Understanding the construction and use of energy cycles for solution and hydration is pivotal in comprehending the thermodynamic principles governing chemical processes. This topic, situated within the chapter on Enthalpies of Solution and Hydration under the unit Chemical Energetics, is essential for AS & A Level Chemistry (9701) students. It provides foundational knowledge for predicting reaction behavior and designing efficient chemical systems.

Key Concepts

Enthalpy of Solution

The enthalpy of solution, denoted as $\Delta H_{\text{sol}}$, refers to the heat change that occurs when a solute dissolves in a solvent to form a solution. This thermodynamic quantity is crucial for understanding endothermic and exothermic dissolution processes. $$\Delta H_{\text{sol}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}$$ Where:

ΔH_lattice: Enthalpy of lattice, the energy required to break the intermolecular forces in the solute.
ΔH_hydration: Enthalpy of hydration, the energy released when solvent molecules surround and interact with solute ions.

The sign of $\Delta H_{\text{sol}}$ indicates whether the dissolution process absorbs or releases heat:

Endothermic Process: $\Delta H_{\text{sol}} > 0$, indicating heat absorption.
Exothermic Process: $\Delta H_{\text{sol}} < 0$, indicating heat release.

**Example:** Dissolving ammonium nitrate in water is an endothermic process: $$\text{NH}_4\text{NO}_3 (s) \rightarrow \text{NH}_4^+ (aq) + \text{NO}_3^- (aq)$$ The process absorbs heat, leading to a cooling effect.

Energy Cycles

Energy cycles are graphical representations that depict the energy changes during chemical processes. They are instrumental in visualizing the steps involved in the dissolution and hydration of solutes. **Construction of an Energy Cycle:**

**Reactants**: Start with the solid solute and solvent separately.
**Dissociation**: Show the breaking of solute's lattice energy.
**Hydration**: Illustrate the formation of interactions between solute ions and solvent molecules.
**Products**: Represent the final solution where solute is fully hydrated.

The energy cycle ensures that the first law of thermodynamics is upheld, showing that the total energy before and after the process remains constant.

Enthalpy Change of Hydration

The enthalpy change of hydration, $\Delta H_{\text{hydration}}$, is the heat released when one mole of ions is solvated by the solvent molecules. It plays a significant role in determining the overall enthalpy of solution. $$\Delta H_{\text{sol}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}$$ Factors Affecting $\Delta H_{\text{hydration}}$:

Charge Density: Ions with higher charge density exhibit stronger hydration due to greater electrostatic attractions with solvent molecules.
Size of Ions: Smaller ions have higher charge densities, leading to more energy released during hydration.

**Example:** Hydration of $\text{Mg}^{2+}$ ions releases more energy compared to $\text{Na}^{+}$ ions due to the higher charge density of $\text{Mg}^{2+}$.

Lattice Enthalpy

Lattice enthalpy, $\Delta H_{\text{lattice}}$, is the energy required to separate one mole of a solid ionic compound into gaseous ions. It reflects the strength of the ionic bonds within the crystal lattice. $$\Delta H_{\text{lattice}} = \text{Energy required to break down the ionic lattice}$$ **Factors Influencing Lattice Enthalpy:**

Charge of Ions: Higher charges result in stronger ionic bonds and greater lattice enthalpy.
Size of Ions: Smaller ions can pack more closely, increasing lattice enthalpy.

**Example:** $\text{MgO}$ has a higher lattice enthalpy compared to $\text{NaCl}$ due to the presence of $\text{Mg}^{2+}$ and $\text{O}^{2-}$ ions, which have higher charges and smaller sizes.

Thermodynamic Stability of Solutions

The thermodynamic stability of a solution is determined by the sign and magnitude of $\Delta H_{\text{sol}}$ and the entropy change $\Delta S$. According to the Gibbs free energy equation: $$\Delta G = \Delta H_{\text{sol}} - T\Delta S$$ For a solution to be thermodynamically favorable, $\Delta G$ must be negative. Even if $\Delta H_{\text{sol}}$ is positive (endothermic), a large positive $\Delta S$ can make $\Delta G$ negative, favoring dissolution. **Example:** Sodium chloride dissolution is endothermic, but the increase in entropy drives the process forward, making it thermodynamically favorable.

Heat Capacity and Temperature Effects

Heat capacity affects the temperature change observed during dissolution. Endothermic processes cause a temperature drop, while exothermic processes result in a temperature rise. $$q = m \cdot c \cdot \Delta T$$ Where:

q: Heat absorbed or released
m: Mass of the solution
c: Specific heat capacity
ΔT: Change in temperature

Understanding heat capacity is essential for calculating the energy changes associated with dissolution and hydration processes.

Solubility and Temperature Relationship

The solubility of a solute in a solvent is often temperature-dependent. For endothermic dissolution processes, solubility typically increases with temperature, while for exothermic processes, solubility decreases as temperature rises. **Example:** $\text{KNO}_3$ solubility increases with temperature (endothermic), whereas $\text{NH}_4\text{Cl}$ solubility decreases with temperature (exothermic).

Calculations Involving Enthalpy Changes

Calculations of enthalpy changes involve using Hess’s Law, which states that the total enthalpy change of a reaction is the same, regardless of the number of steps. **Example Problem:** Calculate the enthalpy change for the dissolution of $\text{NaCl}$ given:

$\Delta H_{\text{lattice}} (\text{NaCl}) = +787 \text{ kJ/mol}$
$\Delta H_{\text{hydration}} (\text{Na}^{+}) = -406 \text{ kJ/mol}$
$\Delta H_{\text{hydration}} (\text{Cl}^{-}) = -363 \text{ kJ/mol}$

**Solution:** $$\Delta H_{\text{sol}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}(\text{Na}^{+}) + \Delta H_{\text{hydration}}(\text{Cl}^{-})$$ $$\Delta H_{\text{sol}} = 787 + (-406) + (-363)$$ $$\Delta H_{\text{sol}} = 787 - 406 - 363$$ $$\Delta H_{\text{sol}} = +18 \text{ kJ/mol}$$ Therefore, the dissolution of $\text{NaCl}$ is slightly endothermic.

