Chemistry Preliminary Course
2015
Acids and Bases
17/09/15
Outline
What are acids and bases?
Can we provide a general
definition of acid and base?
How can we quantify acidity
and basicity?
Can we classify acid and base
strength?
pH concept and pH scale.
Acid/base reactions:
neutralization
How can we monitor an acid/base
reaction in real time?
What is an Acid?
Some Common Acids
Acetic Acid or Ethanoic Acid
Carbonic Acid
Citric Acid
Phosphoric Acid
Some Common Acids
Aspirin
Formic Acid
Sulfuric Acid
Some Common Bases
Ammonia
Sodium Bicarbonate
Oven Cleaners and Drain Unblockers
Caffeine
General Acid/Base Properties
Acids
Have a sharp or sour taste
React with metals to produce hydrogen gas
React with (bi)carbonates to produce CO2 gas
-This results in weathering of buildings, etc
Bases
Have a bitter taste
React with acids to make salts
React with oils to make soap
-So they feel slippery on your hands
The Search for an Acid/Base Theory
Science as a whole tries to DESCRIBE
the world by constructing MODELS.
Models that do not cover everything
are not useless – they are helpful tools
to describe their own situation
As better or more general models are
developed, they replace the old
models
The Classical Acid/Base Model
The first theory of acids and bases was proposed by Arrhenius
His theory was based on how they act in water (or in aqueous solutions)
Under this model…:
An ACID is a H-containing substance that dissociates in
water to produce HYDRONIUM IONS, H3O+
HA + H2O
A- + H3O+
A BASE is an OH-containing substance that dissociates in
water to produce HYDROXIDE IONS, OHBOH
B+ + OH-
The Classical Acid/Base Model
Under this definition, NEUTRALISATION is the reaction of
hydronium (H3O+) and hydroxide (OH-) to form water
H3O+ + OH-
2H2O
The counterparts to the neutralisation reaction generally form a salt
HCl + H2O
NaOH
Cl- + H3O+
Na+
+
OH-
2H2O + NaCl
This is a fairly good model – but doesn’t account for everything
A Note on “Hydronium”
The H3O+ hydronium ion is often
represented simply as “H+”
This is simpler and easier to write, but “H+” is simply a proton –
and an isolated proton simply cannot exist by itself in solution
However, “H3O+” is also a simplification – acidified water is
EXTREMELY complicated, with large and dynamic conglomerates
of water molecules really stabilising the extra protons.
On balance, it’s probably best to write “H3O+” (to demonstrate
you know free protons don’t exist), but don’t be confused if you
see “H+”
The Problems with Arrhenius
For example, ammonia (NH3) has all the properties of a base, but
doesn’t contain any OH groups, so doesn’t fit Arrhenius’s definition
Similarly, Boric Acid (BO3H3) produces hydronium ions, but by taking
on an OH, not by losing a H, so it doesn’t fit Arrhenius’s definition
BO3H3 + H2O
BO4H4- + H3O+
A More General Definition
Bronsted and Lowry replaced the
Arrhenius definition with a more
general definition centred on
protons.
An ACID is a PROTON DONOR
A BASE is a PROTON ACCEPTOR
An acid-base reaction is therefore
the transfer of a proton. Acids and
bases must act concurrently.
A More General Definition
Does this match Arrhenius’s definition?
If I put an acid in water, it can donate a proton to form hydronium:
HA + H2O
A- + H3O+
If I put an base in water, it can accept a proton to form hydroxide:
B + H2O
BH+ + OH-
So Bronsted and Lowry repeat Arrhenius’s observations, but with more general rules:
- Water does not need to be present
- More things fit the definitions of acid and base (eg. NH3 and BO3H3)
A More General Definition
How does water fit in to these definitions of acid/base?
With acid:
HA + H2O
A- + H3O+
So water is
acting as a
base
Accepted a proton
And with base:
B + H2O
BH+ + OH-
So water is
acting as
an acid
Donated a proton
Water is known as an AMPHOTERIC or AMBIPROTIC substance
A More General Definition
Acid-base processes are DYNAMIC. They are not one way streets – the products can turn
back into the reactants. This is known as an EQUILIBRIUM PROCESS, denoted by
HA + H2O
A- + H3O+
Therefore:
After an acid donates its proton, it is known as a CONJUGATE BASE
