Unit-2 : pH and Buffers, Calculation related
to enzyme activity
2.1 Concept of pH and the pH Scale
1. Definition of pH
The pH scale is a measure of how acidic or basic a solution is. The term "pH" stands for
"potential of hydrogen" or "power of hydrogen." It quantifies the concentration of hydrogen ions
(H⁺) in a solution. The pH scale ranges from 0 to 14:
pH < 7: Acidic solutions, with higher concentrations of hydrogen ions.
pH = 7: Neutral solutions, like pure water.
pH > 7: Basic (alkaline) solutions, with lower concentrations of hydrogen ions and
higher concentrations of hydroxide ions (OH⁻).
2. Historical Background
The concept of pH was introduced by Danish chemist Søren Sørensen in 1909. Sørensen was
working at the Carlsberg Laboratory in Copenhagen, and he sought a way to express the acidity
and alkalinity of solutions quantitatively.
3. Importance of pH
pH is crucial in various fields:
Chemistry: pH affects reaction rates, equilibria, and solubility.
Biology: Enzyme activity, cellular processes, and metabolic functions are pH-dependent.
Environmental Science: pH impacts ecosystems, water quality, and soil health.
Medicine: Body fluids' pH, such as blood and urine, are vital for diagnosing and
monitoring health conditions.
Fundamentals of pH
1. The pH Formula
The pH of a solution is calculated using the formula: pH=−log[H+]\text{pH} = -
\log[H^+]pH=−log[H+] where [H+][H^+][H+] is the concentration of hydrogen ions in moles
per liter (M).
2. The Relationship between pH and pOH
, The pH and pOH of a solution are related by the following equation: pH+pOH=14\text{pH} +
\text{pOH} = 14pH+pOH=14 This relationship comes from the ion product constant of water
(KwK_wKw), where Kw=[H+][OH−]=1×10−14K_w = [H^+][OH^-] = 1 \times 10^{-14}Kw
=[H+][OH−]=1×10−14 at 25°C.
3. The Role of Water in pH
Water undergoes self-ionization, forming hydrogen ions and hydroxide ions:
2H2O⇌H3O++OH−2H_2O \rightleftharpoons H_3O^+ + OH^-2H2O⇌H3O++OH− At 25°C,
pure water has a pH of 7 because [H3O+]=[OH−]=1×10−7 M[H_3O^+] = [OH^-] = 1 \times
10^{-7} \text{ M}[H3O+]=[OH−]=1×10−7 M.
The pH Scale in Detail
1. Acidic Solutions
Strong Acids: Completely dissociate in water. Examples include hydrochloric acid (HCl)
and sulfuric acid (H₂SO₄). They have very low pH values (close to 0).
Weak Acids: Partially dissociate in water. Examples include acetic acid (CH₃COOH)
and citric acid. They have higher pH values compared to strong acids but are still less
than 7.
2. Basic Solutions
Strong Bases: Completely dissociate in water. Examples include sodium hydroxide
(NaOH) and potassium hydroxide (KOH). They have high pH values (close to 14).
Weak Bases: Partially dissociate in water. Examples include ammonia (NH₃) and
bicarbonate ion (HCO₃⁻). They have lower pH values compared to strong bases but are
still greater than 7.
3. Neutral Solutions
Pure water at 25°C is neutral with a pH of 7. This neutrality arises from the equal concentrations
of hydrogen ions and hydroxide ions.
4. The pH Scale Range
0 to 3: Strongly acidic solutions (e.g., battery acid).
4 to 6: Weakly acidic solutions (e.g., vinegar).
7: Neutral solutions (e.g., pure water).
8 to 10: Weakly basic solutions (e.g., baking soda).
11 to 14: Strongly basic solutions (e.g., bleach).
Measuring pH
to enzyme activity
2.1 Concept of pH and the pH Scale
1. Definition of pH
The pH scale is a measure of how acidic or basic a solution is. The term "pH" stands for
"potential of hydrogen" or "power of hydrogen." It quantifies the concentration of hydrogen ions
(H⁺) in a solution. The pH scale ranges from 0 to 14:
pH < 7: Acidic solutions, with higher concentrations of hydrogen ions.
pH = 7: Neutral solutions, like pure water.
pH > 7: Basic (alkaline) solutions, with lower concentrations of hydrogen ions and
higher concentrations of hydroxide ions (OH⁻).
2. Historical Background
The concept of pH was introduced by Danish chemist Søren Sørensen in 1909. Sørensen was
working at the Carlsberg Laboratory in Copenhagen, and he sought a way to express the acidity
and alkalinity of solutions quantitatively.
3. Importance of pH
pH is crucial in various fields:
Chemistry: pH affects reaction rates, equilibria, and solubility.
Biology: Enzyme activity, cellular processes, and metabolic functions are pH-dependent.
Environmental Science: pH impacts ecosystems, water quality, and soil health.
Medicine: Body fluids' pH, such as blood and urine, are vital for diagnosing and
monitoring health conditions.
Fundamentals of pH
1. The pH Formula
The pH of a solution is calculated using the formula: pH=−log[H+]\text{pH} = -
\log[H^+]pH=−log[H+] where [H+][H^+][H+] is the concentration of hydrogen ions in moles
per liter (M).
2. The Relationship between pH and pOH
, The pH and pOH of a solution are related by the following equation: pH+pOH=14\text{pH} +
\text{pOH} = 14pH+pOH=14 This relationship comes from the ion product constant of water
(KwK_wKw), where Kw=[H+][OH−]=1×10−14K_w = [H^+][OH^-] = 1 \times 10^{-14}Kw
=[H+][OH−]=1×10−14 at 25°C.
3. The Role of Water in pH
Water undergoes self-ionization, forming hydrogen ions and hydroxide ions:
2H2O⇌H3O++OH−2H_2O \rightleftharpoons H_3O^+ + OH^-2H2O⇌H3O++OH− At 25°C,
pure water has a pH of 7 because [H3O+]=[OH−]=1×10−7 M[H_3O^+] = [OH^-] = 1 \times
10^{-7} \text{ M}[H3O+]=[OH−]=1×10−7 M.
The pH Scale in Detail
1. Acidic Solutions
Strong Acids: Completely dissociate in water. Examples include hydrochloric acid (HCl)
and sulfuric acid (H₂SO₄). They have very low pH values (close to 0).
Weak Acids: Partially dissociate in water. Examples include acetic acid (CH₃COOH)
and citric acid. They have higher pH values compared to strong acids but are still less
than 7.
2. Basic Solutions
Strong Bases: Completely dissociate in water. Examples include sodium hydroxide
(NaOH) and potassium hydroxide (KOH). They have high pH values (close to 14).
Weak Bases: Partially dissociate in water. Examples include ammonia (NH₃) and
bicarbonate ion (HCO₃⁻). They have lower pH values compared to strong bases but are
still greater than 7.
3. Neutral Solutions
Pure water at 25°C is neutral with a pH of 7. This neutrality arises from the equal concentrations
of hydrogen ions and hydroxide ions.
4. The pH Scale Range
0 to 3: Strongly acidic solutions (e.g., battery acid).
4 to 6: Weakly acidic solutions (e.g., vinegar).
7: Neutral solutions (e.g., pure water).
8 to 10: Weakly basic solutions (e.g., baking soda).
11 to 14: Strongly basic solutions (e.g., bleach).
Measuring pH
