acid-base equilibria
DESCRIPTION
Acid-Base Equilibria. REVIEW. Electrolyte : Substances that dissolves in water to produce solutions that conduct electricity Nonelectrolytes : Substances whose aqueous solutions do not conduct electricity Strong and weak relates to the degree of dissociation or ionization - PowerPoint PPT PresentationTRANSCRIPT
Acid-Base EquilibriaAcid-Base Equilibria
REVIEWREVIEW
• Electrolyte: Substances that dissolves in water
to produce solutions that conduct electricity
• Nonelectrolytes: Substances whose aqueous
solutions do not conduct electricity
• Strong and weak relates to the degree of
dissociation or ionization
• In aqueous solutions protons occasionally exist
as hydrated molecules, H(H2O)n+ (n=1)
Definitions:
Arrhenius: An acid is a substance that increases the H+ (or H3O+)
concentration in an aqueous solution.
HCl + H 2O H3O+ + Cl-
A base is a substance that increases the OH- concentrationin an aqueous solution.
HCl H+ + Cl-
NaOH(s) Na+ + OH-
What about Na2CO3 ????
Bronsted-Lowry:
Acid = a proton donor in a Reaction
Base = a proton acceptor in a Reaction
HCl(aq) + NaOH(aq) → HOH + NaCl
Lewis:
An acid is an electron pair acceptor
H+
acid
A base is an electron pair donor
:O:H-....
H:O:H..
..
water
Acid/Base reactions:
Produce water and a salt (and sometimes carbon dioxide).
Hint: concentrate on the water first. Remember, water has the formula HOH.
Complete and balance the following:
HCl + KOH
HCl + Ca(OH)2
HOH + KCl
2
Require equal numbers
2HOH + CaCl2
1. Ba(OH)2 + H3PO4
2. HC2H3O2 + NaOH
3. H2SO4 + KOH
4. H2CO3 + NaOH
5. Na2CO3 + HCl
6. NH4OH + H2SO4
7. NH3 + HCl
Give a definition of an acid:
An acid is a proton donor (H+)
Give a definition of a base:
A base is a proton acceptor
Conjugate acids and Conjugate bases
HCl + KOH HOH + KClacid base conj. acid conj. base
Na2CO3 + 2HCl H2CO3 + 2NaCl base acid conj. acid conj. base
Na2CO3 + 2HCl H2O + CO2(g) + 2NaCl conj. baseconj. acidacidbase
NH3 + HCl NH4+ + Cl-
What is a strong Acid?
An Acid that is 100% ionized in water.
Strong Acids:100% ionized (completely dissociated) in water.
HCl + H2O H3O+ + Cl-
often written as:
HCl H+ + Cl-
Strong Acids:Perchloric HClO4
Chloric, HClO3
Hydrobromic, HBrHydrochloric, HClHydroiodic, HINitric, HNO3
Sulfuric, H2SO4
Strong Acids:100% ionized (completely dissociated) in water.
HCl + H2O H3O+ + Cl-
What is a strong Base?
A base that is completely dissociated in water (highly soluble).
NaOH(s) Na+ + OH-
Strong Bases:
Group 1A metal hydroxides(LiOH, NaOH, KOH,RbOH, CsOH)
Heavy Group 2A metal hydroxides[Ca(OH)2, Sr(OH)2, andBa(OH)2]
Weak Acids: “The Rest”
Strong Acids:100% ionized (completely dissociated) in water.
HCl + H2O H3O+ + Cl-
Note the “one way arrow”.
Weak Acids:Only a small % (dissociated) in water.
HC2H3O2 + H2O H3O+ + C2H3O2
-
Note the “2-way” arrow.
Why are they different?
Strong Acids:
HCl HCl HClHCl HCl
ADD WATER to MOLECULAR ACID
(H2O)
Strong Acids:
(H2O)H3O+
H3O+
H3O+
H3O+
H3O+
Cl-
Cl-
Cl-
Cl-
Cl-
Note: No HCl molecules remain in solution, all have been ionized in water.
