Download - Acids and Bases
Acids and bases. Buffers
Brønsted-Lowry theory of acids and bases
acid base
HCl + H2O H3O+ + Cl- NH3 + H2O NH4
+ + OH-
ammonia
H2SO4 + H2O H3O+ + HSO4
- CH3NH2 + H2O CH3NH3+ + OH-
aminomethane methylammonium ion
HSO4- + H2O H3O
+ + SO42-
CH3COOH + H2O CH3COO- + H3O+
ethanoic acid acetate ion
(acetic acid)
Definitions:
A base is a substance that can
accept a hydrogen ion (H+).
An acid is a substance that can donate a hydrogen ion (H+).
Examples:
Brønsted-Lowry theory of acids and bases
When an acid gives up its H+, the resulting anion is a base. This base is called conjugate base.
Acid H+ + Conjugate base
HCl + H2O H3O+ + Cl-
conjugate base
CH3COOH + H2O H3O++ CH3COO-
acetic acid acetate ion; conjugate base
An acid and its conjugate base form conjugate pair.
HCl/Cl- ;
CH3COOH /CH3COO-
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Brønsted-Lowry theory of acids and bases
A base reacts by accepting H+. The species produced by
this reaction is its conjugate acid.
Base + H+ conjugate acid
NH3(aq) + H2O NH4+ + OH-
base conjugate acid
A base and its conjugate acid form conjugate pair.
NH3/NH4+.
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Brønsted-Lowry theory of acids and bases
Substances that can either lose or gain a hydrogen ion are amphoteric substances.
NH3(aq) + H2O NH4+(aq) + OH-(aq)
H+ donor conjugate base
HCl + H2O Cl- + H3O+
H+ acceptor conjugate acid
Water acts as an acid when it donates hydrogen ions to ammonia.
It acts as a base in its reaction with hydrogen chloride.
Water is an amphoteric substance.
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Strength of acids and bases.
strong acids and bases
weak acids and bases
Strong acids are acids that ionize completely in water – HNO3,
HCl, H2SO4, HClO4.
HCl H+ + Cl-
Acids that are partially ionized in water are weak acids.
HCN – hydrocyanate
H2CO3 – carbonic acid
H3PO4, HF, CH3COOH etc.
Most organic acids are weak acids. The only H atom that does ionize to
some degree is the one, bonded to the oxygen atom of the COOH group.
CH3COOH CH3COO- + H+
Brønsted-Lowry theory of acids and bases
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Brønsted-Lowry theory of acids and bases
HCN H+ + CN- 1 mol acid produces 1 mol H+, this acid is
monoprotic acid
H2CO3 H+ + HCO3
- bicarbonate ion
HCO3- H+ + CO3
2- carbonate ion
1 mol acid produces 2 mols H+, this acid is diprotic acid
Polyprotic acids ionize stepwise.
Problem:
Write the equation for the stepwise ionization of the diprotic acid
H2SO3 (sulphurous acid) .
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Brønsted-Lowry theory of acids and bases
Answer:
H2SO3 H+ + HSO3
- hydrogen sulfite ion (bisulfite ion)
HSO3- H+ + SO3
2- sulfite ion
Strong bases produce OH- ions in high
concentration in solutions.
NaOH → Na+ + OH-
Weak bases have low concentration of OH- ions
in their solutions.
