Acid and bavyjc jo gv hiking ijjses.pptx

Acids, Bases are defined by Four main theories,
1.Traditional theory / concept
2.Arrhenius theory
3.Bronsted and Lowry theory
4.Lewistheory

1. Traditional theory / concept-Acid:
• Acids: are the substances
• Which converts blue litmus paper to red
• Having the PH<7
• Sour taste
• React with bases to form salts and water
• Eg:-Hydrochloric acid(HCl)

1. Traditional theory / concept-Base:
Base: are the substances
• Which converts red litmus paper to blue
• Having the PH>7
• Bitter taste
• React with Acids to form salts and water
• Eg: Sodium Hydroxide (NaOH)

2. Arrhenius theory:
The Swedish chemist Svante Arrhenius proposed the first definition of acids
and bases
• According to Arrhenius Concept, Acid are substance which are capable of
providing
• Hydrogen ions (H+, proton) when dissolved in water and bases are
substances which are capable
• of providing hydroxide ions (OH-, hydroxyl ions) in aqueous solution.
• For example, Hydrochloric acid in the water, HCl undergoes dissociation
reaction to
• produce H+ ion and Cl– ion, as explained below. The concentration of the
H+ ions is increased by
• forming hydronium ion.

• HCl (aq) → H+(aq) + Cl– (aq)
• HCl (aq) + H2O(l) → H3O+(aq) + Cl– (aq)
• Other examples of Arrhenius acids are listed below
• NHO3(aq) + H2O(l) → H3O+ (aq) + NO3–
• In this reaction, nitric acid dissolves in aqueous water to give hydrogen
and nitrate ions.
• Another example of Arrhenius Base is
NaOH + H2O → Na+ + OH- + H2O

Neutralization reaction
Acid react Base
Salt &Water
Eg: Hydrochloric acid react sodium hydroxide
Sodium chloride (Salt) & water
NaOH + HCl------------------- NaCl + H2O

• “Neutralization as the process in which hydrogen ion and hydroxyl ion
combine to form unionized molecule or water”
• NaOH + HCl------------------- NaCl + H2O
• HCl (aq) → H+(aq) + Cl– (aq)
• NaOH + → Na+ + OH-
• H+ + OH- → H2O

Application
1. Non-metallic Oxides are arrhenius acidic in nature
eg CO2, SO2, SO3, N2O5, P4O10
• CO2 + H2O ------------------- H2CO3------ H+ + HCO3
• SO3 + H2O ---------------------H2SO4 -------H+ + HSO4-
2. metallic Oxides are arrhenius bases
eg CaO, Na2O

Limitations of Arrhenius theory:
• The Arrhenius theory is applicable only in aqueous solution; for example, according to the theory, HCl
is an acid in the aqueous solution but not in benzene, even though it donates H+ ion to the benzene.
Also, under Arrhenius’s definition, the solution of sodium amide in liquid
• ammonia is not alkaline, even though amide ion deprotonates the ammonia.
• Basisity of Ammonia (No OH-ion)is not explained
• Acidity of BF3,AlCl3 (No H+ ion)is not explained
• Acidity of oxides of P block element (CO2) is not explained
• Basicity of oxides of S block element (Na2o)is not explained
• Neutralization with out absence of solvent is not explained

3.Bronsted -Lowry concept:
Bronsted -Lowry concept: In 1923 the Danish chemist Johannes Nicolaus Bronsted and the
English chemist Thomas Martin Lowry, proposed the theory.
• According to Bronsted-Lowry theory, An acid is any substance (molecular or ionic) that can
donate a proton to any other substance (molecular or ionic) and a base is any substance that
can accept a proton from anyother substance.
• HCl + H2O H3O+
+ Cl
• In the above example what is the Bronsted acid? What is the Bronsted base?
In reality, the reaction of HCl with H2O is an equilibrium and occurs in both directions,
although in this case the equilibrium lies far to the right.
HCl + H2O H3O+
+ Cl-

• For the reverse reaction Cl- behaves as a Bronsted base and H3O+
behaves as a Bronsted acid.
• The Cl- is called the conjugate base of HCl. Bronsted acids and bases
always exist as conjugate acid-base pairs. Their formulas differ by only
one proton.
• Acid Base conjugate acid conjugate base
HCl + NH3 NH4
+
+ Cl-