After an base accepts its proton, it is known as a CONJUGATE ACID
CONJUGATE
ACID
ACID
HA + H2O
BASE
A- + H3O+
CONJUGATE
BASE
Equilibrium is established very quickly in solution
A More General Definition
CONJUGATE PAIR
CONJUGATE
ACID
ACID
HA + H2O
BASE
A- + H3O+
CONJUGATE
BASE
CONJUGATE PAIR
Reaction 1
HF
+
H2O
F-
+
H3O+
Reaction 2
HCOOH +
CN-
HCOO-
+
HCN
Reaction 3
NH4+
+
CO32-
NH3
+
HCO3-
Reaction 4
H2PO4-
+
OH-
HPO42-
+
H2O
Reaction 5
H2SO4
+
N2H5+
HSO4-
+
N2H62+
Reaction 6
HPO42-
+
SO32-
PO43-
+
HSO3-
Strong and Weak Acids
With some acids, you need to add
more to produce the same effect
The STRENGTH of an acid is the
DEGREE OF IONISATION
So a STRONG ACID will exist mostly as
A- and H3O+
A WEAK ACID will exist mostly as HA
Weak acid is
mostly this
HA + H2O
A-
+
H3O+
Strong acid is
mostly this
Strength and concentration combine to produce effects.
A dilute solution of a strong acid could have the same concentration of H3O+ as a
concentrated solution of a weak acid, and therefore react similarly
Strong and Weak Acids
Strong acids
TEND to be
mineral acids.
These FULLY
DISSOCIATE in
solution.
Weak acids
TEND to be
organic acids.
These only
PARTLY
DISSOCIATE in
solution.
HCl
(Hydrochloric Acid)
H2SO4 (Sulfuric Acid)
HNO3 (Nitric Acid)
But concentration is also important!!
Formic Acid
Acetic Acid
Citric Acid
Strong and Weak Acids
CONJUGATE PAIR
CONJUGATE
ACID
ACID
HA + H2O
BASE
A- + H3O+
CONJUGATE
BASE
CONJUGATE PAIR
A strong acid has a weak conjugate base
A weak acid has a strong conjugate base
Strong and Weak Acids
“Electrolytes” are just charged (ionic)
species – like salt.
Strong acids or bases make
good electrolytes
Weak acids or bases make
bad electrolytes
This has an impact on
electrochemistry – hence
sulfuric acid (strong acid) is
present in car batteries
Quantifying Acid Strength
Need a way to compare the
strength of different acids and
bases
We do this by looking at the
equilibrium concentrations in
solution.
A strong acid will be mostly Aand H3O+, while a weak acid will
be mostly HA
Equilibrium Interlude
DYNAMIC EQUILIBRIUM occurs when a reaction is reversible:
A+B
C+D
These reactions are said to be at equilibrium if the RATES OF REACTION ARE EQUAL
Rate of flow in….
…equals rate of flow out.
Therefore, the CONCENTRATION DOES NOT CHANGE with time
Equilibrium Interlude
The location of the balance point is influenced by many factors (which you’ll
learn about in your chemistry courses), but for now, you just need to know
that it is quantified by an EQUILIBRIUM CONSTANT, K:
A+B
K=
[C][D]
[A][B]
C+D
Where “[x]” represents
“concentration of x” in
units of mol/L
The products go over the reactants
If K is large…
…[C] and [D] are large, or [A] and [B] are small - so solution is mostly C + D
If K is small…
…[C] and [D] are small, or [A] and [B] are large - so solution is mostly A + B
Quantifying Acid Strength
We can use these generic equilibrium ideas to quantify acid or base strength.
For the generic acid dissociation:
A- + H3O+
HA + H2O
K for this process is defined as:
K=
[A-][H3O+]
[H2O][HA]
Where “[x]” represents
“concentration of x” in
units of mol/L
We get rid of the water element (because it doesn’t tell us much about dissociation)
to give an ACID DISSOCIATION CONSTANT, Ka:
Ka=
[H2O]K=
[A-][H3O+]
[HA]
Quantifying Acid Strength
Similarly, a base dissociation constant Kb can be defined as:
B + H2O
Kb=
[H2O]K=
BH+ + OH-
[BH+][OH-]
[B]
A LARGE value of Kb indicates a STRONG BASE
– [BH+] and [OH-] are large, [B] is small
A LARGE value of Ka indicates a STRONG ACID –
[A-] and [H3O+] are large, [HA] is small
Quantifying Acid Strength
Ka and Kb values can vary MASSIVELY – so it’s more beneficial to
use a log scale:
pKa = - log10Ka
and
pKb = - log10Kb
The logarithm is a mathematical operation
that is the inverse of exponentiation.
If x = bn then logb(x)=n where b is the base.
If 10x = y then log10y = x,
Example: 102=10x10=100, then log10(100)=2.
A change of one unit in pKa indicates a tenfold increase in Ka
Quantifying Acid Strength
As Ka goes up (IE. as the acid gets stronger), pKa goes DOWN.
