Acids and Bases
I.
Bronstead-Lowry acids and bases:
A. Transfer of protons reactions according to this view:
1. Bronstead-Lowry acid- a proton donor.
2. Bronstead-Lowry base- a proton acceptor.
3. Used more than Arhenius definition because not all reactions are carried out in
water.
B. Conjugate acids and bases:
1. Conjugate acid-base pair- two substances that differ from each other by only one
proton.
a) Conjugate acid- is the proton donor of the pair.
b) Conjugate base- is the pair’s proton acceptor.
C. Amphoteric substances can behave as either acids or bases:
1. Amphoteric substance- a substance that can either be an acid or a base depending
on the other substance present.
a) Water is a common example, it can either accept a proton, becoming H3O+,or
donate a proton, becoming OH-.
II.
Strengths of Bronstead acids and bases follow periodic trends:
A. Standards:
1. Strength of the acid or base is determined by the degree to which the species can
carry out its reaction to completion.
a) So most of the classifications that held in the Arhenius classification still stand.
2. A reference base or acid is used to establish a standard.
a) Usually water because most reactions are carried out in an aqueous medium,
and it is amphiprotic.
B. H3O+ & OH- are the strongest acid and base that can exist in the presence of water:
1. Because any strong acid when placed in water will always react completely to yield
H3O+.
a) This means that once the acid is placed in solution, the only acid present is
hydronium ion.
b) Therefore, a reference base that is less willing to accept protons.
2. By the same logic, hydroxide ion will be the strongest base in aqueous solution
because the strong base will react to completion leaving only hydroxide ions.
a) Must use a reference acid less willing to donate protons.
C. Compare strength of conjugate pairs:
1. The equilibrium is the indicator to tell the relative strength of the conjugate pairs.
a) The position of acid-base equilibrium favors the weaker acid and base.
2. Stronger acids and bases react with each other to produce weaker conjugates.
a) This results in a reciprocal relationship of strength between conjugate pairs.
Strong acid weak base as a conjugate pair.
Strong base weak acid (conjugate)
D. Binary acid trends:
1. Binary acids- acids formed between hydrogen and nonmetals have the form of HX,
H2X, H3X, etc…
a) The strength of the acid increases as you go from left to right on the same
period.
i.
This is due to the fact that the electronegativity increases in this
manner, and this results in a greater degree of separation between
elements in a bond. This leads to a larger partial positive charge, and
thereby making it easier to abstract the proton.
b) The strength of the acid increases as you descend down a group of the periodic
table.
i.
Opposite the trend you would expect based on electronegativity, due to
the fact that atomic size takes precedence over electronegativity.
a) Smaller atoms form stronger covalent bonds, and therefore, are
harder to break the bond.
b) Larger atoms form weaker bonds and are easier to break
E. Oxoacids trends:
1. Oxoacids- hydrogen, oxygen, and some other element.
2. The strength of these oxoacids with the same number of oxygens depends on the
electronegativity of the atom that is attached to the hydroxide.
a) If the atom is very electronegative, the polarity of the OH bond increases the
charge separation of the bond thereby making it easier to abstract the proton.
3. The number of lone oxygen atoms also dictates the strength:
a) As the number of lone oxygens, oxygens unbound to hydrogen, increases, the
better the donor.
i)
Comparing oxoacids with same central atoms,
ii)
This is due to the electronegativity of the oxygen atoms ability to pull
the electron density away from the central atom results in the central
atom becoming more electronegative. This eases the abstraction of the
hydrogen attached to the hydroxide bond.
b) The number of lone oxygens also dictates the conjugate base strength:
i)
The more lone oxygens the easier it is to share the charge of the
conjugate on base. This results in a decreased ability to accept a
proton, in other words, a weaker conjugate base.
III.
Lewis acids and Bases:
A. Basics:
1. Lewis acid- any ionic or molecular species that can accept a pair of electrons in the
formation of a coordinate covalent bond.
a) May or may not have a filled valence shell.
i.
Some atoms that are able to exceed the octet rule that have one or
more pie bonds, are able to move the electrons around to form a
coordinate bond. They use the partial charge on the central atom to
attract the donating Lewis acid.
2. Lewis base- any ionic or molecular species that can donate a pair of electrons in the
formation of a coordinate covalent bond.
a) These substances usually have complete valence shells and an extra unshared
pair of electrons.
