1. Structure: Classification of Matter

1.1 Functional Groups: Classification of Organic Compounds

1.1.1 IUPAC naming conventions for organic compounds

1.1.2 Introduction to organic chemistry and functional groups

1.1.3 Alcohols, aldehydes, ketones, carboxylic acids, and amines

1.1.4 Isomerism and structural formulas

1.2 The Periodic Table: Classification of Elements

1.2.1 Periodic trends: Atomic radius, ionization energy, electronegativity

1.2.2 Group classification of elements

1.2.3 Transition metals and their unique properties

2. Reactivity: How Much, How Fast, and How Far?

2.1 How Far? The Extent of Chemical Change

2.1.1 Calculating equilibrium constants (Kc)

2.1.2 Shifting equilibrium: Concentration, pressure, and temperature effects

2.1.3 Chemical equilibrium and Le Chatelier’s Principle

2.2 How Much? The Amount of Chemical Change

2.2.1 Stoichiometry and mole-to-mole ratios

2.2.2 Limiting reagents and theoretical yield

2.2.3 Concentration, volume, and number of particles

2.3 How Fast? The Rate of Chemical Change

2.3.1 Factors affecting reaction rates: Concentration, temperature, surface area, catalysts

2.3.2 Rate law and order of reactions

2.3.3 Collision theory and activation energy

3. Structure: Models of Bonding and Structure

3.1 The Ionic Model

3.1.1 Properties of ionic compounds

3.1.2 Lattice structure in ionic solids

3.1.3 Conductivity and solubility of ionic compounds

3.1.4 Ionic bonding and formation of ions

3.2 The Covalent Model

3.2.1 Covalent bonding and electron sharing

3.2.2 Single, double, and triple bonds

3.2.3 Polar vs non-polar covalent bonds

3.2.4 Bonding in simple molecules (e.g., H₂O, CO₂)

3.3 The Metallic Model

3.3.1 Metallic bonding and electron sea model

3.3.2 Electrical conductivity of metals

3.3.3 Properties of metals: Malleability, ductility, high melting points

3.4 From Models to Materials

3.4.1 Solids, liquids, and gases: Structural differences

3.4.2 Physical properties of materials

3.4.3 Nanotechnology and applications

3.4.4 Alloys and their properties

4. Experimental Programme

4.1 Collaborative Sciences Project

4.1.1 Effective communication in scientific collaboration

4.1.2 Group-based scientific inquiry

4.1.3 Interdisciplinary applications of chemistry

4.2 Scientific Investigation

4.2.1 Formulating hypotheses and research questions

4.2.2 Designing experiments and gathering data

4.2.3 Writing and presenting scientific reports

4.3 Practical Work

4.3.1 Laboratory safety protocols and techniques

4.3.2 Common chemical reactions in the laboratory

4.3.3 Data collection, analysis, and interpretation

4.3.4 Precision, accuracy, and error analysis

5. Reactivity: What Drives Chemical Reactions?

5.1 Measuring Enthalpy Change

5.1.1 Enthalpy of reaction, heat capacity, and calorimetry

5.1.2 Exothermic vs endothermic reactions

5.1.3 Hess's Law and enthalpy cycles

5.2 Energy Cycles in Reactions

5.2.1 Thermochemistry and enthalpy diagrams

5.2.2 Born-Haber cycle and lattice enthalpy

5.2.3 Bond dissociation and bond formation energies

5.3 Energy from Fuels

5.3.1 Combustion reactions and energy release

5.3.2 Fuels and their efficiency

5.3.3 Energy density of different fuels (e.g., hydrocarbons, biofuels)

5.4 Entropy and Spontaneity (Additional Higher Level)

5.4.1 Entropy and the second law of thermodynamics

5.4.2 Spontaneous vs non-spontaneous processes

5.4.3 Gibbs free energy and its application to chemical reactions

6. Reactivity: What Are the Mechanisms of Chemical Change?

6.1 Proton Transfer Reactions

6.1.1 Acid-base reactions: Bronsted-Lowry theory

6.1.2 Strong vs weak acids and bases

6.1.3 pH scale and calculations

6.2 Electron Transfer Reactions

6.2.1 Redox reactions and electron transfer

6.2.2 Oxidizing and reducing agents

6.2.3 Electrochemical cells and half-reactions

6.3 Electron Sharing Reactions

6.3.1 Covalent bonds and electron sharing

6.3.2 Polar vs non-polar covalent bonds in molecular compounds

6.3.3 Electronegativity and bond polarity

6.4 Electron-Pair Sharing Reactions

6.4.1 Lewis acids and bases

6.4.2 Coordinate covalent bonding

6.4.3 Complex ion formation

7. Structure: Models of the Particulate Nature of Matter

7.1 Introduction to the Particulate Nature of Matter

7.1.1 Properties of matter

7.1.2 Molecular theory and the nature of matter

7.1.3 Solid, liquid, gas phases and changes of state

7.2 The Nuclear Atom

7.2.1 Structure of the atom: Protons, neutrons, electrons

7.2.2 Atomic number, mass number, isotopes

7.2.3 Atomic models: Bohr model, quantum model

7.3 Electron Configurations

7.3.1 Electron arrangement and atomic orbitals

7.3.2 Aufbau principle, Pauli exclusion principle, Hund's rule