As Kb goes up (IE. as the base gets stronger), pKb goes DOWN.
Acid Name (Formula)
KA at 298 K
pKA
Hydrogen sulfate ion (HSO4-)
1.02 x 10-2
1.991
Nitrous acid (HNO2)
7.1 x 10-4
KA 
3.15
Acetic acid (CH3COOH)
1.8 x 10-5
pK A 
4.74
Hypobromous acid (HBrO)
2.3 x 10-9
8.64
Phenol (C6H5OH)
1.0 x 10-10
10.00
Quantifying Acid Strength
If pKa is less than about 2, its considered a strong acid
The Behaviour of Water
Earlier we talked about how water can act as both ACID AND BASE (amphoteric):
This behaviour can also happen in pure water:
ACID
H2O + H2O
BASE
CONJUGATE
BASE
H3O+ + OHCONJUGATE
ACID
This is known as AUTOIONISATION
The Behaviour of Water
We quantify this behaviour in a similar way, using equilibrium constants:
H3O+ + OH-
H2O + H2O
The equilibrium constant for this reaction is :
[H O+][OH-]
K=
3
[H2O]
Again, we get rid of the [H2O] part, because it’s extremely large and effectively constant:
Kw = K[H2O] = [H3O+][OH-]
“the ionic product of water”
At 250C
Kw = [H3O+][OH-] = 1.0 x 10-14
The Behaviour of Water
At 250C
Kw = [H3O+][OH-] = 1.0 x 10-14
[H3O+] = [OH-]
[H3O+] > [OH-]
Solution Is
neutral
acidic
[H3O+] < [OH-]
basic
If we take logs, like we did earlier:
Kw = [H3O+][OH-] = 1.0 x 10-14
pKw = -log10([H3O+][OH-]) = -log10(1.0 x 10-14) = 14
pKw = -log10[H3O+]– log10[OH-] = 14
In general,
log  a.b   log(a)  log(b)
We define –log10[H3O+] as pH and –log10[OH-] as pH and pOH respectively, giving:
pKw = pH + pOH= 14
The pH concept
On the last slide we defined pH = –log10[H3O+]
If pH = pOH (ie. the solution is NEUTRAL):
pKw = pH + pOH = 2pH = 14
pH = 7
If pH < pOH (ie. pH < 7), the solution is ACIDIC
If pH > pOH (ie. pH > 7), the solution is BASIC
The pH concept
pH<7
pH>7
pH
[H3O+] 100 M
10-14 M
(1.0 M)
pH=7
A pH change of 1 unit implies a 10 fold change in [H3O+]
The pH concept
Some typical pH values:
Remember, concentration and acid strength are both interlinked with pH
pH Indicators
Indicators are compounds (generally weak acids) that change
colour based on pH (generally from donating protons):
HIn (aq)  H 2O  H 3O  (aq)  In 
For Litmus:
Blue
Red
pH Indicators
Universal indicator is a mixture of indicators to give a full range of pH values
A More Accurate Way to Measure
pH meters are a much more accurate and
sophisticated way to measure pH. They consist of
a probe (voltmeter) which measures electrical
potential across a membrane. This potential is
proportional to [H3O+], so pH can be read off.
pKa vs pH
Ka and pKa are measures of an acid’s tendency to dissociate:
Ka=
[A-][H3O+]
[HA]
pKa = - log10Ka
Strong acid has a
large Ka, and a
small pKa
pH is a measure of the amount of acid ([H3O+]) in aqueous solution:
pH = –log10[H3O+]
pH goes down as
[H3O+] goes up
Measuring pH
TITRATIONS are used to determine the concentration of an unknown substance.
For this (Arrhenius) base reaction:
HA  MOH  MA  H 2O
Base (known conc.)
Base is added until the reaction reaches
the EQUIVALENCE POINT – ie. The point at
which the reaction is complete
Ie. The amount of acid = amount of base
Acid (unknown conc.)
An indicator which changes colour at this
pH is added to see if the reaction is
complete
Measuring pH
For Strong Acids/bases:
NaOH (aq) + HCl (aq)
H2O (l) + NaCl (aq)
0.10 M NaOH added to 25 mL of 0.10 M HCl
Measuring pH
For weak acids/bases:
CH3COOH (aq) + NaOH (aq)
CH3COONa (aq) + H2O (l)
Some Recommended Reading

Silberberg, Chemistry, 4th edition.
◦ Chapter 18.
 Acid/base equilibria. pp.766-813.
◦ Chapter 19.
 Ionic equilibria in aqueous systems. pp.814-862.

Kotz, Treichel and Weaver, 7th edition.

Burrows et al. Chemistry3 (OUP), 2009.Ch.6,
pp.263-300.
◦ Chapter 17&18, pp.760-859.
Lecture notes available after course on School of
Chemistry website located at:
http://www.tcd.ie/Chemistry/outreach/prelim/

Some Recommended (Online) Reading





http://www.shodor.org/unchem/basic/ab/
http://chemistry.about.com/od/acidsbases/
http://www.chem.neu.edu/Courses/1221PAM/acidb
ase/index.htm
http://dbhs.wvusd.k12.ca.us/webdocs/AcidBase/Aci
dBase.html
http://www.sparknotes.com/chemistry/acidsbases/f
undamentals/section1.html