Grade 11
  1. Basic concepts of chemistry  1.1. Importance of Chemistry  1.2. Nature and scope of chemistry  1.3. laws of chemical combination  1.3.1. Law of conservation of mass  1.3.2. Law of definite proportions  1.3.3. Law of multiple proportions  1.3.4. Law of gaseous volume  1.3.5. Avogadro's Law  1.4. Dalton's atomic theory  1.5. Mole concept and molar mass  1.6. Percent composition  1.7. Empirical and molecular formula  1.8. Stoichiometry and Stoichiometric Calculations  1.9. Limiting reagent concept
  2. Structure of the atom  2.1. Discovery of the electron, proton, and neutron  2.2. Atomic Model  2.2.1. Thomson's model  2.2.2. Rutherford's model  2.2.3. Bohr's model  2.3. Quantum mechanical model of the atom  2.4. Dual nature of matter and radiation  2.5. Heisenberg's uncertainty principle  2.6. Quantum numbers  2.7. Atomic orbitals and their shapes  2.8. Electronic configuration of elements  2.9. Aufbau principle  2.10. Pauli's exclusion principle  2.11. Hund's maximum multiplicity rule  2.12. de Broglie's hypothesis  2.13. Photoelectric effect
  3. Classification of elements and periodicity in properties  3.1. Historical development of the periodic table  3.2. Modern Periodic Law and Periodic Table  3.3. Periodic trends in properties  3.3.1. Atomic and Ionic Radius  3.3.2. Ionization enthalpy  3.3.3. Electron gain enthalpy  3.3.4. Electronegativity  3.3.5. Valency  3.3.6. Metallic and Non-Metallic Properties  3.3.7. Oxidation states  3.4. Unusual Properties of the First Elements
  4. Chemical Bonding and Molecular Structure  4.1. Kossel–Lewis approach to chemical bonding  4.2. Ionic or electrovalent bond  4.3. Covalent bond and its features  4.4. Bond parameters  4.5. VSEPR Theory  4.6. Valence bond theory  4.7. Hybridization and its types  4.8. Molecular orbital theory  4.8.1. Bond order and stability  4.9. Hydrogen bonding
  5. States of matter  5.1. Intermolecular forces  5.2. Gas Laws  5.2.1. Boyle's law  5.2.2. Charles's Law  5.2.3. Avogadro's Law  5.2.4. Dalton's law of partial pressure  5.2.5. Graham's law of spread  5.3. Ideal Gas Equation and Applications  5.4. Real gases and deviations from ideal behaviour  5.5. Kinetic molecular theory of gases  5.6. Liquefaction of gases  5.7. Properties of Liquids  5.8. Surface tension and viscosity
  6. Thermodynamics  6.1. System and environment  6.2. Types of systems in thermodynamics  6.3. Types of procedures  6.4. First law of thermodynamics  6.5. Internal energy and enthalpy  6.6. Heat capacity and specific heat  6.7. Hess's law of constant heat summation  6.8. Bond dissociation enthalpy  6.9. Enthalpy of formation and combustion  6.10. Enthalpy of solution and neutralization  6.11. Spontaneity and the second law of thermodynamics  6.12. Gibbs free energy and its applications
  7. Balance  7.1. The concept of balance  7.2. Equilibrium constant and its applications  7.3. Le Chatelier's principle  7.4. Ionic equilibrium in solutions  7.5. Acid and Base Theory  7.5.1. Arrhenius concept  7.5.2. Bronsted-Lowry concept  7.5.3. Lewis concept  7.6. pH Scale and pOH  7.7. Common ion effect  7.8. Hydrolysis of salts  7.9. Buffer solutions and their mechanism of action  7.10. Solubility equilibrium of sparingly soluble salts
  8. Redox reactions  8.1. Oxidation and reduction concepts  8.2. Oxidation number and its applications  8.3. Redox reactions in terms of electron transfer  8.4. Balancing redox reactions (oxidation number and ion-electron methods)  8.5. Types of redox reactions  8.6. Electrochemical Cells and Redox Reactions
  9. Hydrogen  9.1. Position of Hydrogen in the Periodic Table  9.2. Isotopes of hydrogen  9.3. Preparation of Hydrogen  9.4. Properties of Hydrogen  9.5. Hydrides and their classification  9.6. Water and heavy water  9.7. Hydrogen peroxide and its properties  9.8. Uses of Hydrogen and Its Compounds
  10. S-block elements (alkali and alkaline earth metals)  10.1. Group 1 Elements - Properties and Trends  10.2. Group 2 Elements - Properties and Trends  10.3. Unusual properties of lithium and beryllium  10.4. Important compounds of sodium and their uses  10.5. Important Compounds of Magnesium and Calcium
  11. Some p-block elements  11.1. General Introduction to p-Block Elements  11.2. Group 13 Elements  11.2.1. Boron and its compounds  11.3. Group 14 Elements  11.3.1. Carbon and its allotropes  11.3.2. Important Compounds of Carbon and Silicon
  12. Organic Chemistry - Some Basic Principles and Techniques  12.1. Classification and Nomenclature of Organic Compounds  12.2. Structural representation of organic compounds  12.3. Classification of organic reactions  12.4. Electronic Effects in Organic Chemistry  12.4.1. Inductive effect  12.4.2. Resonance effect  12.4.3. Hyperconjugation  12.5. Reaction mechanism and intermediate species  12.6. Purification and qualitative analysis of organic compounds
  13. Hydrocarbons  13.1. Hydrocarbons  13.1.1. Preparation and Properties of Alkenes  13.2. alkene  13.2.1. Preparation and Properties of Alkenes  13.2.2. Electrophilic addition reactions  13.3. Alkynes  13.3.1. Preparation and Properties of Alkynes  13.4. Aromatic hydrocarbons  13.4.1. Structure and properties of benzene  13.4.2. Electrophilic substitution reactions of benzene
  14. Environmental Chemistry  14.1. Environmental Pollution  14.2. Atmospheric pollution and the greenhouse effect  14.3. Water pollution and treatment methods  14.4. Soil pollution and its effects  14.5. Industrial waste and its treatment  14.6. Strategies for pollution control