Factors Affecting Enthalpy of Solution

Several factors influence the enthalpy of solution, including:

Nature of Solute and Solvent: Polar solutes dissolve readily in polar solvents due to strong intermolecular interactions.
Temperature: As mentioned, temperature can affect both solubility and the enthalpy change during dissolution.
Pressure: Although pressure has a minimal effect on the dissolution of solids, it significantly affects the solubility of gases in liquids.

Advanced Concepts

Heat of Hydration and Born’s Equation

The heat of hydration can be quantitatively described using Born’s equation, which relates the enthalpy of hydration to the charge and radius of ions: $$\Delta H_{\text{hydration}} = -\frac{C \cdot z^2}{r}$$ Where:

C: A constant
z: Charge of the ion
r: Radius of the ion

This equation illustrates that ions with higher charges ($z$) and smaller radii ($r$) have more exothermic hydration enthalpies. Implications:

Comparing $\text{Mg}^{2+}$ and $\text{Na}^{+}$ ions, $\text{Mg}^{2+}$ has a higher charge and smaller radius, resulting in a more negative $\Delta H_{\text{hydration}}$.
This contributes to the high lattice enthalpy of $\text{MgO}$, requiring substantial energy for dissolution despite substantial hydration energy.

Solution Stoichiometry and Enthalpy Changes

In solution stoichiometry, understanding the stoichiometric ratios aids in calculating the heat changes during dissolution and hydration. Consider the dissolution of ionic compounds where stoichiometry plays a role in determining the total enthalpy change. **Example:** For the dissolution of $\text{CaCl}_2$: $$\text{CaCl}_2 (s) \rightarrow \text{Ca}^{2+} (aq) + 2\text{Cl}^{-} (aq)$$ The total $\Delta H_{\text{sol}}$ is calculated by summing the enthalpies involved, considering the stoichiometric coefficients.

Entropy and Enthalpy Interplay in Dissolution

Entropy ($\Delta S$) plays a crucial role alongside enthalpy in determining the spontaneity of dissolution. An increase in entropy usually favors the dissolution process. $$\Delta G = \Delta H_{\text{sol}} - T\Delta S$$ A positive $\Delta S$ can compensate for a positive $\Delta H_{\text{sol}}$, resulting in a negative $\Delta G$ and making the process spontaneous at higher temperatures. Case Study: Dissolution of $\text{NH}_4\text{NO}_3$ in water is highly endothermic but driven by a significant increase in entropy, making it thermodynamically favorable for cooling applications.

Solvation vs. Hydration

While hydrating refers specifically to solvation in water, solvation is a general term applicable to any solvent. The principles governing solvation are similar, but the enthalpies involved differ based on solvent properties. Comparative Analysis:

Hydration: Specific to water as the solvent, involving hydrogen bonding.
Solvation: Can involve various interactions depending on the solvent, such as dipole-dipole interactions or London dispersion forces.

Applications in Industrial Processes

Understanding energy cycles for solution and hydration is vital in various industrial applications, including:

Pharmaceuticals: Designing drug formulations that maximize solubility and stability.
Material Science: Synthesizing composites and nanomaterials with desired properties.
Environmental Engineering: Managing solubility of pollutants and designing remediation strategies.

Interdisciplinary Connections

The concepts of solution and hydration enthalpies intersect with other scientific disciplines:

Biochemistry: Understanding enzyme-substrate interactions and protein folding relies on solvation principles.
Environmental Science: Solubility of gases like CO₂ in oceans affects climate models.
Physics: Quantum chemistry provides deeper insights into the interactions at the molecular level.

Complex Problem-Solving

Advanced problems often require multi-step reasoning involving Hess’s Law, Born’s Equation, and thermodynamic principles. **Example Problem:** Calculate the enthalpy change for the dissolution of $\text{BaCl}_2$ in water, given:

$\Delta H_{\text{lattice}} (\text{BaCl}_2) = +965 \text{ kJ/mol}$
$\Delta H_{\text{hydration}} (\text{Ba}^{2+}) = -2526 \text{ kJ/mol}$
$\Delta H_{\text{hydration}} (\text{Cl}^{-}) = -363 \text{ kJ/mol}$

**Solution:** $$\Delta H_{\text{sol}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}} (\text{Ba}^{2+}) + 2 \times \Delta H_{\text{hydration}} (\text{Cl}^{-})$$ $$\Delta H_{\text{sol}} = 965 + (-2526) + 2 \times (-363)$$ $$\Delta H_{\text{sol}} = 965 - 2526 - 726$$ $$\Delta H_{\text{sol}} = -2,287 \text{ kJ/mol}$$ The dissolution of $\text{BaCl}_2$ is highly exothermic.

Experimental Determination of Enthalpy Changes

Techniques such as calorimetry are employed to experimentally determine enthalpy changes during dissolution and hydration. Calorimetry Process:

Setup: A calorimeter is used to measure temperature changes during a dissolution process.
Calculation: Using the formula $q = m \cdot c \cdot \Delta T$, the heat absorbed or released is calculated.
Example: Measuring the temperature drop when an endothermic salt dissolves in water to calculate $\Delta H_{\text{sol}}$.

Limitations and Challenges

Despite its usefulness, constructing energy cycles presents challenges:

Complexity of Real Systems: Real solutions may involve multiple solute species and interactions beyond simple models.
Accuracy of Data: Reliable thermodynamic data for all species involved may not always be available.
Assumptions in Models: Simplifications such as treating solvation as an ideal process can limit the applicability of energy cycles.

Advanced Thermodynamic Models

Beyond basic energy cycles, more sophisticated models account for factors like:

Non-Ideal Solutions: Incorporating activity coefficients to describe deviations from ideality.
Temperature Dependence: Modeling how enthalpy and entropy changes vary with temperature.
Pressure Effects: Although minor for solids, pressure can influence solutions involving gases.

