1) Acids cause substances like lemons and food to be sour and can damage materials like teeth and sculptures. Acids have positively charged hydrogen ions and turn litmus red. 2) Bases have negatively charged hydroxide ions, feel slippery, and turn litmus blue. Common bases include hand soaps and drain cleaners. 3) The Brønsted-Lowry concept defines acids as proton donors and bases as proton acceptors in reversible acid-base reactions. Both acids and bases can act as conjugates of each other by gaining or losing protons.

1. Course: Diploma
Subject: Applied science(Chemistry)
Unit: III

2. When we think of acids and bases we tend to think of science labs
and chemicals…but did you know
Acids cause:
•Lemons to be sour
•Acid rain to eat away at sculptures
•Framed paintings to be damaged
•Cavities in your teeth
•Food to digest in your stomach
•Ants and bees use it to sting
Acids and Bases

3. Acids and Bases

4. Acids and BasesAcids and Bases
3

5. Properties of Acids and BasesProperties of Acids and Bases
Acids
◦ turn blue litmus red
◦ taste sour
◦ Acids corrode metals
◦ positively charged hydrogen ions (H+
)
Bases
◦ turn red litmus blue
◦ taste bitter
◦ Negatively charged hydroxide ions (OH–
)
◦ Feel slippery
◦ Most hand soaps and drain cleaners are bases
◦ Strong bases are caustic

6. Acid-Base ConceptsAcid-Base Concepts
In the first part of this chapter we will
look at several concepts of acid-base
theory including:
– The Arrhenius concept
– The Bronsted Lowry concept
– The Lewis concept
This chapter expands on what you learned in
Chapter 3 about acids and bases.

7. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
According to the Arrhenius concept of acids
and bases, an acid is a substance that, when
dissolved in water, increases the
concentration of hydronium ion (H3O+
).
– Chemists often use the notation H+
(aq) for the
H3O+
(aq) ion, and call it the hydrogen ion.
– Remember, however, that the aqueous hydrogen
ion is actually chemically bonded to water, that is,
H3O+
.

8. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
According to the Arrhenius concept of acids and
bases, an acid is a substance that, when
dissolved in water, increases the
concentration of hydronium ion (H3O+
).
The H3O+
is shown
here hydrogen
bonded to three
water molecules.

9. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
A base, in the Arrhenius concept, is a
substance that, when dissolved in
water, increases the concentration
of hydroxide ion, OH-
(aq).

10. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
In the Arrhenius concept, a strong acid is a
substance that ionizes completely in aqueous
solution to give H3O+
(aq) and an anion.
– Other strong acids include HCl, HBr, HI, HNO3 ,
and H2SO4.
– An example is perchloric acid, HClO4.
)()()()( 4324 aqClOaqOHlOHaqHClO
−+
+→+

11. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
In the Arrhenius concept, a strong base
is a substance that ionizes completely
in aqueous solution to give OH-
(aq) and a
cation.
– Other strong bases include LiOH, KOH,
Ca(OH)2, Sr(OH)2, and Ba(OH)2.
– An example is sodium hydroxide, NaOH.
)aq(OH)aq(Na)s(NaOH OH2 −+
+→

12. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
Most other acids and bases that you encounter
are weak.They are not completely ionized
and exist in reversible reaction with the
corresponding ions.
nium hydroxide, NH4OH, is a weak base.
)aq(OH)aq(NH)aq( 4
−+
+
(aq)OHC(aq)OH 2323
−+
+
– An example is acetic acid, HC2H3O2.
)l(OH)aq(OHHC 2232 +

13. Arrhenius Concept of Acids andArrhenius Concept of Acids and
BasesBases
The Arrhenius concept is limited in that it
looks at acids and bases in aqueous
solutions only.
– In addition, it singles out the OH-
ion as the
source of base character, when other
species can play a similar role
– Broader definitions of acids and bases are
discussed in the next sections.

14. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
A base is the species accepting the proton in
a proton-transfer reaction.
– In any reversible acid-base reaction, both forward
and reverse reactions involve proton transfer.
• According to the Brønsted-Lowry concept, an
acid is the species donating the proton in a
proton-transfer reaction.

15. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Consider the reaction of NH3 and H20.
)aq(OH)aq(NH)l(OH)aq(NH 423
−+
++

16. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Consider the reaction of NH3 and H2O.
– In the forward reaction, NH3 accepts a proton
from H2O. Thus, NH3 is a base and H2O is an
acid.
)aq(OH)aq(NH)l(OH)aq(NH 423
−+
++
H+
base acid

17. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Consider the reaction of NH3 and H2O.
– In the reverse reaction, NH4
+
donates a
proton to OH-
. The NH4
+
ion is the acid and
OH-
is the base.
)aq(OH)aq(NH)l(OH)aq(NH 423
−+
++
H+
baseacid

18. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Consider the reaction of NH3 and H2O.
– A conjugate acid-base pair consists of two
species in an acid-base reaction, one acid and
one base, that differ by the loss or gain of a
proton.
)aq(OH)aq(NH)l(OH)aq(NH 423
−+
++
base acid
– The species NH4
+
and NH3 are a conjugate
acid-base pair.

19. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Consider the reaction of NH3 and H2O.
– The Brønsted-Lowry concept defines a species
as an acid or a base according to its function in
the proton-transfer reaction.
)aq(OH)aq(NH)l(OH)aq(NH 423
−+
++
base acid
– Here NH4
+
is the conjugate acid of NH3 and
NH3 is the conjugate base of NH4
+
.

20. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Some species can act as an acid or a base.
– For example, HCO3
-
acts as a proton donor (an acid) in
the presence of OH-
)l(OH)aq(CO)aq(OH)aq(HCO 2
2
33 +→+
−−−
–H+
– An amphoteric species is a species that can act
either as an acid or a base (it can gain or lose a
proton).

21. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
Some species can act as an acid or a base.
– An amphoteric species is a species that can act
either as an acid or a base (it can gain or lose a
proton).
– Alternatively, HCO3 can act as a proton acceptor
(a base) in the presence of HF.
)aq(F)aq(COH)aq(HF)aq(HCO 323
−−
+→+
H+

22. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
The amphoteric characteristic of water is
important in the acid-base properties of
aqueous solutions.
– Water reacts as an acid with the base NH3.
)aq(OH)aq(NH)l(OH)aq(NH 423
−+
+→+
H+

23. Brønsted-Lowry Concept of AcidsBrønsted-Lowry Concept of Acids
and Basesand Bases
The amphoteric characteristic of water is
important in the acid-base properties of
aqueous solutions.
– Water can also react as a base with the acid HF.
)aq(OH)aq(F)l(OH)aq(HF 32
+−
+→+
H+

24. Classes of Bronsted Acids & BasesClasses of Bronsted Acids & Bases
 1. Monoprotic acids: one protons
 E.g., HF, CH3COOH
 2. Polyprotic acids: two or more protons
 E.g., H2S, Oxalic acid
 Bases: 1. Monoprotoc Bases: HS-
, H2O
 2. Polyprotic bases: SO4
-2
, PO4
-3

25. Lewis Concept of Acids and BasesLewis Concept of Acids and Bases
acid : electron pair acceptor
 base : electron pair donor.
Creates covalent bond : complex
–

26. Lewis Concept of Acids and BasesLewis Concept of Acids and Bases
The reaction of boron trifluoride with
ammonia is an example.
– Boron trifluoride accepts the electron pair, so it is a
Lewis acid. Ammonia donates the electron pair,
so it is the Lewis base.
+ N
H
H
H:
::
: B
F
F
F
: :
::
::
::
: B
F
F
F
: :
::
::
N
H
H
H

27. Examples of Lewis reactionsExamples of Lewis reactions
1. Between H+
and NH3
2. H+
and OH-
3. H2O and CH3
4. BF3 and NH3
5. Hydration of Al+3
6. H-
and BH3

28. Self-ionization of WaterSelf-ionization of Water
Self-ionization is a reaction in which two like
molecules react to give ions.
– In the case of water, the following equilibrium is
established.
)aq(OH)aq(OH)l(OH)l(OH 322
−+
++
– The equilibrium-constant expression for this
system is:
2
2
3
c
]OH[
]OH][OH[
K
−+
=

29. Self-ionization of WaterSelf-ionization of Water
– The concentration of ions is extremely
small, so the concentration of H2O remains
essentially constant. This gives:
]OH][OH[K]OH[ 3c
2
2
−+
=
constant

30. Self-ionization of WaterSelf-ionization of Water
– We call the equilibrium value for the ion product
[H3O+
][OH-
] the ion-product constant for water,
which is written Kw.
]OH][OH[K 3w
−+
=
– At 25 o
C, the value of Kw is 1.0 x 10-14
.
– Like any equilibrium constant, Kw varies with
temperature.

31. Self-ionization of WaterSelf-ionization of Water
Self-ionization is a reaction in which two like
molecules react to give ions.
– Because we often write H3O+
as H+
, the ion-
product constant expression for water can be
written:
]OH][H[Kw
−+
=
– Using Kw you can calculate the concentrations of
H+
and OH-
ions in pure water.

