conjugate base) has a different color

4.     If the indicator was added to an acid, the ⇋ would push the first equation more towards reactants, and it would turn blue

5.     If the indicator was added to a basic solution, the ⇋ would readjust, pushing more towards the products, turning the color red

X.          11.10:  ACID-BASE TITRATIONS

A.     Acid-base titration

an experimental technique used to determine the identity of unknown weak acid or base by determining its pKa or pKb; consists of adding strong acid (or base) of known identity and concentration (this is the titrant)

B.     As NaOH is added, the equivalent amount of HF gets neutralized with minimal change in the pH (buffering region)

C.    The ⇋ continues to shift until all the equivalence point is reached, when all the H+s have been neutralized by OHs(acid-base equivalence point)

D.    This should be slightly basic! The conjugate base of a weak acid is basic, so the remaining F in solution are slightly basic

E.     [HF]at half-equiv = [F]at half-equiv  → this is when ½ of the H+s have been neutralized by NaOH

F.     pHat half-equiv = pKa + log([F]at half-equiv /[HF]at half-equiv) = pKa + log1 = pKa

1.     ∴ pKa of HF equals the pH at the half-equivalence point!

MCAT Study Guide Chemistry - Kim Matsumoto

More MCAT Study Guide Chemistry


Ch. 3 Chemistry Basics


Ch. 4 Atomic Structure


Ch. 5 Chemical Bonds


Ch. 6 Enthalpy + Entropy


Ch. 7 Calorimetry + Phase Diagrams


Ch. 8 Ideal Gas Law


Ch. 9 Rate Laws


Ch. 10 Equilibrium


Ch. 11 Acids and Bases

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