ACIDS & BASES
ACIDS:
1. Have a sour taste
2. Contain H that can be exchanged for active metals with
H2 release:
2HCl(aq)+Mg(s)H2+MgCl2(aq)
3. React with bases forming salt & water (neutralization):
HCl(aq)+KOH(aq)KCl(aq)+H2O
or with metal oxides: 2HCl(aq)+MgO(s)  H2O+MgCl2(aq)
4. Change the color of some dyes: “acid-base indicators”
5. Acids are electrolytes: their solutions conduct
electrical current.
ARRHENIUS definition:
Acids are electrolytes that release H+ ions in
the solution:
HCl(aq)  H+(aq) + Cl-(aq)
All reactions are due to H+& can be presented as net ionic
equations with these ions:
Neutralization:
H+ (aq)+OH-(aq) H2O
or:
H+(aq)+MgO  H2O+Mg2+(aq)
reactions with metals:
Mg+2H+(aq)  H2 +Mg2+(aq)
EQUILIBRIUM
Some rxn’s are irreversible, i.e. they can only go one direction
A B, while some other are reversible, i.e. they can go both
directions:
A
B
If at the beginning only A is present, the forward rxn prevails.
As B is accumulated, the reverse rxn starts.
At some point the rate of forward rxn will be equal to the rate
of the reverse rxn:
R=R
This is the state of
EQUILIBRIUM
We’ll apply this concept to electrolyte solutions:
Ionic compounds, when dissolve, completely dissociate into
hydrated ions:
NaCl(s)  Na+(aq) + Cl-(aq)
Ca(OH)2(s)  Ca2+(aq) + 2OH-(aq)
CuSO4(s)  Cu2+(aq) + SO42-(aq)
All ionic compounds are ELECTROLYTES
Because they completely dissociate into ions, they are
STRONG ELECTROLYTES
Many molecular compounds do not form ions in solutions:
glucose C6H12O6, ethanol C2H5OH, acetone H3C-CO-CH3
They are non-electrolytes
Some molecular compounds, although they do not contain ions
as pure substance, can form ions interacting with water.
Most important of those are ACIDS, which are IONIZED by
+
water releasing H
HCl(g) + H2O  H3O+ + Cl-(aq)
acid ionization
in a simplified form (omitting H2O):
HCl(g)  H+ + Cl-(aq)
There are acids that are completely ionized in water – they are
STRONG ACIDS
HCl, HBr, HI (but not HF !), HNO3, H2SO4, HClO4
H2SO4  H+ + HSO4Other acids are ionized partially: their ionization is a
reversible rxn.
There is an equilibrium between their ions & dissolved
molecules:
H2CO3  H+ + HCO3HClO  H+ + ClOHCN H+ + CN- H2SO3  H+ + HSO3Those are WEAK ACIDS
Acids with 1 displaceable H are monoprotic: HCl, HNO3
Acids with 2 displaceable H are diprotic
H2SO4 + Zn  ZnSO4 + H2
Acids with 3 displaceable H are triprotic
2H3PO4 + 3Ca  Ca3(PO4)2 + 3H2
Strong diprotic acids are completely ionized releasing 1 H+:
H2SO4 H+ + HSO4- ,
but further ionization is always reversible, with equilibrium
shifted to the left:
HSO4-  H+ + SO42-
Strong acids displace weak acids from their salts:
NaCN +HCl  NaCl + HCN
Strong
or:
Weak
CN + H  HCN
-
+
Arrhenius acids are
MOLECULAR
ELECTROLYTES
i. e. they are ionized when interact with a solvent
(solvated in general, hydrated in an aqueous solution).
The minimal solvated hydrogen ion is
HYDRONIUM ION:
HCl(g)+H2O H3O+(aq)+Cl(aq)
Naked proton does not exist in a solution
Arrhenius
BASE
Arrhenius BASES are electrolytes producing
-
hydroxyde
ion, OH ,
in a solution:
+
-
KOH(s)K (aq)+OH (aq)
NH3(g)+H2O [NH4OH] NH4+ +OH-
Strong bases or alkali are soluble metal
hydroxides (K2O + H2O2KOH)
-
+
KOH  OH + K
Weak bases:
NH3(g) + H2O  OH-+NH4+
Bases (when soluble) have bitter taste, feel slippery,
color indicators differently from acids
Neutralization with acids:
KOH + HCl  H2O + KCl
or with non-metal oxides:
KOH + CO2  KHCO3
Most of metal hydroxides are insoluble or low
soluble, except:
Alkaline metal hydroxides, NH4+, alkaline earth
metal hydroxides: Mg(OH)2, Ca(OH)2, Ba(OH)2
are soluble
Be(OH)2 is insoluble
Bases are not only those electrolytes that
release OH-, but any species that accepts H+:
NH3 + HCl NH4Cl or: NH3 + H+ NH4+
Bronsted-Lowry BASE: any acceptor of H+.
Bronsted-Lowry ACID: any donor of H+.
B-L bases have a highly electronegative element with a lone electron pair
:N H 3
LEWIS ACIDS & BASES
Acid –Base interaction is considered as electron pair transfer from
one element, with lone el. pair(s) to another one, with an electron
deficit or positive charge:
+
H
H
|
|
..
.. H–C
..l: + :N – H  :C
..l: + H – N – H
|
|
H
H
Electron pair
acceptor
Electron pair
donor
BF3 + :NH3  F3B – NH3
Cu2+ + :NH3  Cu(NH3)2+
Light
Dark
blue
blue
COMMON ACIDS:
Strong acids Binary acids
Aqueous solutions of
hydrogen halides:
HCl(aq), HBr(aq), HI(aq).
Hydrochloric acid
HF(aq) is a weak acid!
Reason: H-bonds: H-F…H-F…HF…
Ternary compounds:
Oxygenic Acids
+
-
Nitric acid: HNO3  H + NO3
Chloric acid: HClO3 H++ClO3Perchloric acid: HClO4H+ +ClO4-
Sulfuric acid:
diprotic acid
+
H2SO4 H +HSO4-2H++SO42Two-step ionization:
strong in first step, much weaker in second step.
Phosphoric acid:
triprotic acid
+
-
H3PO4H +H2PO4 2H +HPO
+
+
23 3 H +P O
4
3H +PO4
Oxyacids may be considered as a combination
of an non-metal oxide & water:
SO3 + H2O  H2SO4
sulfuric anhydride
P2O5 + 3H2O  2H3PO4
N2O5 + H2O  2HNO3
Weak acids:
Binary: HF  H+ + FTernary: HCN  H+ + CNWeak Oxygenic Acids:
Acetic acid:
C H3 COOH  CH3 COO- + H+
Chlorous:
HClO2  H+ + ClO2-
Nitrous:
HNO2  H+ + NO2-
SO2(g)+H2O  H2SO3 H+ + HSO3Sulfurous

2H+ + SO32CO2(g)+H2OH2CO3H++HCO3Carbonic

2H++CO32H2CO3 & H2SO3 are unstable acids: easily decompose to water
& (gaseous) acid anhydride:
H2SO3  SO2(g)+H2O
H2CO3  CO2(g)+H2O
The only soluble but weak base in this course is ammonium
hydroxide: NH3+ H3O  NH4OH NH4++OH-