Mathematical Derivations

Derivations involving Hess’s Law and thermodynamic cycles provide deeper insights into enthalpy changes. **Derivation of Enthalpy of Solution:** Starting from: $$\Delta H_{\text{sol}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}$$ Using Hess’s Law:

Breaking the lattice requires $\Delta H_{\text{lattice}}$.
Hydrating ions releases $\Delta H_{\text{hydration}}$.

Thus, the net enthalpy of solution is the sum of these two steps.

Comparison Table

Aspect	Endothermic Dissolution	Exothermic Dissolution
Enthalpy Change ($\Delta H_{\text{sol}}$)	Positive ($\Delta H_{\text{sol}} > 0$)	Negative ($\Delta H_{\text{sol}} < 0$)
Temperature Effect	Solubility increases with temperature	Solubility decreases with temperature
Heat Exchange	Absorbs heat from surroundings, causing cooling	Releases heat to surroundings, causing warming
Entropy Change ($\Delta S$)	Typically positive to drive dissolution	Can be positive or negative depending on the system
Examples	Ammonium nitrate ($\text{NH}_4\text{NO}_3$)	Sodium hydroxide ($\text{NaOH}$)

Summary and Key Takeaways

Energy cycles visually represent enthalpy changes during dissolution and hydration.
Enthalpy of solution combines lattice enthalpy and hydration enthalpy.
Thermodynamic favorability depends on both enthalpy and entropy changes.
Advanced concepts include mathematical models and experimental techniques.
Understanding these principles is crucial for applications across various scientific fields.

Examiner Tip

Tips

1. **Use Hess’s Law Strategically**: Break down complex enthalpy changes into manageable steps to simplify calculations.
2. **Memorize Key Formulas**: Ensure you know equations like $\Delta H_{\text{sol}} = \Delta H_{\text{lattice}} + \Delta H_{\text{hydration}}$ and Gibbs free energy.
3. **Visualize Energy Cycles**: Drawing energy cycles can help solidify your understanding of the energy transitions during dissolution and hydration.
4. **Practice Diverse Problems**: Engage with a variety of problems to strengthen your ability to apply concepts in different scenarios.

Did You Know

1. The renowned cold packs used in first aid leverage the endothermic dissolution of ammonium nitrate to absorb heat and provide instant cooling.
2. The concept of hydration enthalpy is fundamental in understanding how minerals dissolve in water, influencing geological processes like soil formation.
3. Supercooled solutions, where solutes remain dissolved below their normal freezing points, result from specific hydration dynamics.

Common Mistakes

1. **Confusing Enthalpy of Solution with Heat of Crystallization**: Students often mistake the enthalpy change during dissolution for the reverse process of crystallization. Remember, endothermic dissolution means exothermic crystallization.
2. **Ignoring Stoichiometry in Calculations**: Failing to account for the correct stoichiometric coefficients can lead to inaccurate enthalpy calculations. Always balance the equation and include proper molar ratios.
3. **Overlooking Entropy Contributions**: Assuming that only enthalpy determines spontaneity can result in incomplete analyses. Both enthalpy and entropy changes must be considered using Gibbs free energy.

FAQ

What is the difference between enthalpy of solution and enthalpy of hydration?

The enthalpy of solution ($\Delta H_{\text{sol}}$) is the total heat change when a solute dissolves in a solvent, while the enthalpy of hydration ($\Delta H_{\text{hydration}}$) specifically refers to the heat released when solvent molecules surround and stabilize solute ions.

How does temperature affect the solubility of endothermic and exothermic solutions?

For endothermic solutions, solubility increases with rising temperature as heat is absorbed. Conversely, for exothermic solutions, solubility decreases as temperature increases because heat is released.

Why is lattice enthalpy important in predicting solubility?

Lattice enthalpy indicates the strength of ionic bonds in a solid. High lattice enthalpy means more energy is required to break the ionic lattice, which can make the solute less soluble if not compensated by hydration enthalpy.

Can a solution be spontaneous if the enthalpy of solution is positive?

Yes, if the entropy change ($\Delta S$) is sufficiently positive, the Gibbs free energy ($\Delta G$) can still be negative, making the dissolution process spontaneous despite a positive $\Delta H_{\text{sol}}$.

How is calorimetry used to determine enthalpy changes?

Calorimetry measures the temperature change during a chemical process. Using the formula $q = m \cdot c \cdot \Delta T$, the heat absorbed or released can be calculated, allowing determination of enthalpy changes during dissolution or hydration.