Grade 11 → Balance ↓


pH Scale and pOH


The concepts of pH and pOH are important in understanding chemical equilibrium, especially in aqueous solutions. The pH scale is a measure of hydrogen ion concentration in a solution, while the pOH scale is a measure of hydroxide ion concentration. Both scales are logarithmic, meaning that each whole number change represents a tenfold change in concentration.

pH scale

The pH scale ranges from 0 to 14. A pH of 7 is considered neutral, which means there are equal concentrations of hydrogen ions and hydroxide ions. A pH less than 7 indicates an acidic solution, where there are more hydrogen ions than hydroxide ions. A pH greater than 7 indicates an alkaline solution, where there are fewer hydrogen ions than hydroxide ions. The formula for calculating pH is as follows:

pH = -log[H + ]

Here, [H + ] represents the concentration of hydrogen ions in moles per liter.

For example, if the hydrogen ion concentration in a solution is 0.001 moles per liter, the pH is calculated as follows:

pH = -log(0.001) = 3

This shows that the solution is acidic.

Visual example:

|--------|---------|---------|--------|---------|--------|---------|--------|
0        3         7         10       14
Acidic                            Neutral                            Basic
|--------|---------|---------|--------|---------|--------|---------|--------|

pOH scale

The pOH scale is similar to the pH scale, but it measures hydroxide ion concentration. The relationship between pH and pOH in water at 25°C is given by the formula:

pH + pOH = 14

This means that if you know the pH of a solution, you can easily find its pOH, and vice versa. The formula for calculating pOH is:

pOH = -log[OH - ]

Here, [OH - ] represents the concentration of hydroxide ions in moles per liter.

For example, if the concentration of hydroxide ions in a solution is 0.01 moles per liter, the pOH is calculated as:

pOH = -log(0.01) = 2

If the pOH is 2, the pH can be determined as follows:

pH = 14 - pOH = 14 - 2 = 12

This indicates an alkaline solution as the pH is greater than 7.