NH3 + H2O NH4+ + OH-
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stronger acid than H3O+ weaker base than H2O
conjugate pair
HCl + H2O Cl- + H3O+
conjugate pair
stronger base weaker acid
than Cl- than HCl
Brønsted-Lowry theory of acids and bases
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weaker acid than H3O+ stronger base than H2O
conjugate pair
CH3COOH + H2O CH3COO- + H3O+
conjugate pair
weaker base stronger acid
than CH3COO- than CH3COOH
CH3COO – H …….. CH3COO-
Brønsted-Lowry theory of acids and bases
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Conjugate Acid-Base Pairs
Brønsted-Lowry theory of acids and bases
Acid Base
strongest HClO4 HClO4- weakest
H2SO4 HSO4-
HCl Cl-
HNO3 NO3-
H3O+ H2O
HSO4- SO4
2-
H3PO4 H2PO4-
HF F-
CH3COOH CH3COO-
H2CO3 HCO3-
H2PO4- HPO4
2-
NH4+ NH3
HCO3- CO3
2-
HPO42- PO4
3-
weakest H2O OH- strongest
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stronger acid than H2O weaker base than OH- conjugate pair
CH3COOH + OH- CH3COO- + H2O
conjugate pair
stronger base weaker acid
than CH3COO- than CH3COOH
weaker base than OH- stronger acid than H2O
conjugate pair
NH3 + H2O NH4+ + OH-
conjugate pair
weaker acid stronger base
than NH4+ than NH3
Brønsted-Lowry theory of acids and bases
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Brønsted-Lowry theory of acids and bases
The strengths of acids are measured by their tendencies to transfer protons to a base, usually water. For an acid with the general formula HA,
the equilibrium constant for ionization is obtained by equation:
HA + H2O H3O+ + A-
K = [H3O+] [A-]
[HA] [H2O]
Water concentration - [H2O], is so large compared to the concentrations of the ions formed in the equilibrium, that is why [H2O] is included in another equilibrium constant, the acid ionization constant, Ka.
Ka = K [H2O] = [H3O+] [A-]
[HA]
The larger the value of Ka, the stronger the acid.
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Brønsted-Lowry theory of acids and bases
To compare bases we use their tendency to accept protons from water. For the base with the general formula B, the base ionization reaction and the related equilibrium constant expression are:
B + H2O BH+ + OH-
K = [BH+] [OH-]
[B] [H2O]
As in the case of the acid ionization constant the [H2O] is included in the base ionization constant, Kb.
Kb = K [H2O] = [BH+] [OH-]
[B]
The larger the value of Kb, the stronger the base.
Ka and Kb are related in this way:
Ka . Kb = Kw Kw = 1.10 -14
Therefore, if you know one of them, you can calculate the other.
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Problem
Lactic acid is a monoprotic organic acid produced in metabolic reactions. Its Ka is 1.10-4. Calculate the Kb of its conjugate base.
Solution:
Lactic acid – HL,
than formation of its conjugate base can be represent with the
equation:
HL + H2O L- + H3O+
acid conjugate base
Ka.Kb = Kw
1.10-4.Kb = 1.10-14
Kb =1.10 -14 = 1.10 -10
1.10-4
Brønsted-Lowry theory of acids and bases
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Values of Ka and Kb are very small that is way it is convenient
to define pKa and pKb.
pKa = -logKa
pKb = -logKb
pKa + pKb = pKw = 14
pKa and pKb are used to determine if the acid or base is
strong or weak.
The larger the value of pKa (pKb) , the weaker the acid
(base).
Problem
For CH3COOH/CH3COO─ pKa + pKb = 14. Calculate the
pKb(CH3COO¯ ) if pKa(CH3COOH) = 4,76. Is the acetate ion
(CH3COO¯ ) a strong base?
Brønsted-Lowry theory of acids and bases
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Ionic product of water
Water is amphoteric and can act as an acid and a base with itself.
Water self ionizes at 250C to produce hydronium ions and hydroxide ions.
H2O + H2O H3O+ + OH-
acid base conjugate conjugate
acid base
One water molecule behaves as an acid, and another behaves as a base in the reaction.
The equilibrium constant expression is
K = [H3O+].[OH-]
[H2O]2
The concentration of water is so large that it remains virtually constant. It is included in an alternate equilibrium constant, ion product constant of water, Kw.
[H2O] = constant Kw = K.[H2O]2 = [H3O+].[OH-]
Kw = [H3O+].[OH-]
Kw – ion product constant of water
Kw = 1.10-14 at 25°C, it depends on temperature.
In all water solutions Kw = 1.10-14
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In pure water [H+] = [OH-] = 1.10-7M.
Depending on [H+] solutions are acidic, neutral and basic.
