# L5 - Acid-Base Properties of Salts and Lewis Definition

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```					   Acid-Base Properties of Salt             A salt is simply another name for
Solutions and Lewis Acid-           an ionic compound.
Base Theory                      Remember, that most “salts” are
strong electrolytes that completely
dissociate in solution.
By dissociation, salts can
sometimes affect pH by increasing H+
or OH- concentrations, donating or
accepting a proton, in side reactions.

The reactions of ions with water     Salts that form Neutral Solutions
are frequently called hydrolysis            We know from the Bronsted-
reactions.                              Lowry theory that conjugate bases of
strong acids have no affinity for
For Example:                            protons in water.
NaF(s) → Na+(aq) + F-(aq)        Example:
No apparent change in pH; However,         HNO3 + H2O → H3O+ + NO3-
If NO3- had any affinity for H+,
F- + H2O ⇄ HF + OH-              then this reaction would be an
equilibrium reaction.

By this logic, we should              Salts That Form Basic Solutions
understand that the addition of salts   Notice what happens when we put
containing cations of strong bases      sodium acetate in solution
and anions of strong acids have no
NaCH3CO2 → Na+ + CH3CO2-
effect on the pH of a solution.
And then,
i.e.
CH3CO2- + H2O ⇄ CH3CO2H + OH-
AgNO3 → Ag+ + NO3-
What happed to the pH when the

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In this situation, the acetate anion is   1. What is the pH of a solution made
acting as a weak base. Because we         by adding 3.0 g of sodium acetate to
know the Ka of the conjugate acid,        make 5.0 L solution?
acetic acid, we can easily calculate
the Kb of the acetate anion from:
Notice, the solution is very basic.
Kb = Kw / Ka
= 1.0x10-14 / 1.8x10-5   = 5.6x10-10   Salts derived from a strong base and
By knowing the dissociation           a weak acid create basic solutions.
constant, we can now make                 Examples: NaClO, LiF, etc…
calculations of the solution knowing
the initial concentration.

Salts That Form Acidic Solutions               A second type of salt that produces
an acid solution is one that contains a
Observe:                                  highly charged metal ion.
i.e. Al3+
NH4Cl + H2O → NH4+ + Cl-
And then,                          The presence of
aluminum ions in solution
NH4 + + H O ⇄ NH + H O+                results in the formation of
2         3 3
coordination bonds
So, we see that salts derived from a      between the ion and
unshared pairs of electrons
weak base and strong acid create          on the water molecules
acidic solutions                          These compounds are known as metal hydrates.

Because metal ions are positively      Now, Look at Fe3+
charged, they attract the unshared               Fe3+ + 6H2O → Fe(H2O)63+
electron pairs of water molecules.
What is the geometry of the
Notice, Aluminum is of the third       iron hydrate?
period and can have an expanded octet.
It has accepted 6 coordination bonds.

Al(NO3)3 → Al(H2O)63+ + NO3-
Draw the orbital notation for Al3+

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When a water molecule interacts
with the positively charged metal ion,
electron density is drawn from the
oxygen causing the O-H bond to
become more polarized; as a result,      Al(H2O)63+ ⇄ Al(OH)(H2O)52+ + H+
water molecules bound to the metal
ion act as a source of protons:
Al(NO3)3 + 6H2O → Al(H2O)63+ + NO3-

Al(H2O)63+ ⇄ Al(OH)(H2O)52+ + H+

From a known Ka value, the pH          Metal Hydrates increase in their acid
for solutions containing hydrated ions   strength as their charge increases and
can be easily calculated.                atomic radius decreases
Fe3+ is a fairly strong acid compared
2. Calculate the pH and pOH of a         to Al3+
1.0x10-4 M AlCl3 solution. (Ka for
Al(H2O)63+ = 1.4x10-5)                Fe3+ + 6H2O → Fe(H2O)63+ Ka = 2.0x10-3

Wow, metal ions act as Acids!
Except, in most cases, Heavy Group I and
Group II metals.

Salts of binary hydrides (HS-) and
polyprotic oxyacids often have anions
that can form salts that are amphoteric.
Example:
anions of PO4 - PO43- (B-L Base)
- H2PO42-
(B-L Acid or Base)

Therefore, NaH2PO4 could be
acidic or basic depending on relative Ka
and Kb values

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3. Predict whether Na2HPO4 will be           4. Determine if the following salts
acidic or basic in aqueous solution.            will result in a neutral, acidic, or
basic solution when dissolved
a. Ammonium acetate
b.Lithium chloride
c. Iron (II) sulfate
d.Ammonium fluoride
e. Lithium phosphate

Above, we have seen the formation        G.N. Lewis (1920’s)
of a metal hydrate. Interestingly,
metals other than heavy alkali and               Noticed that all bases
contained an unshared electron
alkaline-earth metals form acidic
pair and coordination bonds
hydrates.
always form between this
This behavior is difficult to predict    electron pair and the protons.
however using Arrhenius or Bronsted-
Lowry definitions of acids/bases.            Lewis acid – Electron pair acceptor
Lewis base – Electron pair donor

Observe:                                                 Lewis Acids

BF3 + NH3
..

→ BF3NH3
Notice, no protons       F          H
are transferred       F                     • Lewis acids are defined as electron-pair
and H+ does not               B N             acceptors.
increase                                H   • Atoms with an empty valence orbital can
F
H         be Lewis acids.

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Lewis Bases

• Lewis bases are defined as electron-pair
donors.
• Anything that could be a Brønsted–
Lowry base is a Lewis base.
• Lewis bases can interact with things
other than protons, however.

5. Identify the Lewis acids and bases        6. Describe the electron configuration
in the following                          and geometry of the Beryllium
hydrate using the Lewis theory of
a. Ni 2+ + 6NH3 → Ni(NH3)6 2+                acids and bases.
b.H+ + H2O ⇄ H3O+

c. Fe(ClO4)3(s) + 6H2O(l) ⇄
Fe(H2O)63+ (aq) + 3ClO4- (aq)

5

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