Acid and Base Chemistry
Arrhenius Theory of Acids
• An acid is a substance that gives off H+ in water.
This is a hydronium ion or proton.
• H2SO4 is sulfuric acid.
H2SO4  2H+ + SO4 2• has one proton and one electron. H+ has lost the
electron and so we call it a proton.
H+ + H2O H3O+ Hydronium Ion
H2SO4 + 2H2O  2H3O+ + SO4 2-
Arrhenius Bases
• According to the Arrhenius Theory, a base
is a substance that dissolves in water to
make OH-.
NaOH(s)  Na+(aq) + OH-(aq)
Ca(OH)2(s)  Ca+2(aq) + 2OH-(aq)
Overall Arrhenius Theory
• Acid + Base → Salt + Water + Heat
• Salt is an ionic compound that results from the
reaction between an acid and a base.
• A solution is acidic if there are more moles of H+
than OH-.
• A solution is basic if there are more moles of OH - than
H+.
• A solution is neutral if the number of moles of H+
and OH- are the same. They then neutralize, or
cancel out, each other.
Examples of Neutralization Equations
• HCl + NaOH → NaCl + H2O
• 2HBr + Ca(OH)2 → CaBr2 + 2H2O
• HNO3 + KOH → KNO3 + H2O
• 3H2SO4 + 2Al(OH)3 → Al2(SO4)3 + 6H2O
Acid Nomenclature
• Binary Acids = H and one nonmetal element.
• Name always begins hydro-, root is based on the
nonmetal name with the suffix -ic, followed by
the word acid.
HCl = Hydrochloric Acid
HBr = Hydrobromic Acid
H2Se = Hydroselenic Acid
Oxyacids
These are “ternary” acids, and contain 3 elements:
H, O and a nonmetal.
• Example: H2SO4
• Ternary or oxyacids do not have hydro as a prefix.
• Name starts with root. The suffix depends on the
anion from which the acid is derived.
• -ate anion = -ic acid.
• -ite anion = -ous acid.
Examples
• HClO3 = H+ + ClO3 (Chlorate ion) Chloric Acid
• HNO2 = Nitrous Acid
• H2CO3 = Carbonic Acid
Titration
Titration is a process in which the amount of acid
and base are carefully measured to exactly
produce a neutral solution.
Technically speaking, the substance being added
(either acid or base) is slowly added until the
solution is neutral, and the volume that is added is
carefully measured and recorded.
Indicator
Substance that changes color as the relative
concentration of acid and base changes.
The best indicator will change color from
acid to neutral or from base to neutral.
• Example--phenolthalein. It is clear in an
acidic solution or neutral solution, and pink
in a basic solution.
Math on Titration Problems
In a titration problem, the moles of H+ and the
moles of OH- must be equal.
• Molarity • Volume = moles, so
MaVa = MbVb
Ma = molarity of H+
Va = volume of acid
Mb = molarity of OHVb = volume of base
Practice Problem
23.1 mL of a .100 M HCl solution is titrated with
15.7 mL of a NaOH base solution. Find the
concentration of the base.
• MaVa = MbVb
• (.100 M)•(23.1mL)=Mb•(15.7mL)
• Mb= .147 M
Logarithms
Logarithm Defined
• X is the logarithm (or log) of A if A can be
written as 10X.
• A = 10X X = log A
• What is the log of 1,000,000?
• Can 1,000,000 be written as a power of 10?
• Yes, 1,000,000 = 10e6.
• So log 1,000,000 = 6.
pH
• Definition: pH is the - log of the [H+]
• pH = - log [H+]
• pH is a measure of the acidity or basicity of
a solution.
• For a neutral solution, pH = 7.
• For an acidic solution, pH < 7.
• For a basic solution, pH > 7.
pH and Logarithms
104 = 10,000 4 is the logarithm of 10,000
• 10e3 = 1,000 3 is the logarithm of 1,000
• 10e2 = 100 2 is the logarithm of 100
• 10e1 = 10
1 is the logarithm of 10
• 10e0 = 1
0 is the logarithm of 1
• 10e-1 = .1 -1 is the logarithm of .1
• 10e-2 = .01
• 10e-3 = .001
What is the log of 4,258?
• 4,258 = 10X
• What is X?
Use calc.---- log(4258)=3.63
• 10e3.63 = 4,258
Other important equations
pH = - log [H+]
• pOH = - log [OH-]
• Kw = [H+][OH-] = 1 x 10-14
• pH + pOH = 14
• [H+] = 10e-pH
• [OH-] = 10e-pOH
Dissociation Equilibrium for Water
• H2O(l) ↔ H+(aq) + OH-(aq)
• Kw = [H+][OH-] = 1 x 10-14
So as water dissociates, [H+]=[OH-].
• Then, [H+]2 = 1 x 10-14, and
• [H+] = [OH-] = 1 x 10-7.
• This is true for a neutral solution
Practice Problem
A solution is .20 M in HCl. Find the pH,
pOH and the [OH-].