Concentration of H+ pH
Acidic solution [H+] > 1.10-7M < 7
Neutral solution [H+] = 1.10-7M = 7
Basic solution [H+] < 1.10-7M >7
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pH value
• pH is defined as the negative logarithm of the molarity of the hydrogen ions (hydronium ions).
pH = log [H+]
• for [H+] = 1.10-7M
pH = -log(1.10-7) = - (-7) = 7
This is the pH of neutral solution,
• pH of acidic solution is <7
and pH of basic solution is >7
• pOH = -log[OH-] pKw = - logKw
• as you know Kw = [H+].[OH-] = 1.10 -14 then
pKw = pH + pOH = 14
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Problem
Blood sample contains 4,5 . 10-8 M hydronium ions. What is
the pH?
Solution
[H3O+] = 4,5.10-8 M
pH = log [H3O+]
pH = log (4,5.10-8) = -log 4,5 + (-log 10-8)= -0,65 + 8 = 7,35
Problem
A 0,0050 M morphine solution has [OH-] = 8,8.10-5 M. What is
the pH of the solution?
Solution
[OH-] = 8,8.10-4 M
pOH = log [OH-]
pOH = log (8,8.10-4) = -log 8,8 + (-log 10-4)= -0,94 + 4 = 3,06
pH + pOH = 14
pH = 14 – pOH = 14 – 3,06 = 10,94
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pH
2 3 4 5 6 7 8 9 10 11 12
neutral @ 25oC
(H+) = (OH-)
distilled water
acidic
(H+) > (OH-)
basic or alkaline
(H+) < (OH-)
natural waters
pH = 6.5 - 8.5
normal rain (CO2)
pH = 5.3 – 5.7
acid rain (NOx, SOx)
pH of 4.2 - 4.4
0-14 scale for the chemists
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pH
1 2 3 4 5 6 8 9 10 11
The biological view in the human body
acidic basic/alkaline
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Methods to measure pH
Using indicators
Hind H+ + Ind ¯
colour A colour B
Paper tests like litmus paper and pH paper
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Using pH meter
• Tests the voltage of the electrolyte
• Converts the voltage to pH
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Buffer – a solution that prevents a drastic pH change when
either H+ or OH- is added to it.
Each buffer consists of a weak acid and its conjugate base
(salt of that acid)
or weak base and its conjugate acid (salt of that base)
dissolved in water.
Examples:
acetic acid and sodium acetate dissolved in water
CH3COOH / CH3COONa;
ammonia and ammonium chloride dissolved in water
NH3 / NH4Cl
Buffers
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Acetate buffer - CH3COOH / CH3COONa
determines pH of the buffer solution
CH3COOH CH3COO- + H+
CH3COONa CH3COO- + Na+
CH3COOH / CH3COO-
The buffer reacts to neutralize any strong acid or base that is
added to the solution.
could alter pH
HCl H+ + Cl-
H+ + Cl- + CH3COO- + Na+ CH3COOH + Na+ + Cl-
For short: H+ + CH3COO- CH3COOH
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NaOH Na+ + OH- could change pH
Na+ + OH- + CH3COOH CH3COO- + Na+ + H2O
Short equation: OH- + CH3COOH CH3COO- + H2O
! In general:
Added acid reacts with the base that present into a solution.
Added base reacts with the acid that present into a solution.
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Problem
How would the ammonium buffer (NH4OH / NH4Cl) neutralize
the effect of the addition of NaOH and HCl so that the pH value
remains the constant?
Solution:
NH4OH is the base
NH4+ from NH4Cl is its conjugate acid
So, added NaOH must react with the acid
NaOH + NH4Cl NH4OH + NaCl
For short: OH- + NH4+ NH4OH
HCl added reacts with the base
HCl + NH4OH NH4Cl + H2O
H+ + OH- H-OH
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Phosphate buffer:
H2PO4- / HPO4
2-
How this buffer reacts to neutralize addition of
HCl?
HPO42- + HCl H2PO4
- + Cl-
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pH of the buffer
What does the pH of buffer depend on? How to calculate
pH of buffer?
