1.
The Nature of Acids and Bases - Acid Strength and the Acid Ionization Constant (Ka)
Rachel Pietrow

2. The Nature of Acids and Bases
General properties of acids
Sour taste, ability to dissolve many metals, turn blue litmus paper red, neutralize bases
Some common acids
Hydrochloric acid, HCl
Sulfuric Acid, H2SO4
Nitric Acid, HNO3
Acetic Acid, HC2H3O2
General properties of bases
Bitter taste, slippery feel (react with oils on skin to form soap-like substances), turn red litmus paper blue, neutralize acids
Alkaloids - organic bases found in plants, often poisonous
Some common bases
Sodium hydroxide, NaOH
Potassium hydroxide, KOH
Sodium bicarbonate, NaHCO3
Ammonia, NH3

3. Structure of Acids and Bases
Binary acids have hydrogens attached to a nonmetal
Ex: HCl, HF
Oxy acids have hydrogens attached to an oxygen atom
Ex: H2SO4, HNO3
Carboxylic acids have the -COOH group
Ex: HC2H3O2
Most bases contain OH- ions
Ex: NaOH, KOH
Some bases contain CO32- ions
Ex: NaHCO3, CaCO3

4. The Arrhenius Definition
According to Arrhenius
Acid: A substance that produces H+ ions in aqueous solution
Base: A substance that produces OH- ions in aqueous solution
HCl (aq) → H+ (aq) + Cl- (aq)
According to this definition, HCl is an acid because it produces H+ ions in solution
Since HCl is a covalent compound it does not contain ions. However, in water it ionizes completely to form H+ and Cl- ions
These H+ ions are highly reactive and in aqueous solutions they bond to water to form H3O+ (hydronium ion)
H+ (aq) and H3O+ (aq) are often used interchangeably to mean the same thing - an H+ has been dissolved in water
NaOH (aq) → Na+ (aq) + OH- (aq)
NaOH is a base according to this definition because it produces OH- ions in solution
NaOH is ionic and when added to water dissociates into Na+ and OH- ions
H+ (aq) + OH- (aq) →H2O (l)
Under this definition, acids and bases combine to form water
This is known as a neutralization reaction

5. The Brønsted-Lowry Definition
This definition focuses on the transfer of H+ ions
Acid: proton (H+ ion) donor
Base: proton (H+ ion) acceptor
HCl (aq) + H2O (l) →H3O+ (aq) + Cl- (aq)
According to this definition, HCl is an acid because it donates a proton to water
The H+ ion associates with a water molecule to form a hydronium ion (H3O+)
NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq)
Applies nicely to bases such as NH3 that do not inherently contain OH- but still produce them in solutions. According to this definition, NH3 is a base because it accepts a proton from water
Acids (proton donors) and bases (proton acceptors) always occur together in an acid-base reaction
Ex: HCl (aq) + H2O (l) →H3O+ (aq) + Cl- (aq)
HCl is the proton donor (acid) and H2O is the proton acceptor (base)
Ex: NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq)
H2O is the proton donor (acid) and NH3 is the proton acceptor (base)
An H+ ion is referred to as a proton because it is a hydrogen atom without its one electron.

6. Brønsted-Lowry Definition Continued
NH3 (aq) + H2O (l) ⇌ NH4+ (aq) + OH- (aq)
HCl (aq) + H2O (l) →H3O+ (aq) + Cl- (aq)
Water is acting as an acid in the first equation and as a base in the second
According to the Brønsted-Lowry definition, substances that can act as acids or bases (such as water) are amphoteric
Let’s reverse the first equation: NH4+ (aq) + OH- (aq) ⇌ NH3 (aq) + H2O (l)
Here, NH4+ is the proton donor (acid) and OH- is the proton acceptor (base). The substance that was the base (NH3) has become the acid (NH4+ ) and vice versa
NH4+ and NH3 are often referred to as a conjugate acid-base pair - two substances related to each other by the transfer of a proton
A conjugate acid is any base to which a proton has been added
A conjugate base is any acid from which a proton has been removed
Summary:
A base accepts a proton and becomes a conjugate acid
An acid donates a proton and becomes a conjugate base
Base Acid Conjugate Conjugate
acid base