7.3.3 Periodicity of electron configurations

7.4 Counting Particles by Mass: The Mole

7.4.1 Definition of the mole and Avogadro's constant

7.4.2 Molar mass and its calculation

7.4.3 Conversions between moles, mass, and number of particles

7.4.4 Empirical and molecular formulas

7.5 Ideal Gases

7.5.1 Ideal gas law (PV = nRT)

7.5.2 Gas laws: Boyle's law, Charles's law, Avogadro's law

7.5.3 Real gases vs ideal gases

7.5.4 Applications of the ideal gas law

Born-Haber cycle and lattice enthalpy

Topic 2/3

Revision Notes
Flashcards
Past Paper Analysis
Questions
Videos

Your Flashcards are Ready!

15 Flashcards in this deck.

Born-Haber Cycle and Lattice Enthalpy

Introduction

The Born-Haber cycle and lattice enthalpy are fundamental concepts in inorganic chemistry, particularly within the context of ionic compound formation. Understanding these principles is crucial for IB Chemistry SL students as they provide insights into the energetics of ionic bonds and the stability of crystalline structures. This article explores the intricacies of the Born-Haber cycle and lattice enthalpy, elucidating their roles in chemical reactions and their significance in predicting compound formation.

Key Concepts

Born-Haber Cycle

The Born-Haber cycle is a thermochemical cycle that illustrates the steps involved in the formation of an ionic compound from its constituent elements in their standard states. It applies Hess's Law to calculate the lattice enthalpy by breaking down the overall formation reaction into a series of intermediate steps, each with its associated enthalpy change. The cycle typically includes the following steps:

Formation of Gaseous Ions: The metal atom is ionized to form a cation, and the non-metal molecule is dissociated into atoms and then gains electrons to form anions.
Formation of Sublimated Metal: The metal is sublimated from its solid state to a gaseous state.
Dissociation Enthalpy: The energy required to dissociate the non-metal molecules into individual atoms.
Ionization Energy: The energy needed to remove electrons from the metal atom to form cations.
Electron Affinity: The energy change when electrons are added to the non-metal atoms to form anions.
Lattice Enthalpy: The energy released when gaseous ions combine to form the solid ionic lattice.

By applying Hess's Law, the sum of the enthalpy changes of these steps equals the overall enthalpy of formation of the ionic compound. The Born-Haber cycle is instrumental in determining lattice enthalpy, which is otherwise challenging to measure directly.

Lattice Enthalpy

Lattice enthalpy ($ \Delta H_{lattice} $) refers to the energy released when gaseous ions come together to form an ionic solid. It is a measure of the strength of the ionic bonds within the crystal lattice. A higher lattice enthalpy indicates a more stable and tightly bound ionic compound. Lattice enthalpy is influenced by two primary factors:

Charge of the Ions: Higher charges on the cations and anions lead to stronger electrostatic attractions, resulting in greater lattice enthalpy.
Size of the Ions: Smaller ions can pack closer together, enhancing the attraction between them and increasing the lattice enthalpy.

Mathematically, lattice enthalpy can be estimated using Coulomb's Law: $$ \Delta H_{lattice} = \frac{K \cdot Q_1 \cdot Q_2}{r} $$ where $ K $ is a proportionality constant, $ Q_1 $ and $ Q_2 $ are the charges of the ions, and $ r $ is the distance between the ion centers.

Application of Born-Haber Cycle

The Born-Haber cycle is widely used to calculate lattice enthalpy indirectly. For instance, to determine the lattice enthalpy of sodium chloride (NaCl), the cycle involves the following steps:

Sublimation of Na(s) to Na(g): $ \Delta H_{sublimation} $
Dissociation of Cl$_2$(g) to 2Cl(g): $ \frac{1}{2} \Delta H_{dissociation} $
Ionization of Na(g) to Na$^+$(g) + e\(^-$): $ \Delta H_{ionization} $
Electron gain by Cl(g) to form Cl$^-$(g): \( \Delta H_{electron affinity} $
Formation of NaCl(s) from Na$^+$(g) and Cl\(^-$(g): \( \Delta H_{lattice} $

By applying Hess's Law, the sum of the enthalpy changes for steps 1 through 4 equals the overall enthalpy of formation of NaCl, allowing the calculation of $ \Delta H_{lattice} $.

Lattice Enthalpy and Properties of Ionic Compounds

Lattice enthalpy significantly influences the physical properties of ionic compounds. High lattice enthalpy correlates with:

High Melting and Boiling Points: Strong ionic bonds require more energy to break, resulting in higher melting and boiling points.
Solubility: Ionic compounds with high lattice enthalpy may be less soluble in water, as more energy is needed to overcome the lattice forces during dissolution.

