分析化学英文课件04 酸碱平衡Acid-Base Equilibria

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Autoprotolysis
As the elemental unit of positive charge, the proton possesses a charge density which makes its independent existence in a solution extremely unlikely. Thus, in order to transform HB into B-, a proton acceptor must be present. Often as in the dissociation of acetic acid in water, this base may be the solvent(H2O) itself.
HAc = H+ + Ac- , H2O + H+ = H3O+
Combination of this two equations, HAc + H2O = H3O+ + Ac-
The protonated water molecule or hydrated proton H3O+, may be called a “hydronium ion”, but it is usually designated simply by “hydrogen ion” and often written as “H+”.
According to Arrhenius theory, acids dissociate into
hydrogen ions and anions, and bases dissociate into
hydroxide ions and cations.
HA = H+ + A BOH = OH- + B+
Dissociation constants of various acids or bases (Ka or Kb) quantitatively describe the strengths of their acidity or basicity. Consider HAc:
HAc + H2O = H3O+ + Ac-,
Conjugate pairs:
Acids and bases that differ only in number of protons (H+) are called conjugate pairs.
Here are some examples: (Conjugate Base)
HCN (acid) and CNHCl and Cl-
acid1 = H+ + base1, base2 + H+ = acid2
We have the general equation for the reaction of acid1 with base2:
acid1 + base2 = acid2+ base1
Obviously, the products must be more weakly acidic and basic than the reactants.
Kb = Kw / Ka =10-14 / 5.6 10-10 = 1.8 10- 5
For polyprotic acids, such as H3A, the
corresponding relations between Ka and Kb should be:
Ka1
H3A
H+ + H2A-
This simple definition omits one type of acids, often called Lewis acids, that do not give up protons, yet have the properties and undergo the reactions of acids. Lewis acids are substances such as AlCl3 and BF3, that have unfilled electron shells and react by accepting an electron pair from a base.
Ka = [H +][A -] / [HA] Kb =[B+][OH -] / [BOH]
Arrhenius theory has long been applied to aqueous
systems, but it cannot be extended to nonaqueous
solutions. In addition, the theory cannot clearly
• Determine which of several possible species predominate at a given pH. • Aid in selecting the regions of buffer effectiveness for mixtures of acids or bases and their salts. • Calculate the concentration of a particular species in the solution at a given pH. •Possibility of stepwise titration and its error.
Ka
[H ] [Ac ] [Ac ]
For the conjugate base, Ac-:
Ac-+ H2O = HAc + OH -,
[HAc]ween Ka and Kb
Obviously, relationship between Ka and Kb for the acidbase conjugate pair can be expressed as :
KaKb = [H+][OH -] = Kw = 10-14
Example
Dissociation of NH4+, Ka= 5.6 10-10, calculate the dissociation constant Kb of the conjugate base NH3 .
Solution:
The status of dissociation of common acid-base species is a function of pH, can be expressed by graphs of distribution.
Application of the distribution graphs:
Kb3
H2A-
Ka2
Kb2
HA2-
Ka3
HA2-
Kb1
H++A3-
Ka1 Kb3 = Ka2 Kb2 = Ka3 Kb1 = Kw
§4.3 Distribution of Acid-Base Species as a Function of pH
It may be seen from acid-base dissociation that H3+O and various species of acid or base exist in an aqueous system as the conjugate acid-base pairs are at equilibrium.
The Bronsted definition of acids and bases 1eads to the simple relationship:
acid1 = H+ + base1,
base2 + H+ = acid2
Combining these two equations:
The Bronsted definition of acids and bases 1eads to the simple relationship:
of accepting a proton: HB + S = HS+ + B -
Similarly, for the proton transferring process in an
aqueous solution of the base, the solvent molecule(H2O) must be also present. For instance:
Fundamental acid-base equilibria are important in understanding acid-base titration and the effect of acids on chemical species and reactions.
In this chapter, we discuss the principles of acidbase equilibria and the related calculations.
NH3 H2O NH3
+ = +
H+ = H+ +
H2O
NOHH4+= OH
-
+
NH4+
A substance that can act either as an acid or as a base
is said to be amphiprotic. Many solvents are amphiprotic.
constant, KW is called autoprotolysis or self-ionization constant of water.
KW = [H+][OH - ], at 25C,
Kw = 10 -14 And pKw= 14
§4.2 Dissociation Equilibria of Acids and Bases
Analytical Chemistry 分析化学
Chapter 4 Acid-Base Equilibria
§4.1 Aqueous Acid-Base Theories §4.2 Dissociation Equilibria of Acids And
Bases §4.3 Distribution of Acid-Base Species as a
Function of pH §4.4 pH Calculations for Aqueous Solutions §4.5 Activity and Activity Coefficient
§4.1 Aqueous Acid-Base Theories
The acidity or basicity of a solution is frequently an important factor in chemical reactions.
For example, water acts as a base toward acids and as an
acid toward bases.
The proton transferring reaction between the water
molecules is called autoprotolysis, and the equilibrium
describe the basic properties of substances such as NH3 and Na2CO3.
2. Bronsted theory
Acid-base definition, According to Bronsted, an acid is a substance that can give up protons and a base is any compound or ion that can accept protons.
Among the theories for describing acid-base properties of substances, Arrhenius theory and Bronsted theory are widely adopted for aqueous systems.
1. Arrhenius theory
Water is not the only solvent to which acids can
transfer their protons, and we may write a general
dissociation equation where S is any solvent capable
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