8. Practice
Identify the Brønsted-Lowry acid, the Brønsted-Lowry base, the conjugate acid, and the conjugate base for this reaction.
H2SO4 (aq) + H2O (l) → HSO4- (aq) + H3O+ (aq)
H2SO4 (aq) + H2O (l) → HSO4- (aq) + H3O+ (aq)
Since H2SO4 donates a proton to H2O it is the acid (proton donor). After it donates this proton, it becomes HSO4-, the conjugate base. Since H2O accepts a proton, it is the base (proton acceptor). After it accepts the proton, it becomes H3O+, the conjugate acid.
Acid Base Conjugate Conjugate
base acid

9. Practice
Identify the Brønsted-Lowry acid, the Brønsted-Lowry base, the conjugate acid, and the conjugate base for this reaction.
HCO3- (aq) + H2O (l) ⇌ H2CO3 (aq) + OH- (aq)
HCO3- (aq) + H2O (l) ⇌ H2CO3 (aq) + OH- (aq)
Since H2O donates a proton to HCO3- it is the acid (proton donor). After it donates this proton, it becomes OH-, the conjugate base. Since HCO3- accepts a proton, it is the base (proton acceptor). After it accepts the proton, it becomes H2CO3, the conjugate acid.
Base Acid Conjugate Conjugate
acid base

10. Acid Strength and the Acid Ionization Constant (Ka)
A strong acid completely ionizes in solution
A weak acid only partially ionizes
The strength of an acid depends on the equilibrium
HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq)
HA is a generic formula for an acid
If the equilibrium lies far to the right, the acid completely ionizes and is strong
If the equilibrium lies to the left, only a small percentage of the molecules ionizes and the acid is weak

11. Strong Acids
HCl is an example of a strong acid, one that completely ionizes in solution
HCl (aq) + H2O (l) → H3O+ (aq) + Cl- (aq)
So since HCl is a strong acid, an HCl solution contains basically no intact HCl; it has essentially completely ionized to form H3O+ (aq) and Cl- (aq)
A 1.0 M HCl solution has [H3O+] = 1.0 M
Examples of strong acids (know these!)
Hydrochloric acid, HCl
Hydrobromic acid, HBr
Hydriodic acid, HI
Nitric acid, HNO3
Perchloric acid, HClO4
Sulfuric acid, H2SO4 (diprotic)
Except for sulfuric acid, all of these examples are monoprotic acids; they only contain one ionizable proton
Sulfuric acid is diprotic, meaning it contains two ionizable protons

12. Weak Acids
HF is a weak acid, one that only partially ionizes in solution
HF (aq) + H2O (l) ⇌ H3O+ (aq) + F- (aq)
Since HF is a weak acid, an HF solution contains a large number of intact HF molecules and only some H3O+ (aq) and F- (aq) molecules
A 1.0 M HF solution has a [H3O+] that is much less than one, since only some of the HF molecules ionized to form H3O+
Examples of weak acids
Hydrofluoric acid, HF
Acetic acid, HC2H3O2
Formic acid, HCHO2
Sulfurous acid, H2SO3 (diprotic)
Carbonic acid, H2CO3 (diprotic)
Phosphoric acid, H3PO4 (triprotic)
Triprotic - three ionizable protons

13. Why?
The strength or weakness of an acid depends on the attraction between the anion (conjugate base) and the H+ ion
HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq)
The degree to which this generic reaction proceeds in the forward direction (how completely the molecule ionizes) depends on the strength of the attraction between H+ and A-
If the attraction is weak, the reaction favors the forward direction and the acid is strong
If the attraction is strong, the reaction favors the reverse reaction and the acid is weak
In general, the stronger the acid, the weaker the conjugate base and vice versa.
In HCl (a strong acid), the conjugate base (Cl-) has a fairly weak attraction to H+
In HF (a weak acid), the conjugate base (F-) has a greater attraction to H+
If the forward reaction (that of the acid) is likely to occur, then the reverse reaction (that of the conjugate base) has a lower tendency to occur.
In HCl, the reverse reaction has a lower tendency to occur
In HF, the reverse reaction occurs to a significant degree

14. The Acid Ionization Constant (Ka)
The acid ionization constant (Ka), the equilibrium constant for the ionization reaction of a weak acid, is used to quantify the relative strength of an acid
HA (aq) + H2O (l) ⇌ H3O+ (aq) + A- (aq)
HA (aq) ⇌ H+ (aq) + A- (aq)
The equilibrium constant for these two equivalent reactions is: Ka= [H3O+][A-]/[HA] = [H+][A-]/[HA]
Since [H3O+] is equal to [H+], these expressions are equal
The ionization constants for all weak acids are relatively small, but there is some variance in magnitude
The smaller Ka, the less the acid ionizes and the weaker the acid
The larger Ka, the more the acid ionizes and the stronger the acid