HC2H3O2
HC2H3O2
HC2H3O2
HC
2H3O
2
HC2H3O2
(H2O)
Weak Acid Ionization:
Add water to MOLECULES of WEAK Acid
HC2H3O2
HC2H3O2
HC2H3O2
HC
2H3O
2
HC2H3O2
H30
+ C2H3O2-
(H2O)
Weak Acid Ionization:
Note: At any given time only a small portion of the acid molecules are ionized and since reactions are running in BOTH directions the mixture composition stays the same.
This gives rise to an Equilbrium expression, Ka
H30+ C2H3O2
-HC2H3O2
ACID-BASE CONCEPTSACID-BASE CONCEPTS
• Acid-base reactions may be one of the most
important class of reactions
• The most basic of the acid-base concepts is
the Arrhenius theory
• Acids are substances that dissociate in water
to produce hydronium ions, H3O+ and bases are
substances that dissociate in water to produce
hydroxide ions, OH-
ACID-BASE STRENGTHACID-BASE STRENGTH
• A strong acid is almost completely dissociated
or ionized
• A weak acid is only partially dissociated or
ionized
HCl(aq) H+(aq) + Cl- (aq)
HF(g) + H2O(l) H3O+(aq) + F- (aq)
HCl(aq) H+(aq) + Cl-(aq)
• So each reaction is an equilibrium with a K
STRONG ACIDS AND BASESSTRONG ACIDS AND BASES
• Common strong acids are either monoprotic or
diprotic
HA(aq) + H2O(l) H3O+(aq) + A-(aq) 100%
• Means [HA] = [H3O+]
• Similar situation for strong bases
• Typical strong bases: Group 1A metal and Ca,
Sr, and Ba hydroxides
• Strong acids: HX, HNO3, HClO3, HClO4, H2SO4
WEAK ACIDSWEAK ACIDS
• Weak acids and bases are weak electrolytes
• For a weak acid, HA
• Ka: the dissociation or ionization constant
• For acetic acid
[H3O+][A- ]
Ka = [HA]
HA(aq) + H2O(l) H3O+(aq) + A- (aq)
CH3COOH(aq) + H2O(l) CH3COO- (aq) + H3O
+(aq)
Ka = [H3O+][CH3COO-]/[CH3COOH]
• Acids with larger ionization constants ionize or
dissociate to a greater extent than acids with
smaller ionization constants
• The larger the value of Ka, the higher [H3O+]
and the stronger is the acid
HIO3; Ka = 1.6 x 10-1
CH3COOH; Ka = 1.8 x 10-5
HCN; Ka = 6.2 x 10-10
AUTOIONIZATION OF WATERAUTOIONIZATION OF WATER
• Reaction called autoionization of water
K = [H3O+][OH-]/[H2O]2
K[H2O]2 = [H3O+][OH-]
Kw = [H3O+][OH-]
• Kw is called ion product of water
• At 25 °C, Kw = 1.0 x 10-14
• Valid also for dilute aqueous solutions
H2O(l) + H2O(l) H3O+(aq) + OH- (aq)
THE pH SCALETHE pH SCALE
• The hydronium ion concentration is a measure
of a solution’s acidity
• Usually small numbers
• The pH scale is used express acidity and
basicity
pH = -log[H3O+] so [H3O+] = 10-pH
• Note that as pH increases [H3O+] decreases
• Value of Kw varies with temperature
• “Acidity” depends on the concentration of
H3O+ ions
Acidic: [H3O+] > [OH-]
Basic: [H3O+] < [OH-]
Neutral: [H3O+] = [OH-]
• Notice that neutral does NOT necessarily
mean pH 7
• pH is usually quoted with the same number of
significant digits as the concentration
• pH can be measured using a pH meter or by an
acid-base indicator
• An indicator is usually an organic acid that have
different colors in solutions of different pH
• Indicators exist that cover the entire pH scale
e.g. bromothymol blue: 6.0 – 7.6
phenolphthalein: 8.0 – 10
HI n(aq) + H2O(l) H3O+(aq) + I n- (aq)