3. Neutraliztion- the formation of a coordinate covalent bond between donor(base)
and acceptor(acid)
IV.
Elements and their Oxides acid-base properties:
A. Locations:
1. Elements most likely to form acids are found in the upper right-hand corner of the
periodic table.
2. Elements most likely to form basic hydroxides are grouped in one general location in
the table, among the metals, particularly those in groups IA, alkali metals, and IIA,
alkali earth metals
3. We can classify the elements based their oxide’s abilities to react with water to form
acids and bases:
a) Metal oxides react with water to form bases,
i.
Some metal oxides have their oxygen so tightly bound that the oxygen
remains bound to the crystal even in water. So a stronger acid must be
present, for example in the removal of rust, Fe2O3 a metal oxide.
B. Hydrated metal ions can behave as weak acids:
1. When metal ions are placed in water, the water acts as a Lewis base, donates a pair
of electrons to form a bond with the metal ion, the Lewis acid.
a) The resulting hydrated ion is a Bronstead-Lowry acid with the fundamental
reaction as follows:
M(H2O)n+ + H2O MOH(n-1)+ + H3O +
*Therefore, a hydrated metal donates protons*
2. The degree to which metal ions produce acidic solutions depends on two principal
factors:
a) The amount of charge the cation has,
i.
Again, the larger the positive charge, the easier it is to abstract the
proton.
b) The size of the ion,
i.
the smaller the ion, the more concentrated the charge. The more
concentrated the charge the easier the ion is able to abstract the
electrons from the hydroxide during hydration of the ion.
3. Charge density is a way of considering size and volume of the ion:
Charge density =
*The higher the charge density the more acidic the solution*
*Very small ions with large positive charges have high charge densities*
C. Periodic trends in the acidity of metal ions:
1. The ions size follows the same trends as the elements themselves, Inc as you go left
in a period and down a group.
2. Top of a group has the smallest volume and the largest charge density, and
therefore are the most acidic in the group.
3. Some of the transition metals are acidic, if the charge of the ion is +3 as in Fe or Cr
in water will behave as the hydrated metal ions(BL acid)
D. Acid-base properties depend on the oxidation number:
1. The metal oxides that have small positive charges behave as bases,
a) In actuality the hydrated metals are amphoteric depending on their charge,
V.
pH:
A. Water as very weak electrolyte:
1. Weak conductor of electricity because little self- ionization,
H2O + H2O H3O+ + OH –
a) With a little manipulation of the equilibrium law, and knowing that pure water
has a molar concentration of 55.6 M leads to the following KW:
KW = [H+][OH-] = 1 X 10-14 at 25 degrees Celsius.
2. Solutes will alter the proton and hydroxide concentration, however, the KW, the
product of the [H+][OH-] , remains unaffected, 1 X 10-14 at 25 degrees Celsius.
B. Differentiating between acidic, basic, and neutral solutions:
1. Neutral solution- is a solution in which the molar concentrations of H3O+ and OHare equal.
2. Other classifications:
a) Acidic- [H3O+]>[OH-]
b) Basic- [H3O+]<[OH-]
C. pH is the common scale for finding pH:
1. This is an easier way to determine the acidity of a solution:
pH = - log [H+]
2. The basicity of a solution can be obtained in a similar manner:
pOH = - log [OH-]
3. By taking the log of the KW equation listed above, the equation can be rewritten as
follows:
pH + pOH = 14
4. pH as a means of grouping:
a) pH < 7.00 is an acidic solution,
b) pH = 7.00 is a neutral solution,
c) pH > 7.00 is a basic solution.
5. Measure using electrodes that are sensitive to [H+] concentrations that are accurate
up to or dyes responsive to different pH’s such as litmus paper.
VI.
Strong acids and bases are fully dissociated in solution:
A. Strong acids and bases are fully dissociated : (for now consider monoprotic acids only)
1. Take the concentration of the acid to be that of the concentration of the protons:
a) [H+] = [HA]
2. Take the concentration of the base to be that of the concentration of the hydroxide
ion:
a) [OH-] = [BOH]
3. Assume that the concentration of protons and hydroxide ions from the
autoionization of water are negligible. This is due to the fact that these amounts of
ions are reduced when the solute is introduced into the solution.
a) The protons and hydroxide ions produced by the solute increases the overall
concentrations of these ions, and according to LeChatlier, the reaction shifts the
equilibrium to the removal of these ions, in other words, suppress the further
formation of ions.