1.Amphoteric: a species that can act as an acid or a base water is an
example of an amphoteric species.
2.Conjugatebase: species that remains after an acid donates its H+.
3.Conjugateacid: species that forms after a base accepts a H+
• 14

• Every Arrhenius Acid is Bronsted Acid
• Every Arrhenius Base is not Bronsted Base e.g., NaOH is Arrhenius base
because it gives
• OH- ion in aqueous solution but not a Bronsted base because it cannot accept
proton.
• Limitations of Bronsted Lowry Concept:
• The protonic definition cannot be used to explain the reactions occurring in
non-protonic solvents such as COCl2, SO2, N2O4, etc.
• Substances like BF3, AlCl3 etc, do not have any hydrogen and hence cannot
give a proton but are known to behave as acids

Lewis Theory
• In 1923 of scientist G.N. Lewis proposed the theory in terms of chemical
structure.
• Lewis Acids:
• Lewis acids accept an electron pair. Lewis Acids are Electrophilic meaning
that they are electron attracting.
• Various species can act as Lewis acids. All cations are Lewis acids since
they are able to accept electrons. (e.g., Cu2+
,Fe2+
,Fe3+
)
• Lewis acids- H+
, NH4+
, Na+
, K+
, Cu2+,
Al3+,
etc.

• Lewis Bases
• Lewis Bases donate an electron pair. Lewis Bases are Nucleophilic meaning that they “attack” a
positive charge with their lone pair. An atom, ion, or molecule with a lonepair of electrons can thus
be a Lewis base.
• Lewis base- NH3, H2O, OH-
, Cl-
, CN-
, S2-
, etc.
• (Lewis base) (Lewis acid)
• Boron trifluoride accepts the electron pair, so it is a Lewis acid. Ammonia makes available (donate)
the electron pair, so it is the Lewis base.

Importance of acids and bases in pharmacy
• Acids, bases and their reaction play vital role in pharmacy practice. Some of the
main application of the these are as follows:
• Acid-base neutralization reaction finds use in preparative procedures for the
preparation of suitable salt, and for conversion of certain salts into more suitable
forms.
• Acid-base is used in analytical procedure which is involving acid-base titrations.
• Acids and bases find use as therapeutic agents in the control of and adjustment of
pH of the GI tract, body fluids and urine.

Buffers:
• A buffer is a solution that can resist pH change upon the addition of an
acidic or basic components.
• It is able to neutralize small amounts of added acid or base, thus
maintaining the pH of the solution relatively stable.
• •This is important for processes and/or reactions which require specific
and stable pH ranges.

Strong Acids vs. Weak Acids
Strong acids are assumed to be 100%
ionized in solution (good proton donors).
Weak acids are usually less than 5%
ionized in solution (poor proton donors).
HCl H2SO4 HNO3
H3PO4 HC2H3O2 Organic acids

• Which of the following "molecular" pictures best
represents a concentrated solution of the weak acid
HA?
Concept Test
A B

pH Scale
0
7
INCREASING
ACIDITY NEUTRAL
INCREASING
BASICITY
14
• pH
• a measure of the concentration of H3O+
ions in solution
• measured with a pH meter or an indicator with a wide color
range

Bases have
a pH
greater
than 7

pH Scale
pH of Common Substances

• Buffers: Buffers are defined as a compound or a mixture of compounds
that resists the pH upon the addition of small quantities of acid or alkali.
Buffer have definite pH value.
• The pH will not change after keeping it for a long period of time. The pH
value altered negligibly by the addition of small quantities of acid or base.
• Buffer action: The resistance to a change in pH is known as buffer action.
So buffers can be added to show buffer action.
• Buffer capacity: The amount of acid/base required to produce a unit
change in pH in a solution is called buffer capacity.

Buffers system:
• A buffer system can be made of a weak acid and its salt or a weak base
and its salt.
• A classic example of a weak acid based buffer is acetic acid
(CH3COOH) and sodium acetate(CH3COONa).
• A common weak base buffer is made of ammonia (NH3) and ammonium
chloride (NH4Cl)

Types of Buffers :
Generally buffers are of two types:
1. Acidic buffers
2. Basic buffers
There are some other buffer system:
3. Two salts acts as acid-base pair. Ex- Potassium hydrogen phosphate and potassium
dihydrogen phosphate.
4. Amphoteric electrolyte. Ex- Solution of glycine.
5. Solution of strong acid and solution of strong base. Ex- Strong HCl with KCl.