1. Group 17

1.1 Physical Properties of the Group 17 Elements

1.1.1 Trends in Bond Strength of Halogen Molecules

1.1.2 Volatility Explained by Intermolecular Forces

1.1.3 Colours and Volatility of Chlorine, Bromine and Iodine

1.2 The Chemical Properties of the Halogen Elements and the Hydrogen Halides

1.2.1 Relative Reactivity of Halogen Elements as Oxidising Agents

1.2.2 Reactions of Halogens with Hydrogen

1.2.3 Thermal Stability of Hydrogen Halides

1.3 Some Reactions of the Halide Ions

1.3.1 Relative Reactivity of Halide Ions as Reducing Agents

1.3.2 Reactions of Halide Ions with Aqueous Silver Ions and Ammonia

1.3.3 Reactions of Halide Ions with Concentrated Sulfuric Acid

1.4 The Reactions of Chlorine

1.4.1 Reaction of Chlorine with Aqueous Sodium Hydroxide (Cold and Hot)

1.4.2 Use of Chlorine in Water Purification

2. Electrochemistry

2.1 Redox Processes: Electron Transfer and Changes in Oxidation Number

2.1.1 Calculating Oxidation Numbers

2.1.2 Using Oxidation Numbers to Balance Equations

2.1.3 Redox, Oxidation, Reduction, Disproportionation: Definitions

2.1.4 Oxidising and Reducing Agents: Definitions

2.1.5 Indicating Oxidation Numbers Using Roman Numerals

3. Chemical Energetics

3.1 Enthalpies of Solution and Hydration

3.1.1 Calculations Involving Solution and Hydration Energies

3.1.2 Effect of Ionic Charge and Radius on Hydration Enthalpy

3.1.3 Definition and Use of Enthalpy Change of Solution and Hydration

3.1.4 Construction and Use of Energy Cycles for Solution and Hydration

3.2 Entropy Change ΔS

3.2.1 Definition of Entropy

3.2.2 Prediction and Explanation of Entropy Changes

3.2.3 Calculation of Entropy Change for Reactions

3.3 Gibbs Free Energy Change ΔG

3.3.1 Gibbs Free Energy Equation

3.3.2 Calculations Using ΔG = ΔH – TΔS

3.3.3 Prediction of Reaction Feasibility

3.3.4 Effect of Temperature on Reaction Feasibility

3.4 Lattice Energy and Born–Haber Cycles

3.4.1 Definition and Use of Enthalpy Change of Atomisation and Lattice Energy

3.4.2 First Electron Affinity and Trends in Electron Affinity

3.4.3 Construction and Use of Born–Haber Cycles

3.4.4 Calculations Involving Born–Haber Cycles

3.4.5 Effect of Ionic Charge and Radius on Lattice Energy

4. Group 2

4.1 Magnesium to Barium, and Their Compounds

4.1.1 Variation in Solubility and Enthalpy Change of Solution of Hydroxides and Sulfates

4.1.2 Trend in Thermal Stability of Group 2 Nitrates and Carbonates

4.1.3 Effect of Ionic Radius on Polarisation of Large Anions

5. Atoms, Molecules and Stoichiometry

5.1 Formulas

5.1.1 Using Appropriate State Symbols

5.1.2 Empirical and Molecular Formulas: Definitions

5.1.3 Anhydrous, Hydrated and Water of Crystallisation

5.1.4 Calculation of Empirical and Molecular Formulas

5.1.5 Writing Ionic Compound Formulas from Ionic Charges and Oxidation Numbers

5.1.6 Predicting Ionic Charges from Periodic Table Positions

5.1.7 Common Ions: Names and Formulas

5.1.8 Writing and Balancing Chemical and Ionic Equations

5.2 Reacting Masses and Volumes (of Solutions and Gases)

5.2.1 Calculations Involving Reacting Masses and Percentage Yield

5.2.2 Calculations Involving Volumes of Gases

5.2.3 Calculations Involving Volumes and Concentrations of Solutions

5.2.4 Limiting Reagent and Excess Reagent Calculations

5.2.5 Deducing Stoichiometric Relationships from Calculations

5.3 Relative Masses of Atoms and Molecules

5.3.1 Unified Atomic Mass Unit

5.3.2 Relative Atomic Mass (Ar)

5.3.3 Relative Isotopic Mass

5.3.4 Relative Molecular Mass (Mr)

5.3.5 Relative Formula Mass

5.4 The Mole and the Avogadro Constant

5.4.1 Definition and Use of the Mole

5.4.2 Understanding the Avogadro Constant

6. Nitrogen Compounds

6.1 Primary and Secondary Amines

6.1.1 Production of Primary and Secondary Amines from Halogenoalkanes

6.1.2 Reduction of Amides and Nitriles to Amines

6.1.3 Reaction of Amines with Acyl Chlorides

6.1.4 Basicity of Aqueous Solutions of Amines

6.2 Phenylamine and Azo Compounds

6.2.1 Preparation of Phenylamine from Nitrobenzene

6.2.2 Reactions of Phenylamine with Bromine and Diazotisation

6.2.3 Relative Basicities of Ammonia, Ethylamine and Phenylamine

6.2.4 Formation and Use of Azo Compounds as Dyes

6.3 Amides

6.3.1 Production of Amides from Acyl Chlorides and Ammonia/Amines

6.3.2 Hydrolysis of Amides with Acids or Alkalis

6.3.3 Reduction of Amides to Amines

6.3.4 Weak Basic Nature of Amides

6.4 Amino Acids

6.4.1 Acid–Base Properties and Zwitterions

6.4.2 Formation of Peptide Bonds Between Amino Acids

6.4.3 Electrophoresis of Amino Acids and Peptides

7. Halogen Compounds

7.1 Halogenoalkanes

7.1.1 Elimination Reactions of Halogenoalkanes

7.1.2 SN1 and SN2 Mechanisms of Nucleophilic Substitution

7.1.3 Reactivity of Halogenoalkanes and Bond Strength

7.1.4 Production of Halogenoalkanes: Free-Radical Substitution and Electrophilic Addition

7.1.5 Classification of Halogenoalkanes: Primary, Secondary, Tertiary

7.1.6 Nucleophilic Substitution Reactions of Halogenoalkanes

7.1.7 Identification of Halogens by Reaction with Silver Nitrate

8. Chemistry of Transition Elements

8.1 General Physical and Chemical Properties of the First Row of Transition Elements