• The p-scale can also be applied to
ionization
constants
pKa = - log Ka
• The larger the value of Ka the smaller
the value of pKa and the stronger the acid
• From the acid-dissociation constant we
can
calculate equilibrium concentrations as
well as
pH
BRØNSTED-LOWRY THEORY
• An acid is a proton donor
• A base is a proton acceptor
• An acid-base reaction is the transfer of a
proton from an acid to a base
• Reactions can be described in terms of what
are called conjugate acid-base pairs
- species which differ by a proton
- a charge
- B is a proton acceptor; it is a base
- BH+ is a proton donor; it is an acid
So A- is the conjugate base of HA and
BH+ is the conjugate acid of B
• HF/F- and H3O+/H2O are conjugate pairs
HA + B BH+ + A-
Conjugate acid- basepairs
HF(g) + H2O(l) H3O+(aq) + F- (aq)
Relationship of Ka and concentration
• In dilute solutions of weak acids, the assumption
- is that all the H3O+ is coming from the acid
- the concentration change of a species is small
compared to the initial concentration of that
species
- at equilibrium, [HA] [HA]init
• Always be cautious when using the assumptions
given above
%-dissociation
• Amount of acid that dissociates can be
expressed as a percent
• In general %-dissociation increases with the
value of Ka
• For any weak acid, HA, %-dissociation
increases with dilution
%- Dissociation = ___________ x 100[HA]dissociated
[HA]initial
POLYPROTIC ACIDSPOLYPROTIC ACIDS
• Polyprotic acids: Acids that provide more than
one hydronium ion in solution
e.g. H2SO4, H3PO4
• Polyprotic acids ionize in a stepwise manner
• Consider sulfuric acid, H2SO4
H2SO4(aq) + H2O(l) H3O+(aq) + HSO4
- (aq)
HSO4- (aq) + H2O(l) H3O
+(aq) + SO42- (aq)
• In general, Ka1 > Ka2
WEAK BASESWEAK BASES
• Weak bases undergo equilibria in water
• For a general weak base, B
• Kb is called the base-dissociation constant
NH3(aq) + H2O(l) NH4+(aq) + OH- (aq)
B(aq) + H2O(l) BH+(aq) + OH- (aq)
[BH+][OH- ]Kb = [B]
• For conjugate acid-base pairs the product of
their equilibrium constants is the ionization
constant for water
Ka x Kb = Kw
HN
OH
H
Hydroxylamine
HN
CH3
HMethylamine
FACTORS AFFECTING ACID STRENGTH
• The strength of an acid depends on the polarity
of the H-E bond
• The polarity of the bond is related to the bond
strength of H-E
• The weaker the H-E bond the stronger the
acid
• Take the hydrohalic acids:
HX; X = F, Cl, Br, I
HF << HCl < HBr < HI
• For binary acid in the same group H-A bond
strength determines acid strength
• Same applies to other groups:
H2O < H2S < H2Se
• For binary acids in the same row polarity of the
H-E bond determines acid strength
:- acid strength increase with the
electronegativity of E
• The oxoacids are also important
e.g. HNO3, H2SO4, HClO4
• The acidity is dictated by factors that affect
the O-H bond strength
1. The electronegativity of the central element
2. The oxidation number of the central element
SO
OO
O
H
HO N
O
O
H
H O I H O Br H O Cl< <hypoiodous acid 2.5
hypobromous acid 2.8
hypochlorous acid 3.0
HClO4 > HClO3 > HClO2 > HClO
+7 +5 +3 +1
• Some small, highly charged metal ions are quite
acidic