32. Self-ionization of WaterSelf-ionization of Water
These ions are produced in equal numbers in
pure water, so if we let x = [H+
] = [OH-
]
– Thus, the concentrations of H+
and OH-
in pure
water are both 1.0 x 10-7
M.
– If you add acid or base to water they are no longer
equal but the Kw expression still holds.
C25at)x)(x(100.1 o14
=× −
714
100.1100.1x −−
×=×=

33. Understanding the pH ScaleUnderstanding the pH Scale
o pH stands for (presence of Hydrogen)
o Numbered from 0 to 14.
o The lower the pH number – the higher Acid
o That means more Hydrogen Ions (H+)
o The higher the pH - the higher the Base
o That means less Hydrogen Ions (H+)
2

34. The pH of a SolutionThe pH of a Solution
Although you can quantitatively describe
the acidity of a solution by its [H+
], it is
often more convenient to give acidity in
terms of pH.
– The pH of a solution is defined as the negative
logarithm of the molar hydrogen-ion concentration.
]Hlog[pH +
−=

35. The pH of a SolutionThe pH of a Solution
For a solution in which the hydrogen-ion
concentration is 1.0 x 10-3
, the pH is:
– Note that the number of decimal places in
the pH equals the number of significant
figures in the hydrogen-ion concentration.
00.3)100.1log( 3
=×−= −
pH

36. The pH of a SolutionThe pH of a Solution
 In a neutral solution, whose hydrogen-ion
concentration is 1.0 x 10-7
, the pH = 7.00.
• For acidic solutions, the hydrogen-ion
concentration is greater than 1.0 x 10-7
, so the
pH is less than 7.00.
• Similarly, a basic solution has a pH greater
than 7.00.
• Figure 16.6 shows a diagram of the pH scale
and the pH values of some common solutions.

37. The pH ScaleThe pH Scale

38. A Problem to ConsiderA Problem to Consider
A sample of orange juice has a hydrogen-ion
concentration of 2.9 x 10-4
M.What is the pH?
]Hlog[pH +
−=
)109.2log(pH 4−
×−=
54.3pH =

39. A Problem to ConsiderA Problem to Consider
The pH of human arterial blood is 7.40.What
is the hydrogen-ion concentration?
)pHlog(anti]H[ −=+
)40.7log(anti]H[ −=+
M100.410]H[ 840.7 −−+
×==

40. The pH of a SolutionThe pH of a Solution
A measurement of the hydroxide ion
concentration, similar to pH, is the pOH.
– The pOH of a solution is defined as the
negative logarithm of the molar hydroxide-
ion concentration.
]OHlog[pOH −
−=

41. The pH of a SolutionThe pH of a Solution
A measurement of the hydroxide ion
concentration, similar to pH, is the pOH.
– Then because Kw = [H+
][OH-
] = 1.0 x 10-14
at 25 o
C, you can show that
00.14pOHpH =+

42. A Problem to ConsiderA Problem to Consider
An ammonia solution has a hydroxide-ion
concentration of 1.9 x 10-3
M.What is the
pH of the solution?
You first calculate the pOH:
72.2)109.1log(pOH 3
=×−= −
Then the pH is:
28.1172.200.14pH =−=

43. Buffer SolutionsBuffer Solutions
 DEFINITION: A buffer solution contains a weak acid mixed
with its conjugate base (or weak base and conjugate acid)
 Buffers resist changes in pH when a small amount of a
strong acid or base is added to it.
HA ∏ H+
+ A-

44.  If a small amount of a strong acid (H+
) is added eqm
shifts to the left as [H+
] increases so system adjusts to
increase [HA] and reduce [H+
] again.
HA ∏ H+
+ A-

45.  A small amount of a strong base will react with H+
to
form H2O and eqm
will shift to the right to increase [H+
]
again.
HA ∏ H+
+ A-

46. Buffering agent pKa useful pH range
Citric acid 3.13, 4.76, 6.40 2.1 - 7.4
Acetic acid 4.8 3.8 - 5.8
KH2PO4, 7.2 6.2 - 8.2
CHES 9.3 8.3–10.3
Borate 9.24 8.25 - 10.25

47. Acid Base IndicatorAcid Base Indicator
Acid Base titration: add base from burette in to
an acid.
Amount of base and acid are equal : equivalence
point or end point.
End point: shown by color change of indicator.
‘An Acid base indicator : organic dye that
signals the end point by a visual change in
color.’

48. Phenolphthalein – pink in base
Phenolphthalein – colorless in acid
Methyl orange-red in acid
Methyl orange – yellow in base.

49. ReferencesReferences
1.Essentials of Physical Chemistry by Bahl
and tuli
2.https://sites.google.com/a/wyckoffschools.
org/ems-biochemistry/acids-and-bases
3. https://www.mhhe.com