Acids and Bases
Definitions:
1. ArrheniusAcid- substance that dissociates in water to
produce hydrogen ions - H+
Examples: HCl, HNO3, H2SO4, etc
Base- substance that dissociates in water to
produce hydroxide ions- OHExamples: NaOH, KOH, Mg(OH)2
According to this definition, H must be present in an
acid, and OH in a base
• HNO3 + H2O
H3O+ + NO3H3O+ is the Hydronium ion
You may see HNO3
H+ + NO3• HCl + KOH
KCl + H2O
Neutralization reaction
Acid + Base
a Salt + water
2. Brønsted-LowryAcid- substance that can donate H+ ions
Base- substance that can accept H+ ions
Substances do not need to be in water, and the base
does not need to have OH.
This expands the Arrhenius definition so that other
substances can be considered as acids and bases.
H+ is simply a proton, so the definition of an acid can
now be a proton donor and a base is a proton
acceptor.
• NH3 + H2O
NH4+ + OHWhat is the acid?
H2O – it donated the H+
What is the base?
NH3 – it accepted the H+
Conjugate Acid- Base Pairs: Many times a reaction
occurs in a forward and reverse direction, so what is
an acid in the forward direction becomes the
conjugate base in the reverse direction, and the base
becomes the conjugate acid in the reverse reaction
H3O+
+
Cl conjugate
conjugate
acid
base
Monoprotic acid- an acid that will donate 1 H+
HNO3
Diprotic acid- an acid that donates 2 H+
H2SO4
Triprotic acid- an acid that donates 3 H+
H3PO4
Amphoteric- a substance that can act as both an
acid or a base- Water
HCl
acid
+
H2O
base
3. Lewis Acids and BasesAcid- substance that can accept a pair of
electrons to form a covalent bond
Base- substance that can donate a pair of
electrons to form a covalent bond
This now allows other substances to be considered
an acid or a base
Acid – Base Definitions
Type
Acid
Arrhenius
H+ producer
Bronsted-Lowry H+ donor
Lewis
Electron- pair
acceptor
Base
OH- producer
H+ acceptor
Electron-pair
donor
Properties of Acids and Bases
• Acids:
-sour or tart taste
-can burn skin if stronger, may burn if you
get it in a cut
-react strongly with most metals, usually to
produce hydrogen gas
-form weak or strong electrolytes
-change blue litmus paper to red
• Bases:
-bitter taste
-feels smooth, slippery
-not reactive with most metals
-weak or strong electrolytes
-turns litmus paper from red to blue
Strengths of Acids and Bases
• Strong acid- easily dissociates H+ to water
- becomes strong electrolyte
• Weak acid- does not easily dissociate H+ ions
- weak electrolyte
Common strong acids: HCl, HBr, HI, HNO3, H2SO4,
HClO4
Weak acids:HC2H3O2, HCN, HNO2, HF, HClO, HCO3-
• Strong Bases: substances with a strong affinity for
H+, those with OHWeak Bases- ions that react only partially in water
to form OH- ions
Common strong bases: CaO, NaOH, KOH, Ca(OH)2
Weak bases- NH3, H2NNH2, CO32-, PO43Strong bases are strong electrolytes, weak bases
are weak electrolytes
Self-Ionization of Water
• When you have pure water, due to the motion of
the water molecules and the high polarity of water,
there is a very small amount of H3O+ and OH- that
exist
H2O + H2O
H3O+ + OH-
• In pure water, the [H+] = [OH-], where [X] is equal
to the concentration of X in solution, usually
measured in M
• When the concentrations are equal, the substance
is known as a neutral solution
• [H+] =[OH-] = 1.0 x 10-7 M
• So, [H+] x [OH-] = 1.0 x 10-14M2 = Kw = ion product
constant of water
• Since these concentration values are multiplied to
give a constant, if [H+] increases, [OH-] must
decrease
• Because Kw is a constant, if you know the
concentration of H+ or OH-, you can calculate the
other.
• Example: If [H+] = 1.0 x 10-2, then
(1.0 x 10-2) [OH-] = 1.0 x 10-14
[OH-] = 1.0 x 10-14
1.0 x 10-2
= 1.0 x 10-12
When [H+] > [OH-], acidic solution
When [H+] < [OH-], basic solution or alkaline solution
• Because these concentrations are very small,
another scale, the pH scale is used to describe [H+],
and pOH is used to describe [OH-]
• pH = -log[H+]
• In a neutral solution, [H+] = 1.0 x 10-7, so
pH = - log (1.0 x 10-7)
= -(0 + -7)
=7
• pH < 7 = acidic solution
• pH = 7 = neutral solution
• pH > 7 = basic solution
•
•
•
•
•
pOH = - log [OH-]
pH + pOH = 14
To calculate pH, use – log [H+]
To calculate pOH, can use 14 – pH or – log [OH-]
To calculate the concentration from a pH value, use
the antilog or 10x button