8.1.1 Titanium to Copper: Definition of Transition Elements

8.1.2 Titanium to Copper: Sketching 3dxy and 3dz² Orbitals

8.1.3 Titanium to Copper: General Properties: Variable Oxidation States, Catalytic Behavior, Formation of

8.1.4 Titanium to Copper: Explanation of Variable Oxidation States and Catalysis

8.2 General Characteristic Chemical Properties of the First Set of Transition Elements

8.2.1 Formation of Complexes and Ligand Types

8.2.2 Shapes and Coordination Numbers of Complexes

8.2.3 Ligand Exchange Reactions

8.2.4 Redox Reactions Involving Transition Elements

8.2.5 Redox Calculations Using E° Values

8.3 Colour of Complexes

8.3.1 Degenerate and Non-Degenerate d Orbitals

8.3.2 Splitting of d Orbitals in Octahedral and Tetrahedral Complexes

8.3.3 Explanation of Colour Based on Light Absorption

8.3.4 Effect of Ligand Types on Colour

8.4 Stereoisomerism in Transition Element Complexes

8.4.1 Geometrical (cis-trans) and Optical Isomerism in Complexes

8.4.2 Deduction of Overall Polarity of Complexes

8.5 Stability Constants

8.5.1 Definition and Expression of Stability Constants (Kstab)

8.5.2 Ligand Exchange and Kstab Values

9. Hydroxy Compounds

9.1 Alcohols

9.1.1 Production of Alcohols: Steam Addition, Oxidation, Reduction and Hydrolysis

9.1.2 Reactions of Alcohols: Combustion, Substitution, Oxidation, Dehydration, Esterification

9.1.3 Classification of Alcohols: Primary, Secondary, Tertiary

9.1.4 Tri-iodomethane Test for CH₃CH(OH)– Group

9.1.5 Acidity of Alcohols Compared to Water

10. Equilibria

10.1 Chemical Equilibria: Reversible Reactions and Dynamic Equilibrium

10.1.1 Reversible Reactions and Dynamic Equilibrium

10.1.2 Le Chatelier’s Principle

10.1.3 Effect of Temperature, Concentration, Pressure and Catalysts on Equilibrium

10.1.4 Equilibrium Constant Expressions (Kc)

10.1.5 Mole Fraction and Partial Pressure

10.1.6 Equilibrium Constant Expressions (Kp)

10.1.7 Calculations Using Kc and Kp

10.1.8 Changes Affecting the Value of Equilibrium Constants

10.1.9 Industrial Applications: Haber Process and Contact Process

10.2 Brønsted–Lowry Theory of Acids and Bases

10.2.1 Common Acids and Alkalis: Names and Formulas

10.2.2 Brønsted–Lowry Theory: Acids and Bases

10.2.3 Strong and Weak Acids and Bases

10.2.4 pH Scale and Differences Between Strong and Weak Acids

10.2.5 Neutralisation and Salt Formation

10.2.6 Acid–Base Titration Curves

10.2.7 Choosing Suitable Indicators for Titrations

11. Nitrogen and Sulfur

11.1 Nitrogen and Sulfur

11.1.1 Lack of Reactivity of Nitrogen Explained by Bond Strength and Polarity

11.1.2 Basicity of Ammonia and Structure of the Ammonium Ion

11.1.3 Displacement of Ammonia from Ammonium Salts

11.1.4 Natural and Man-made Sources of Nitrogen Oxides

11.1.5 Role of Nitrogen Oxides in Photochemical Smog Formation

11.1.6 Role of Nitrogen Oxides in Acid Rain Formation

12. Carbonyl Compounds

12.1 Aldehydes and Ketones

12.1.1 Production of Aldehydes and Ketones by Oxidation of Alcohols

12.1.2 Reduction of Aldehydes and Ketones

12.1.3 Reaction of Aldehydes and Ketones with Hydrogen Cyanide

12.1.4 Mechanism of Nucleophilic Addition Reactions

12.1.5 Detection of Carbonyl Compounds Using 2, 4-DNPH

12.1.6 Distinguishing Aldehydes from Ketones Using Fehling’s and Tollens’ Tests

12.1.7 Tri-iodomethane Test for CH₃CO– Group

13. Chemical Bonding

13.1 Electronegativity and Bonding

13.1.1 Definition of Electronegativity

13.1.2 Factors Affecting Electronegativity

13.1.3 Trends in Electronegativity Across a Period and Down a Group

13.1.4 Predicting Bond Type Using Electronegativity Differences

13.2 Ionic Bonding

13.2.1 Definition of Ionic Bonding

13.2.2 Ionic Bonding Examples: Sodium Chloride, Magnesium Oxide, Calcium Fluoride

13.3 Metallic Bonding

13.3.1 Definition of Metallic Bonding

13.4 Covalent Bonding and Coordinate (Dative Covalent) Bonding

13.4.1 Definition of Covalent Bonding

13.4.2 Covalent Bonding in Common Molecules (H₂, O₂, N₂, Cl₂, HCl, CO₂, NH₃, CH₄, C₂H₆, C₂H₄)

13.4.3 Expanded Octet in Period 3 Elements (SO₂, PCl₅, SF₆)

13.4.4 Definition of Coordinate (Dative Covalent) Bonding

13.4.5 Coordinate Bonding in NH₄⁺ and Al₂Cl₆

13.4.6 Sigma (σ) and Pi (π) Bonds: Orbital Overlap

13.4.7 Hybridisation: sp, sp² and sp³ Orbitals

13.4.8 Bond Energy and Bond Length Concepts

13.5 Shapes of Molecules

13.5.1 Shapes and Bond Angles Using VSEPR Theory (Examples: BF₃, CO₂, CH₄, NH₃, H₂O, SF₆, PF₅)

13.5.2 Predicting Shapes and Bond Angles of Similar Molecules and Ions

13.6 Intermolecular Forces, Electronegativity and Bond Properties

13.6.1 Hydrogen Bonding: Examples in Water and Ammonia

13.6.2 Anomalous Properties of Water Due to Hydrogen Bonding

13.6.3 Bond Polarity and Dipole Moments

13.6.4 Van der Waals' Forces: Instantaneous Dipole–Induced Dipole and Permanent Dipole–Dipole Forces

13.6.5 Comparison of Bond Strengths: Ionic, Covalent, Metallic vs. Intermolecular Forces

13.7 Dot-and-Cross Diagrams

13.7.1 Using Dot-and-Cross Diagrams to Represent Ionic, Covalent and Coordinate Bonding

14. Electrochemistry

14.1 Electrolysis

14.1.1 Prediction of Products During Electrolysis