1. Acidic Buffers:
• An acidic buffer is a combination of weak acid and its salt with a strong
base. i.e. Weak acid & salt with strong base (conjugate base).
• EXAMPLES:
CH3COOH / CH3COONa
H2CO3 / NaHCO3
H3PO4 / NaH2PO4
HCOOH / HCOONa

2. Basic Buffers:
• A basic buffer is a combination of weak base and its salt with a strong
acid.i.e. Weak base & salt with strong acid (conjugate acid).
• EXAMPLES:
NH4OH / NH4Cl
NH3 / NH4Cl
NH3 / (NH4)2CO3

Mechanism of Buffer action:
Mechanism of Buffer action:
• The resistance of a buffer solution to a change in pH is known as buffer action.
• In a buffer solution, the components interact with each other and produce a
dynamic equilibrium.
• When a small quantity of acid or base is added, the dynamic equilibrium shifts
and nullifies the effect of the addition.

 Mechanism of Action of acidic buffers:
• Consider a buffer system of CH3COOH (Weak electrolyte) and CH3COONa (Strong
electrolyte). There will be a large concentration of Na+ ions, CH3COONa – ions, and
undissociated CH3COOH molecules.
When an acid is added:
• If a strong acid (HCl) is added in CH3COOH / CH3COONa buffer, the changes that will occur
may be represented as:
CH3COONa Na + COO H + Cl
CH3COOH
• The hydrogen ions yielded by the HCl are quickly removed as unionized acetic acid, and the
hydrogen ion concentration is therefore only slightly affected (because acetic acid produced is
very weak as compared to HCl added).
-
- +
+

When a base is added:
• If a strong base (NaOH) is added in NH4OH / NH4Cl buffer, the changes that
will occur may be represented as:
CH3COOH CH3COO + H OH + Na NaOH
H2O
• The hydroxyl ions yielded by the NaOH are therefore removed as water. The
supply of hydrogen ions needed for this purpose being constantly provided
by the dissociation of acetic acid..
+
+
-
-

 Buffer equation-Henderson-
Hasselbalch equation:
• The buffer equation is also known as Henderson-Hasselbalch equation,
with the help of this equation it is possible to calculate the pH of a buffer
solution of known concentration or to makebuffer solution of known pH.
• Two separate equations are obtained for each type of buffer, acidic and
basic
• pH of acidic buffer: The hydrogen ion concentration obtained from the
dissociation of weak acid HA is given by equation,

HA H+ + A-
𝐾𝑎
Ka = equilibrium constant
𝐾𝑎
Taking logarithms of both sides of the equation & multiplying throughout by -1
gives
-log -log -log
𝐾𝑎
pH pKa + log
𝑝𝐻 = + log
𝑝𝐾𝑎

pH of an alkaline buffer: The ionization of a weak base BOH is given by,
BOH B+ + OH-
𝐾b
Kb = equilibrium constant
𝐾b
Taking logarithms of both sides of the equation & multiplying throughout by -1
gives
-log -log -log
𝐾𝑎
pH pKa + log
𝑝𝐻 = + log
𝑝𝐾𝑎

Buffer capacity:
• Buffer capacity may also be defined as “The maximum amount of either
strong acid or strong base that can be added before a significant change in
the pH will occur”.
• The maximum amount of strong acid that can be added is equal to the
amount of conjugate base present in the buffer whereas the maximum
amount of base that can be added is equal to the amount of weak acid present
in the buffer.

• Buffer capacity is depend on the factors:
1. The concentration of the acid or base component of the buffer (Direct relation)
2. The pH of the buffer
• Buffer can act best at pH = pKa and buffering range is pH = pKa +1
• Or It may be defined as the moles of strong acid or strong base required to
change the pH of 1000 ml of buffer solution by one unit.
• The magnitude of the resistance of a buffer to pH changes is referred to as the
buffer capacity, β.
• Where, ΔB is the small increment in gram equivalents (g Eq)/liter of strong base
added to the buffer solution and ΔpH: change in a pH

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