Visual example:

|--------|---------|---------|--------|---------|--------|---------|--------|
0        2         7         12       14
Basic                             Neutral                            Acidic
|--------|---------|---------|--------|---------|--------|---------|--------|

Neutralization and water balance

When an acid is mixed with a base, they neutralize each other. In the neutralization reaction between strong acids and bases, the products are usually water and salt. The equation for the neutralization reaction is as follows:

H + (aq) + OH - (aq) → H 2 O(l)

This reaction demonstrates the concept of equilibrium in water. In pure water at 25°C, the concentrations of hydrogen ions and hydroxide ions are equal at 1.0 x 10 -7 M. The equilibrium constant for water, K w, is given by:

K w = [H + ][OH - ] = 1.0 x 10 -14 at 25°C

From this we can conclude that:

pH = -log[H + ]
pOH = -log[OH - ]
pH + pOH = 14

Further exploration of the pH scale

The pH scale, being logarithmic, provides a wide range of values over which the hydrogen ion concentration can vary. For example, a pH change from 3 to 2 represents a tenfold increase in acidity. Understanding this scale helps us classify substances accurately:

  • Strong acids: Solutions such as hydrochloric acid (HCl) and sulfuric acid ( H2SO4 ) often have very low pH values , typically ranging from 1 to 3.
  • Weak acid: Acetic acid ( CH3COOH ) has a high pH close to the neutral point, usually ranging from 4 to 6.
  • Strong bases: Sodium hydroxide (NaOH) solutions have a high pH value, usually above 11.
  • Weak base: The pH value of ammonia ( NH3 ) solution ranges from 8 to 10.

A more detailed look at the examples helps to make this clear:

Example 1: Find the pH value of 0.1 M HCl solution.

[H + ] = 0.1 M
pH = -log(0.1) = 1

Example 2: Find the pH of 1.0 x 10 -4 M acetic acid solution, knowing that it has a low degree of ionization.

The dissociation of acetic acid is:

CH 3 COOH ⇌ H + + CH 3 COO -

[H + ] and then using the acid ionization constant ( K a ) to find the pH.

Understanding weak acid and weak base balance

Weak acids and bases do not completely dissociate in water, maintaining a balance between uncombined molecules and ions:

HA ⇌ H + + A -
BOH ⇌ B + + OH -

The equilibrium constants for these reactions are called K a for acids and K b for bases:

K a = [H + ][A -] / [HA]
K b = [B + ][OH -] / [BOH]

Example: Calculate the pH of a solution of acetic acid if K a = 1.8 x 10 -5 and [HA] = 0.1 M.

The expression for ionization constant is:

K a = [H + ][CH 3 COO -] / [CH 3 COOH]

Let [H + ] = [CH 3 COO - ] = x:

1.8 x 10 -5 = x² / 0.1-x
≈ x² / 0.1
x ≈ √(1.8 x 10 -5 * 0.1) = 1.34 x 10 -3
pH = -log(x) ≈ 2.87

Relation between pH and pOH

As discussed earlier, the relationship between pH and pOH is straightforward:

pH + pOH = 14

This relationship allows one to determine one value when another value is known, simplifying calculations in a variety of chemical contexts:

Example: If the pH value of a solution is 5, find its pOH.

pOH = 14 - pH = 14 - 5 = 9

Role in chemical equilibrium

Understanding pH and pOH is important in predicting the behavior of substances in equilibrium, particularly in the context of Le Chatelier's principle. The principle states that if changing conditions disturbs the dynamic equilibrium, the equilibrium position moves to counteract the change. Changes in pH can significantly affect equilibrium conditions:

  • In acid-base reactions, a change in the concentration of any of the reacting species causes a shift to restore equilibrium.
  • Control of pH is important in industrial and biological processes where sensitive reactions occur. Changes in pH can alter the rates and yields of reactions.

Practical applications

Understanding the pH and pOH scale is helpful in a variety of practical and industrial applications:

  • Environmental science: pH levels are important in assessing water quality, soil fertility, and pollution levels. Understanding pH can help predict and mitigate the effects of acid rain on ecosystems.
  • Medicine: Blood pH is very important for bodily functions. Maintaining proper pH balance is vital for health, and deviations in it can be a sign of medical problems.
  • Agriculture: Soil pH affects nutrient availability and microbial activity. Adjusting soil pH through additives creates conditions favorable for plant growth.
  • Food science: The processes of fermentation, preservation, and spoilage prevention depend on pH management. Controlling pH ensures safety and quality in the food industry.

Conclusion

The pH and pOH scales play an essential role in chemistry, especially in the study of equilibrium. Their logarithmic nature enables easy comparison and classification of solutions as acidic or alkaline, facilitating accurate prediction and control over chemical behavior in both natural and artificial systems.

Precise understanding and manipulation of these scales allows us to efficiently use chemical reactions in a variety of scientific, medical, and industrial fields, highlighting the relevance of these concepts in our daily lives.


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pH Scale and pOH

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