14.1.2 Faraday Constant and Relationship with Avogadro Constant and Electron Charge

14.1.3 Calculations Involving Charge, Mass and Volume in Electrolysis

14.1.4 Determination of Avogadro Constant by Electrolysis

14.2 Standard Electrode Potentials and the Nernst Equation

14.2.1 Definition of Standard Electrode Potential and Standard Cell Potential

14.2.2 Description of the Standard Hydrogen Electrode

14.2.3 Measurement of Standard Electrode Potentials

14.2.4 Calculation of Standard Cell Potentials

14.2.5 Deduction of Electron Flow and Reaction Feasibility Using Electrode Potentials

14.2.6 Reactivity Predictions Using E° Values

14.2.7 Construction of Redox Equations from Half-Equations

14.2.8 Variation of Electrode Potential with Concentration

14.2.9 Using the Nernst Equation for Electrode Potentials

14.2.10 Relationship Between ΔG° and E°cell

15. Introduction to Organic Chemistry

15.1 Formulas, Functional Groups and the Naming of Organic Compounds

15.1.1 Definition of Hydrocarbons

15.2 Formulas

15.2.1 Functional Groups and Physical/Chemical Properties

15.2.2 Interpretation of General, Structural, Displayed and Skeletal Formulas

15.2.3 Systematic Nomenclature of Simple Aliphatic Organic Compounds

15.2.4 Deducing Molecular and Empirical Formulas

15.3 Characteristic Organic Reactions

15.3.1 Terminology of Organic Reactions: Homologous Series, Saturated/Unsaturated, Fission, Radicals, Nucle

15.3.2 Types of Organic Reactions: Addition, Substitution, Elimination, Hydrolysis, Condensation, Oxidation

15.3.3 Mechanisms of Organic Reactions: Free Radical Substitution, Electrophilic Addition, Nucleophilic Sub

15.4 Shapes of Organic Molecules; σ and π Bonds

15.4.1 Shapes and Bond Angles in Organic Molecules

15.4.2 Sigma (σ) and Pi (π) Bond Arrangements

15.4.3 Planarity in Organic Molecules

15.5 Isomerism: Structural Isomerism and Stereoisomerism

15.5.1 Types of Structural Isomerism: Chain, Positional, Functional Group

15.5.2 Types of Stereoisomerism: Geometrical (cis-trans) and Optical

15.5.3 Geometrical Isomerism in Alkenes

15.5.4 Chirality and Optical Isomerism

15.5.5 Identifying Chiral Centres and Isomer Types

16. Genetic technology

16.1 Genetically modified organisms in agriculture

16.1.1 GM crops and animals and their implications

16.2 Primary Amines

16.2.1 Production of Primary Amines by Reaction of Halogenoalkanes with Ammonia

16.3 Nitriles and Hydroxynitriles

16.3.1 Production of Nitriles by Reaction of Halogenoalkanes with KCN

16.3.2 Production of Hydroxynitriles by Reaction of Aldehydes and Ketones with HCN

16.3.3 Hydrolysis of Nitriles to Produce Carboxylic Acids

17. Atomic Structure

17.1 Particles in the Atom and Atomic Radius

17.1.1 Structure of the Atom: Nucleus and Electron Shells

17.1.2 Subatomic Particles: Protons, Neutrons, Electrons

17.1.3 Atomic Number, Proton Number, Mass Number, Nucleon Number

17.1.4 Mass and Charge Distribution in Atoms

17.1.5 Deflection of Charged Particles in Electric Fields

17.1.6 Calculations Involving Protons, Neutrons, Electrons

17.1.7 Trends in Atomic and Ionic Radii Across Periods and Down Groups

17.2 Isotopes

17.2.1 Definition and Notation of Isotopes

17.2.2 Chemical Properties of Isotopes

17.2.3 Physical Properties of Isotopes: Mass and Density Differences

17.3 Electrons

17.3.1 Energy Levels, and Atomic Orbitals: Shells, Sub-shells and Orbitals

17.3.2 Energy Levels,and Atomic Orbitals: Principal Quantum Number (n) and Ground State Configurations

17.3.3 Energy Levels and Atomic Orbitals: Order of Sub-shell Energy Levels (s, p, d)

17.3.4 Energy Levels and Atomic Orbitals: Full and Shorthand Electronic Configurations

17.3.5 Energy Levels and Atomic Orbitals: Electron Box Diagrams

17.3.6 Energy Levels and Atomic Orbitals: Shapes of s and p Orbitals

17.3.7 Energy Levels and Atomic Orbitals: Free Radicals and Unpaired Electrons

17.4 Ionisation Energy

17.4.1 First Ionisation Energy: Definition and Equations

17.4.2 Trends in Ionisation Energy Across Periods and Down Groups

17.4.3 Successive Ionisation Energies and Electron Shells

17.4.4 Factors Affecting Ionisation Energy: Nuclear Charge, Radius, Shielding

17.4.5 Deducing Electronic Configurations Using Ionisation Energy Data

17.4.6 Position of Elements Using Ionisation Energy Patterns

18. Polymerisation (Condensation Polymers)

18.1 Condensation Polymerisation

18.1.1 Formation of Polyesters from Diols and Dicarboxylic Acids or Dioyl Chlorides

18.1.2 Formation of Polyamides from Diamines and Dicarboxylic Acids or Dioyl Chlorides

18.1.3 Formation of Polymers from Amino Acids

18.2 Predicting the Type of Polymerisation

18.2.1 Identifying the Type of Polymerisation from Monomers or Polymer Structure

18.3 Degradable Polymers

18.3.1 Chemical Inertness and Disposal of Poly(alkenes)

18.3.2 Biodegradability of Polyesters and Polyamides

18.3.3 Degradation by Light

19. Carboxylic Acids and Derivatives

19.1 Carboxylic Acids

19.1.1 Production of Carboxylic Acids: Oxidation, Hydrolysis of Nitriles and Esters

19.1.2 Reactions of Carboxylic Acids with Metals, Alkalis and Carbonates

19.1.3 Esterification Reactions with Alcohols