:- they are hydrated and transfers a proton
to a water molecule
SALTS: ACID-BASE PROPERTIESSALTS: ACID-BASE PROPERTIES
• Salts: an ionic compound that is formed when
an acid neutralizes a base
NaOH(aq) + HCl(aq) NaCl(aq) + H2O
• Aqueous solutions of salts can be neutral, acidic
or basic
Strong acid + strong base neutral solutions
Strong acid + weak base acidic solutions
Strong base + weak acid basic solutions
• When a salt dissolves in water, its constituent
ions may react with water – reaction called
hydrolysis
NEUTRAL SOLUTIONS
• Salts of strong acids and strong bases
e.g. NaCl
• Because the ions do not hydrolyze
• Cl- is the conjugate base of HCl – it is a weak
base
• The same argument is made for Na+
• Essentially [H3O+]/[OH-] ratio does not
change
• {Group 1A metals, Ca2+, Sr2+, Ba2+} and {I-,
Br-, Cl-, NO3-, ClO4
-} does not hydrolyze
Cl-(aq) + H2O(l) HCl(aq) + OH- (aq)
Cannot occur
Acidic Solutions
NH3(aq) + HCl(aq) NH4Cl(aq)
• Salts from weak bases and strong acids
• In aqueous solution NH4+ undergo hydrolysis
- the chloride ion does not
• The generation of H3O+ from the reaction
makes these solutions acidic
NH4+(aq) + H2O(l) NH3(l) + H3O
+(aq)
Basic Solutions
• Salts from strong bases and weak acids give
basic solutions
• This basic anions of weak acids hydrolyze to
form hydroxide ions
NaCH3COO CH3COO- + H3O+
CH3COO- (aq) + H2O(l) CH3COOH(aq) + OH- (aq)
SALTS FROM WEAK ACID AND BASESSALTS FROM WEAK ACID AND BASES
e.g. NH4CH3COO, ammonium acetate
• Aqueous solution of the salts may be basic,
acidic or neutral
• The pH depends on the relative Ka and Kb of
the parent acid and base
• Consider the case when Ka = Kb
NH4CH3COO(aq) NH4+(aq) + CH3COO-(aq)
• Both ions can undergo hydrolysis
• Hydrolysis constant for the acetate ion (Kbh) is
equal to the hydrolysis constant for the
ammonium ion (Kah)
• The same concentration of H3O+ as OH- is
produced solution is NEUTRAL
NH4+(aq) + H2O(l) NH3(l) + H3O
+(aq)
CH3COO- (aq) + H2O(l) CH3COOH(aq) + OH- (aq)
• If the parent Ka > Kb: solutions acidic
e.g. NH4F
Ka(HF): 7.2 x 10-4 > Kb (NH3): 1.8 x 10-5
Kb(F-): 1.4 x 10-11 < Ka(NH4+): 5.6 x 10-10
• So NH4+ hydrolyzes to a greater extent than F-
more H3O+ is produced than OH-
NH4+(aq) + H2O(l) NH3(l) + H3O
+(aq)
F- (aq) + H2O(l) HF(aq) + OH- (aq)
• If the parent Ka < Kb: solutions basic
e.g. NH4CN
Ka(HCN): 4.9 x 10-10 < Kb (NH3): 1.8 x 10-5
Kb(CN-): 2.0 x 10-10 > Ka(NH4+): 5.6 x 10-10
• So NH4+ hydrolyzes to a greater extent than F-
more OH- is produced than H3O+
NH4+(aq) + H2O(l) NH3(l) + H3O
+(aq)
CN-(aq) + H2O(l) HCN(aq) + OH-(aq)
LEWIS THEORYLEWIS THEORY
• The Lewis acid-base theory is the most
complete that were are going to see
• In this theory:
- an acid is any species that can accept an
electron pair
- a base is any species that can donate an
electron pair
• The Lewis acid typically have vacant orbitals
• Common Lewis acids: metal ions and electron
deficient molecules such as AlCl3, BF3, etc.
• A typical Lewis base would be ammonia, NH3
as well as the common anions
• Acid-base reactions:
BF3 + NH3 → [F3B←NH3]
acid base acid-base adduct
AlCl3 + Cl- →[AlCl4]-