 As a reminder, water molecules are polar and are in
continuous motion
 Sometimes, water molecules can be moving quickly
enough to collide with each other and transfer a
hydrogen ion from one water molecule to the next
 A water molecule that loses a hydrogen becomes a
negatively charged hydroxide ion (OH-)
 A water molecule that gains a hydrogen becomes a
positively charged hydronium ion (H3O+)
 Self-Ionization: the process in which water molecules
produce ions
H2O(l)
H+(aq) + OH-(aq)
 In water/aqueous solution, hydrogen ion (H+) are
always joined to a hydronium ion (H3O+)
 H+ and H3O+ will both represent hydrogen ions in
aqueous solution
 In pure water @ 25 °C, the equilibrium concentration
of H+ and OH- is 1.0 x 10-7 M
 Concentrations of H+ and OH- are equal in pure water
NEUTRAL SOLUTION
 Remember Le Chatelier’s principle????
 What happens to the reaction when the concentration
of a substance is increased?
 When H+ increases, OH- decreases
 When H+ decreases, OH- increases
H+(aq) + OH-(aq)
H2O(l)
The equilibrium shifts, so more water molecules are
produced
 For aqueous solutions, the product of the H+ and OH-
concentrations equals 1 x 10-14
[H+] x [OH-] = 1.0 x 10-14
 This equation is true for all dilute aqueous solutions @
25 °C
 Even if the concentrations of H+ and OH- change when
substances are added to water, the product of H+ and
OH- will always be 1.0 x 10-14
ION-PRODUCT CONSTANT FOR WATER, KW
Kw = [H+] x [OH-] = 1.0 x 10-14
 Acidic Solution: A solution in which the H+
concentration is greater than the OH- concentration
 The [H+] is greater than 1.0 x 10-7
 HCl(g)
H+ (aq) + Cl-(aq)
 Basic/Alkaline Solution: A solution in which the OH-
concentration is greater than the H+ concentration
 The [H+] of a basic solution is less than 1.0 x 10-7
 NaOH(s)
Na+ (aq) + OH-(aq)
Log10 x = y → 10y = x
Log10 x
Meaning: 10 to what power gives me X
Log10 x → log x
 The pH scale ranges from 0 to 14
 Neutral solutions have a ph of 7
 pH of 14 is strongly basic
 pH of 0 is strongly acidic
 Calculating pH
 The pH of a solution is the negative logarithm of the
hydrogen-ion concentration
 pH = -log[H+]
 If [H+] of a solution is greater than 1 x 10-7 than the pH is
less than 7.0 and its acidic
 If [OH-] of a solution is less than 1 x 10-7 than the pH is
more than 7.0 and its basic
 Calculating pOH
 The pOH of a solution is the negative logarithm of the
hydroxide-ion concentration
 pOH = -log[OH-]
 A neutral solution has a pOH of 7
 A basic solution has a pOH less than 7
 An acidic solution has a pOH greater than 7
 Relationship between pH and pOH
 pH + pOH = 14
 pH = 14 – pOH
 pOH = 14 –pH
 Using pH or pOH to determine concentration
 If pH = -log [H+]
 Than [H+] = 10-pH
 Remember that the base of log is 10 and we
can think: 10 to the what power ‘y’will give me
(x)
log (x) = y
10y = x
Always express concentrations in scientific
notation
NH₃ + H₂O → NH₄⁺ + OH⁻
Base
Acid
C. A.
C. B.
NH₄⁺ + OH⁻ → NH₃ + H₂O
Acid
Base C. B.
C. A.
 Strength of Conjugate Acid-Base Pairs
 Inversely proportional
 The stronger the acid, the weaker its conjugate base
The Hydronium Ion – H₃O⁺
H⁺ + H₂O → H₃O⁺
(H⁺ = H₃O⁺)
HCl + H₂O → H⁺(aq) + Cl⁻(aq)
Acid
Base ↳ H₃O⁺
NH₃ + H₂O → NH₄⁺(aq) + OH⁻(aq)
Base Acid
Amphoteric – a substance that acts as either an acid or a
base.
strong acid: completely dissociates in water
HCl + H2O → H3O⁺ + Cl⁻
weak acid: partially dissociates in water
HC2H3O2 + H2O → H3O⁺ + C2H3O2⁻
strong base: completely dissociates in water
NaOH + H2O → Na⁺(aq) + OH⁻(aq)
weak acid: partially dissociates in water
NH3 + H2O → NH4⁺(aq) + OH⁻(aq)
IONIZATION OF WATER:
H2O(ℓ) + H2O(ℓ) ⇄ H3O⁺(aq) + OH⁻(aq)
Kw = [H3O⁺] [OH⁻]
[H3O⁺] = 1.0 x 10-7M
[OH⁻] = 1.0 x 10-7M
Kw = [H3O⁺] [OH⁻] = (1.0 x 10-7)2 = 1.0 x 10-14
Kw = 1.0 x 10-14
pH equation:
pH = -log [H3O+]
 Important equations to know for determining the
pH, [H3O+], and [OH-].
Kw = [H3O⁺] [OH⁻] = (1.0 x 10-7)2 = 1.0 x 10-14
pH = -log [H3O+]
[H3O+] = 10-pH
pOH = -log [OH-]
[OH-] = 10-pOH
pOH + pH = 14
The Acid Dissociation Constant (Ka)
HA(aq) + H2O(ℓ) ↔ H3O+(aq) + A-(aq)
Ka = [H3O+] [A-]
[HA]
Ka = acid dissociation constant
- measures the strength of acids
↑Ka = Stronger acid
↓Ka = weaker acid
The Approximation vs. The Quadratic
- If [HA]/ Ka > 1000 use approximation
- If [HA]/ Ka < 1000 use quadratic
% Ionization:
% Ionization = [H3O+]/ [HA] x 100%
The Base Dissociation Constant (Kb)
B(aq) + H2O(ℓ) ↔ HB+(aq) + OH-(aq)
Kb = [HB+] [OH-]
[B]
Kb = base dissociation constant
- measures the strength of bases
↑Kb = stronger base
↓ Kb = weaker base
Titrations – determines the concentration of a
weak acid and/or a weak base.
Equation: mLaMa = mLbMb
Standard solution: acid/base with a known
concentration
Indicator: Changes color at a certain pH
Equivalence point: solution becomes neutral
End point: when solution changes color
-total moles H+ ions = total moles of OH- ions
Example:
A volume of 41.6mL of 1.00M NaOH was required
to neutralize a 50.0mL volume of vinegar. What is
the concentration of acetic acid in vinegar?
(0.832M)
Example:
A volume of 25mL of 0.120M H2SO4 neutralizes
40mL of NaOH solution. What is the
concentration of NaOH?
(0.075M)