• HCl → H+ + Cl• .20 M HCl = .20 M H+ = .20 M ClHCl → H+ + Cl• pH = -log [.20 M] = .70
• [OH-] = 1x10e-14/.20 = 5.0x10e-14
• pOH = -log [5.0x10e-14] = 13.30 or
pH + pOH = 14
pOH= 14-.70= 13.30
A solution has a pH of 3.20. Find the
concentration of H+ and OH-. Also find the pOH.
• [H+] = 10e-pH = 10e-3.20 or antilog of -3.20
• [H+] = 6.31 x 10e-4
• pH + pOH = 14
pOH = 10.80.
• [OH-] = 10e-pOH or [H+][OH-]=1x10e-14
• [OH-] = 1.58 x 10-11
Write an equation for dissolving and
dissociating the following acids in water:
HCl, HNO3, H2SO4, H3PO4
HCl → H+(aq) + Cl-(aq) or HCl + H2O → H3O+(aq) + Cl-(aq)
HNO3 → H+(aq) + NO3-(aq) or HNO3 + H2O → H3O+(aq) + NO3-(aq)
H2SO4 → 2H+(aq) + SO4 2-(aq) orH3PO4 ↔ 3H+(aq) + PO4 3-(aq)
Mono-, Di- and Triprotic Acids
A monoprotic acid gives off one H+ for each acid
molecule. HCl and HNO3 are examples of
monoprotic acids.
• A diprotic acid gives off two H+ ions for each acid
molecule. H2SO4 is an example of a diprotic acid.
• A triprotic acid gives off three H+ ions for each
acid molecule. H3PO4 is an example of a triprotic
acid.
Concentration Calculations with Di- and
Triprotic Acids
When you have di- and triprotic acids, the
concentration of H+ is greater than the
concentration of the original acid.
• Find the molarity H+ and the pH of a .200 M
H2SO4 solution.
H2SO4 → 2H+(aq) + SO4 2-(aq)
(.200 moles H2SO4) (2 moles H+) = .400 M H+
( 1 liter soln.) (1 mole H2SO4)
pH = - log (.400) = .398
Dibasic and Tribasic Substances
Bases that have more than one OH- per
molecule follow a similar pattern to the
acids.
Ca(OH)2 → Ca 2+ + 2OH-
Al(OH)3 → Al3+ + 3OH-
Bronsted-Lowry Theory of Acids and
Bases
A Bronsted-Lowry Acid is a substance the donates
a proton or H+ ion to another molecule or ion.
• A Bronsted-Lowry Base is a substance that
accepts a proton or H+ ion from another molecule
or ion.
• An Acid-Base reaction must include both a proton
donor and a proton acceptor.
Examples of Bronsted-Lowry Acid Base
Reactions
Acetic Acid reacts with Ammonia to
produce Ammonium Acetate.
CH3CO2H + NH3 ↔ NH4CH3COo
CH3CO2H + NH3 ↔ NH4+ + CH3COo -
Lewis Acid and Base
• Lewis Acid-an atom, ion, or molecule that accepts
an electron pair to form a covalent bond
H+ + NH3 -> NH4+
• Lewis Base-an atom, ion, or molecule that
donates an electron pair to form a covalent bond
BF3 + F- -> BF4+
• Lewis Acid-Base Reaction-the formation of one or
more covalent bonds between an electron-pair
donor and an electron-pair acceptor
• Do page 482
Conjugate Acids and Bases
A conjugate acid is an acid (as a product) that
results from a Bronsted acid-base reaction.
• A conjugate base is a base (as a product) that
results from a Bronsted acid-base reaction.
• A Bronsted Base accepts a proton to produce its
conjugate acid.
• A Bronsted Acid donates a proton to produce its
conjugate base.
Example
CH3CO2H + NH3  NH4+ + CH3COoThe Bronsted acid is CH3COoH.
The Bronsted base is NH3.
The conjugate acid is NH4+.
The conjugate base is CH3COo-.
Weak Acids and Bases
Weak Acids do not fully dissociate when
dissolved in water. This means that the
concentration of the H+ ion is less than expected
due to the stoichiometry.
• Examples of Weak Acids:
Acetic Acid (CH3CO2H)
Phosphoric Acid (H3PO4)
• Weak Bases do not fully accept one H+ per
molecule as expected.
Strong and Weak Acids and Bases
Strong Acids fully dissociate in water, and
donate all possible H+ ions readily.
• Strong Bases fully accept all donated H+
ions readily.
• Weak Acids do not readily donate all
possible H+ ions.
• Weak Bases do not readily accept all
possible H+ ions.
Buffers
• A buffer is a solution that resists changes in the
pH.
• A buffer is made from a weak acid and its
conjugate base or a weak base and its conjugate
acid.
• Since it is made of both a weak acid and a weak
base, there is both acid and base in solution to
react with any strong acid or base that is added to
it. This minimizes the effect of adding acid or
base.
A .001 M acetic acid solution is
5% dissociated. What is the H+
concentration?
% Dis. = M of dissociated acid(100)
M of acid total
CH3CO2H ↔ H+ + CH3CO2% Dis. = [CH3CO2-](100)
[CH3CO2H]
Ionization Constant
HA ↔ H+ + AKa= [H+]•[A-]
[HA]
Ka = acid dissociation constant
A- + H2O ↔ HA + OH
Kb= [HA]•[OH-]
[A-]
Kb = base dissociation constant