19.1.4 Reduction of Carboxylic Acids

19.1.5 Relative Acidity of Carboxylic Acids, Phenols and Alcohols

19.1.6 Acidity of Chlorine-Substituted Carboxylic Acids

19.2 Esters

19.2.1 Production of Esters by Condensation of Carboxylic Acids and Alcohols

19.2.2 Hydrolysis of Esters with Acid or Alkali

20. Introduction to A Level Organic Chemistry

20.1 Formulas, Functional Groups and the Naming of Organic Compounds

20.1.1 Functional Groups in Aromatic and Extended Organic Molecules

20.1.2 Use of Structural, Displayed and Skeletal Formulas

20.1.3 Systematic Nomenclature for Aliphatic and Aromatic Molecules

20.2 Characteristic Organic Reactions

20.2.1 Electrophilic Substitution and Addition–Elimination Mechanisms

20.3 Shapes of Aromatic Organic Molecules; σ and π Bonds

20.3.1 Shape and Bonding in Benzene: sp² Hybridisation and Delocalised π System

20.4 Isomerism: Optical

20.4.1 Physical and Biological Properties of Optical Isomers

20.4.2 Concepts: Optically Active Compounds, Racemic Mixtures, Chiral Catalysts

21. Reaction Kinetics

21.1 Rate of Reaction

21.1.1 Definition of Rate of Reaction

21.1.2 Collision Theory: Effective and Non-Effective Collisions

21.1.3 Effect of Concentration and Pressure on Reaction Rate

21.1.4 Calculating Reaction Rate from Experimental Data

21.2 Effect of Temperature on Reaction Rates and the Concept of Activation Energy

21.2.1 Definition of Activation Energy

21.2.2 Boltzmann Distribution and Activation Energy

21.2.3 Effect of Temperature on Reaction Rate and Collision Frequency

21.3 Homogeneous and Heterogeneous Catalysts

21.3.1 Definition and Mechanism of Catalysis

21.3.2 Catalysts Lowering Activation Energy

21.3.3 Reaction Pathway Diagrams With and Without Catalysts

22. Organic Synthesis

22.1 Organic Synthesis

22.1.1 Identification of Functional Groups in Organic Molecules

22.1.2 Prediction of Properties and Reactions of Organic Molecules

22.1.3 Designing Multi-Step Synthetic Routes

22.1.4 Analysis of Synthetic Routes and Possible By-products

23. Hydrocarbons (Arenes)

23.1 Arenes

23.1.1 Substitution Reactions of Benzene and Methylbenzene

23.1.2 Mechanism of Electrophilic Substitution in Arenes

23.1.3 Predicting Halogenation: Side-Chain vs. Aromatic Ring

23.1.4 Directing Effects of Substituents in Electrophilic Substitution

24. Equilibria

24.1 Acids and Bases

24.1.1 Conjugate Acids and Bases

24.1.2 pH, Ka, pKa and Kw: Definitions and Calculations

24.1.3 Calculations of pH for Strong Acids, Strong Alkalis and Weak Acids

24.1.4 Definition and Formation of Buffer Solutions

24.1.5 Buffer Solutions in pH Control and Blood Chemistry

24.1.6 Calculations Involving Buffer Solutions

24.1.7 Solubility Product (Ksp) and Calculations

24.1.8 Common Ion Effect and Solubility Calculations

24.2 Partition Coefficients

24.2.1 Definition and Calculation of Partition Coefficient (Kpc)

24.2.2 Factors Affecting Partition Coefficients Based on Solute and Solvent Polarity

25. Analytical Techniques

25.1 Thin-Layer Chromatography

25.1.1 Terms and Interpretation: Stationary Phase, Mobile Phase, Rf Value, Solvent Front, Baseline

25.1.2 Explanation of Differences in Rf Values

25.2 Gas / Liquid Chromatography

25.2.1 Terms and Interpretation: Stationary Phase, Mobile Phase, Retention Time

25.2.2 Composition Analysis and Retention Time Explanation

25.3 Carbon-13 NMR Spectroscopy

25.3.1 Interpretation of C-13 NMR Spectra

25.3.2 Prediction of Number and Type of C Environments

25.4 Proton (¹H) NMR Spectroscopy

25.4.1 Interpretation of ¹H NMR Spectra: Chemical Shifts, Peak Areas, Splitting Patterns (n+1 Rule)

25.4.2 Prediction of Spectral Features

25.4.3 Use of TMS and Deuterated Solvents

25.4.4 Identification of O–H and N–H Protons by D₂O Exchange

26. Halogen Compounds (Arenes)

26.1 Halogen Compounds

26.1.1 Production of Halogenoarenes by Substitution

26.1.2 Reactivity Comparison: Halogenoalkanes vs. Halogenoarenes

27. Hydroxy Compounds (Phenol)

27.1 Alcohols

27.1.1 Reaction of Alcohols with Acyl Chlorides to Form Esters

27.2 Phenol

27.2.1 Directing Effects of Hydroxyl Group in Phenol Reactions

27.2.2 Reactions of Phenol with Bases, Sodium and Diazonium Salts

27.2.3 Nitration and Bromination of Phenol

27.2.4 Acidity of Phenol Compared to Water and Ethanol

27.2.5 Different Reaction Conditions for Phenol vs. Benzene

27.2.6 Production of Phenol from Diazonium Salts

28. Polymerisation

28.1 Addition Polymerisation

28.1.1 Description of Addition Polymerisation

28.1.2 Deduction of Repeat Units from Monomers

28.1.3 Identification of Monomers from Polymer Structures

28.1.4 Challenges of Disposal: Non-biodegradability and Harmful Combustion Products

29. Hydrocarbons

29.1 Alkanes

29.1.1 Production of Alkanes: Hydrogenation and Cracking

29.1.2 Combustion of Alkanes: Complete and Incomplete

29.1.3 Free-Radical Substitution of Alkanes

29.1.4 Mechanism of Free-Radical Substitution: Initiation, Propagation, Termination

29.1.5 Cracking for Useful Alkanes and Alkenes

29.1.6 Environmental Impact of Alkane Combustion

29.2 Alkenes

29.2.1 Production of Alkenes: Elimination and Dehydration Reactions