First we have to know [H+] in buffer solution. If the buffer
consists of weak acid HA and its salt MA obviously H+ are
produced from HA.
HA H+ + A-
Ka = [H+][A-] [H+] = Ka [HA]
[HA] [A-]
Thus pH=-log[H+] pH=-logKa + log [A-]
[HA]
pH = pKa + log [A-] (1)
[HA]
HA - acid
A- - conjugate base
pH depends on the strength of acids and the ratio [A-]
[HA]
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In the case of buffer consisting of a base and its salt the
equation used for pH calculating is:
pH = 14 – pKb + log [B]___ (2)
[HB+]
pH depends on strength of acids or bases
and on the ratio [A-]
[HA]
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The equations ( 1 ) and ( 2 )
are called Henderson – Hasselbalch’s equations.
pH = pKa + log [A-] (1)
[HA]
pH = 14 – pKb + log [B]___ (2)
[HB+]
They are used to calculate pH of buffers.
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Buffer capacity is defined as a moles of strong acid or
strong base which must be added into 1 L of buffer solution
in order to change its initial pH with 1 pH unit.
= number of moles of acid (base)
pH
Buffer capacity is higher when solutions are concentrated.
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Problem
Express the Henderson-Hasselbalch’s equation of
carbonate buffer (H2СO3/HСO3-). Calculate the pH value of
carbonate buffer with [H2СO3] = 0,02M and [HСO3-] = 0,02M,
pKa = 6,36.
Solution
pH = pKa + log [A-] pKa = 6,36
[HA] H2СO3 is the acid
HСO3- - is its conjugate base, so
pH = pKa + log [HCO3-]
[H2CO3]
pH = 6,36 + log 0,02 = 6,36
0,02
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H2CO3 / HCO3- - bicarbonate buffer; in the blood stream
H2PO4- / HPO4
2- - phosphate buffer; involved in kidney
functions
protein buffer: plasma protein and hemoglobin
How do these buffers work?
Bicarbonate buffer
1. HCO3- + H3O
+ H2CO3 + H2O
(aq) (aq) (aq) (l )
H2CO3 H2O + CO2
(aq) (aq) (g)
2. H2CO3 + OH- HCO3- + H2O
(aq) (aq) (aq) (l )
[H2CO3 ] / [HCO3- ] = 1:20
Buffers in the Body
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Phosphate buffer – H2PO4- / HPO4
2-
H2PO4- + OH- HPO4
2- + H2O
HPO42- + H3O
+ H2PO4- + H2O
[ H2PO4- ]/ [HPO4
2- ]=1:4
Hemoglobin buffer: HHb / Hb-O2-
HHb + O2 ↔ HbO2- + H+
oxyhemoglobin
HCO3- + H+ → H2CO3
H2CO3 → CO2 + H2O
Net oxygenation reaction is
HHb + O2 + HCO3- + H+ + H2CO3 ↔ HbO2
- +H+ + H2CO3 +CO2 + H2O
HHb + O2 + HCO3- ↔ HbO2
- + CO2 + H2O
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HHb-NH2 + CO2 → HHb-NH-COO- + H+
carbaminohemoglobin
H2O + CO2 ↔ H2CO3
H2CO3 → HCO3- + H+
Hb-O2- + H+ → HHb + O2
The net equation for removal of carbon dioxide is obtained by
summing the three equations.
Hb-O2- + CO2 + H2O ↔ HHb + O2 + HCO3
-
deoxygenation reaction
This reaction illustrates the effect of hemoglobin buffer.
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Acidosis – pH of blood lower than 7.28 for a period
of time
Alkalosis - pH of blood higher than 7.40 for a period
of time
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Terms
Acid
Base
Amphoteric
Conjugate pair (asid-base pair)
Acid ionization constant Ka
Base ionization constant Kb
Ion product constant of water Kw
pKa, pKb, pKw
pH, pOH, pH of biological fluids
Indicators
Buffers
How to calculate pH of the buffers
Buffer capacity
Buffers in the human body
Brønsted-Lowry theory of acids and bases
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