ACIDS AND BASES
Topic 8
8.1 Reactions of acids and bases
• Acids with metals
• Produces a salt and hydrogen gas
• Mg + 2HCl  MgCl2 + H2
• Acids with carbonates and
hydrogencarbonates
• Produces salt + carbon dioxide + water
• Na2CO3 + H2SO4  Na2SO4 + H2O + CO2
Con’t
• Acids with bases and alkalis
• Bases are metal oxides
• Produce a salt and water
• CuO + H2SO4  CuSO4 + H2O
• Alkalis are bases that dissolve in
water
• Produce a salt and water
• NaOH + HNO3  NaNO3 + H2O
8.2 Definitions of acids and bases
• BrØnsted-Lowry definitions
• Acid proton (H+) donor
• Base/ alkal proton (H+) acceptor
• Conjugate base the base formed when
an acid reacts and donates a proton to
become a base
• Conjugate acid the acid formed when a
base reacts and accepts a proton to
become an acid
• Referred to as conjugate acid- base pair
Conjugate acid-base pairs
• CH3COOH + H2O  CH3COO- + H3O+
• Which is the beginning acid? Base?
• Which is the acid’s conjugate?
• The base’s?
• Water is called amphoteric. What is
that?
Can act as an acid or base
Another way to phrase it…
• In the forward reaction the CH3COOH
acts as the acid and the H2O acts as
the base
• In the reverse reaction the CH3COOacts as the base and the H3O+ acts as
the acid
Lewis theory of acids and bases
• Acid electron pair acceptor
• Base electron pair donor
• Must understand the Lewis structure of the
compound to know which substance will
accept the electrons
• Ex. NH3 + H+  NH4+
• Which substance gained electrons? Which
donated?
Con’t
• A dative covalent bond is always formed in a
Lewis acid-base reaction
• What is a dative covalent bond?
Both electrons come from the same atom
• For a substance to act as a Lewis base, it must
have a lone pair of electrons
• For a substance to act as a Lewis acid, it must
have space to accept a pair of electrons
8.3 Strong and weak acids and bases
• When acid reacts with water it dissociates or
ionizes
• Can use the Bronsted-Lowry theory to understand this
• Strong acids completely dissociate in aqueous
solution
• Which direction does the equilibrium dominantly
lie?
To the right (products)
• HA  H+ + A• Uses a non-reversible arrow
Strong acids
• HCl is considered a monoprotic acid it
dissociates to form one proton per molecule
• H2SO4 is considered diprotic dissociates to
form two protons per molecule
• H2SO4 + H2O  HSO4- + H3O+
• HSO4- + H2O  SO42- + H3O+
• Sulfuric acid is only considered a strong acid in
the first dissociation
Weak acids
• Only partially dissociate in aqueous
solution
• The equilibrium arrow is used for these
equations
• HA  H+ + A• Ex. Carbonated water is acidic due to
dissolved CO2, which acts as a weak acid
Bases
• Strong bases ionize completely in aqueous
solution
• Ex. NaOH  Na+ + OH• The group 1 hydroxides are strong bases;
along with Ba(OH)2
• Weak bases ionize partially in aqueous
solution
• Equilibrium arrows are used in these
equations
• Ex: NH3 + H2O  NH4+ + OH-
Distinguishing experimentally between
strong and weak acids and bases
• Solutions of strong acids conduct electricity better than
weak acids
• Why?
There is a large concentration of ions to carry the electrical
current
• Can also be called strong electrolytes or weak electrolytes
• The same concept is true for strong and weak bases
Con’t
• Strong acids have a lower pH than weak acids
• What does pH measure?
The concentration of H+ ions in solution
• Lower pH = more H+ ions
• Would the pH for strong bases be higher or
lower?
Higher
• Why?
There are very few H+ ions in the solution of
strong bases
Con’t
• Strong acids react more violently with
metals or carbonates
• The higher concentration of free H+ ions
cause a more rapid reaction with metal to
form H2(g)
• There is a similar effect when a carbonate
is added
Con’t
• strength vs. concentration
• Concentration refers to the number of moles of
acid in a certain volume (i.e. mol dm-3)
• Strength refers to what?
How much the acid dissociates
• Ex. Ethanoic acid is considered a weak acid. No
matter how concentrated the acid solution is, it
will still not fully dissociate.
• Similar for bases
8.4 pH
• Definition: pH is the negative logarithm to base 10
for the hydrogen ion concentration in aqueous
solution
pH= -log10 [H+(aq)]
• The pH scale is used to indicate how acidic or
alkaline a solution is
• The scale is from 1 to 14
• One being the most acidic
• Fourteen is the most alkaline
• Seven is neutral
pH
• Since pH is on a log scale, a 1 unit change
in pH means there is a tenfold change in H+
ion concentration
• Calculating [H+] from pH
[H+]= 10-pH
• This is the inverse of the previous equation
Calculating pH of a strong acid
• It can be assumed that a strong acid fully
dissociates and the [H+] is equal to [acid]
• Ex: calculate the pH of a 0.00150 mol dm-3
solution of HCl.
• pH=-log10[H+]= -log[0.00150]= 2.82
• Just plug in the [acid] for hydronium ions
pH is not a measure of acid strength
• This is the measure of what?
[H+] ions
• It is possible for a dilute solution of a strong
acid to have a higher pH than a
concentrated solution of a weak acid
• pH can be used to compare the strength of
acids,ONLY IF THE CONCENTRATIONS
OF THE ACIDS ARE EQUAL