29.2.2 Electrophilic Addition Reactions of Alkenes

29.2.3 Oxidation of Alkenes: Cold Dilute and Hot Concentrated KMnO₄

29.2.4 Test for C=C Bond Using Aqueous Bromine

29.2.5 Mechanism of Electrophilic Addition in Alkenes

29.2.6 Stability of Carbocations and Markovnikov’s Rule

30. States of Matter

30.1 The Gaseous State: Ideal and Real Gases and pV = nRT

30.1.1 Origin of Pressure in Gases

30.1.2 Ideal Gas Assumptions

30.1.3 Ideal Gas Equation: pV = nRT

30.1.4 Applications of the Ideal Gas Equation

30.2 Bonding and Structure

30.2.1 Lattice Structures of Crystalline Solids

30.2.2 Structures and Bonding: Effect on Physical Properties

30.2.3 Deducing Structure and Bonding from Data

31. The Periodic Table: Chemical Periodicity

31.1 Periodicity of Physical Properties of the Elements in Period 3

31.1.1 Trends in Atomic Radius, Ionic Radius, Melting Point and Electrical Conductivity

31.1.2 Explanation of Melting Point and Conductivity Based on Structure and Bonding

31.2 Periodicity of Chemical Properties of the Elements in Period 3

31.2.1 Reactions of Elements with Oxygen, Chlorine and Water

31.2.2 Oxidation Numbers of Oxides and Chlorides

31.2.3 Reactions of Oxides with Water and pH of Solutions

31.2.4 Acid-Base Behaviour of Oxides and Hydroxides

31.2.5 Reactions of Chlorides with Water and pH of Solutions

31.2.6 Trends Explained by Bonding and Electronegativity

31.2.7 Bonding Types from Chemical and Physical Properties

31.3 Chemical Periodicity of Other Elements

31.3.1 Predicting Properties Based on Group Trends

31.3.2 Deducing Element Identity from Physical and Chemical Data

32. Organic Synthesis

32.1 Organic Synthesis

32.1.1 Identification of Organic Functional Groups

32.1.2 Predicting Properties and Reactions

32.1.3 Designing Multi-Step Synthetic Routes

32.1.4 Analyzing Synthetic Routes and Identifying By-products

33. Reaction Kinetics

33.1 Simple Rate Equations, Orders of Reaction and Rate Constants

33.1.1 Terms: Rate Equation, Order of Reaction, Overall Order, Rate Constant, Half-life, Rate-determining S

33.1.2 Deducing Orders from Graphs or Experimental Data

33.1.3 Interpretation of Concentration–Time and Rate–Concentration Graphs

33.1.4 Calculating Initial Rate from Data

33.1.5 Half-life of First-Order Reactions and Calculations

33.1.6 Calculating Rate Constants Using Different Methods

33.1.7 Predicting Reaction Mechanisms and Rate-Determining Steps

33.1.8 Identifying Catalysts and Intermediates from Mechanisms

33.1.9 Effect of Temperature on Rate Constants

33.2 Homogeneous and Heterogeneous Catalysts

33.2.1 Mode of Action of Heterogeneous Catalysts (Adsorption and Desorption)

33.2.2 Examples of Industrial Heterogeneous Catalysis

33.2.3 Mode of Action of Homogeneous Catalysts

33.2.4 Examples of Homogeneous Catalysis in Atmospheric Chemistry

34. Analytical Techniques

34.1 Infrared Spectroscopy

34.1.1 Analysis of Infrared Spectra to Identify Functional Groups

34.2 Mass Spectrometry

34.2.1 Interpretation of Mass Spectra: m/e Values and Isotopic Abundances

34.2.2 Calculation of Relative Atomic Mass and Molecular Mass

34.2.3 Identification of Molecules by Fragmentation Patterns

34.2.4 Determination of Carbon Atom Number Using [M+1]+ Peak

34.2.5 Detection of Bromine and Chlorine Atoms Using [M+2]+ Peak

35. Chemical Energetics

35.1 Enthalpy Change ΔH

35.1.1 Exothermic and Endothermic Reactions

35.1.2 Reaction Pathway Diagrams: Enthalpy Change and Activation Energy

35.1.3 Standard Conditions and Standard Enthalpy Changes

35.1.4 Enthalpy Changes: Reaction, Formation, Combustion, Neutralisation

35.1.5 Energy Transfer in Breaking and Making Bonds

35.1.6 Calculations Using Bond Energies

35.1.7 Calculations Using Experimental Data: q = mcΔT and ΔH = –mcΔT/n

35.2 Hess’s Law

35.2.1 Application of Hess’s Law to Construct Energy Cycles

35.2.2 Calculations Using Hess’s Law and Bond Energy Data

36. Group 2

36.1 Magnesium to Barium, and their Compounds

36.1.1 Reactions of Group 2 Elements with Oxygen, Water and Acids

36.1.2 Reactions of Group 2 Oxides, Hydroxides and Carbonates with Water and Acids

36.1.3 Thermal Decomposition of Group 2 Nitrates and Carbonates

36.1.4 Trends in Physical and Chemical Properties

36.1.5 Variation in Solubility of Hydroxides and Sulfates

37. Carboxylic Acids and Derivatives (Aromatic)

37.1 Esters

37.1.1 Production of Esters from Alcohols and Acyl Chlorides

37.2 Acyl Chlorides

37.2.1 Comparison of Hydrolysis of Acyl Chlorides, Alkyl Chlorides and Aryl Chlorides

37.2.2 Production of Acyl Chlorides from Carboxylic Acids

37.2.3 Addition–Elimination Mechanism of Acyl Chlorides

37.2.4 Reactions of Acyl Chlorides with Water, Alcohols, Phenols, Ammonia and Amines

37.3 Carboxylic Acids

37.3.1 Further Oxidation of Methanoic Acid and Ethanedioic Acid

37.3.2 Reactions of Carboxylic Acids with PCl₃, PCl₅ and SOCl₂

37.3.3 Effect of Chlorine Substitution on Acid Strength

37.3.4 Relative Acidities of Carboxylic Acids, Phenols and Alcohols

37.3.5 Production of Benzoic Acid from Alkylbenzene

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Construction and Use of Energy Cycles for Solution and Hydration

Introduction