Conversely, compounds with lower lattice enthalpy tend to have lower melting points and higher solubility.

Factors Affecting Lattice Enthalpy

Several factors influence lattice enthalpy:

Ion Charge: As the magnitude of the ion charges increases, lattice enthalpy increases due to stronger electrostatic forces.
Ionic Radius: Smaller ions have higher lattice enthalpy because the distance between ion centers decreases, enhancing attraction.
Coordination Number: Higher coordination numbers can lead to increased lattice enthalpy due to more overlapping interactions between ions.

Calculating Lattice Enthalpy

Lattice enthalpy can be calculated using the Born-Haber cycle by rearranging the cycle's energy changes. For example, the lattice enthalpy for MgO can be determined as follows: $$ \Delta H_{formation} = \Delta H_{sublimation} + \frac{1}{2} \Delta H_{dissociation} + \Delta H_{ionization} - \Delta H_{electron affinity} - \Delta H_{lattice} $$ Rearranging for $ \Delta H_{lattice} $: $$ \Delta H_{lattice} = \Delta H_{sublimation} + \frac{1}{2} \Delta H_{dissociation} + \Delta H_{ionization} - \Delta H_{electron affinity} - \Delta H_{formation} $$ By substituting known enthalpy values, the lattice enthalpy can be accurately computed.

Importance in Predicting Compound Formation

Understanding lattice enthalpy and utilizing the Born-Haber cycle aids in predicting the feasibility and stability of ionic compounds. High lattice enthalpy often signifies a compound's strong ionic bonds and thermodynamic stability, which are essential factors in material science, mineralogy, and various industrial applications.

Limitations of Born-Haber Cycle

While the Born-Haber cycle is a powerful tool, it has certain limitations:

Assumption of Stepwise Processes: The cycle assumes that all processes occur in distinct steps, which may not accurately reflect the complexities of real chemical reactions.
Requires Accurate Data: Precise values for each enthalpy change are necessary, and inaccuracies can lead to erroneous lattice enthalpy calculations.
Applicability to Only Ionic Compounds: The Born-Haber cycle is primarily applicable to simple ionic compounds and may not be suitable for more complex or covalent compounds.

Comparison Table

Aspect	Born-Haber Cycle	Lattice Enthalpy
Definition	A thermochemical cycle that calculates the lattice enthalpy of an ionic compound using Hess's Law.	The energy released when gaseous ions form an ionic solid, indicating the strength of ionic bonds.
Purpose	To determine the lattice enthalpy indirectly by breaking down the formation of ionic compounds into measurable steps.	To quantify the stability and strength of the ionic lattice in a compound.
Application	Used in thermodynamic calculations to predict the feasibility of ionic compound formation.	Used to explain physical properties like melting point, boiling point, and solubility of ionic compounds.
Key Factors	Enthalpy of sublimation, dissociation, ionization, electron affinity, and lattice enthalpy.	Charge of ions and their ionic radii.
Limitations	Assumes stepwise processes and requires accurate data; applicable mainly to simple ionic compounds.	Does not account for covalent character in bonds and may vary with coordination number.

Summary and Key Takeaways

The Born-Haber cycle is essential for understanding the formation and stability of ionic compounds.
Lattice enthalpy quantifies the strength of ionic bonds and influences physical properties.
High lattice enthalpy correlates with higher melting points and lower solubility in ionic solids.
The Born-Haber cycle utilizes Hess's Law to calculate lattice enthalpy indirectly.
Factors such as ion charge and size significantly affect lattice enthalpy.

Examiner Tip

Tips

Use Mnemonics: Remember the Born-Haber steps with the mnemonic SIDE ELL (Sublimation, Ionization, Dissociation, Electron affinity, Lattice enthalpy).

Practice Problems: Regularly solve different Born-Haber cycle problems to reinforce your understanding and improve calculation accuracy.

Visual Aids: Create flowcharts of the Born-Haber cycle to visualize each step and its corresponding enthalpy change, aiding memory retention.

Did You Know

Historical Insight: The Born-Haber cycle was developed by the German chemist Max Born and the Austrian physicist Fritz Haber in the early 20th century to explain the formation of ionic solids.

Real-World Application: Understanding lattice enthalpy is crucial in the pharmaceutical industry for designing stable drug compounds with desired solubility properties.

Unexpected Fact: Lattice enthalpy plays a significant role in determining the hardness and melting points of minerals, influencing the mining and material science industries.

Common Mistakes

Misapplying Hess's Law: Students often forget to consider all steps in the Born-Haber cycle, leading to incorrect lattice enthalpy calculations. Incorrect: Ignoring the electron affinity step. Correct: Including all relevant enthalpy changes in the cycle.

Confusing Ionization Energy and Electron Affinity: Mixing up these terms can lead to errors in understanding energy changes. Incorrect: Assigning ionization energy values to electron affinity. Correct: Clearly distinguishing between the energy required to remove an electron and the energy released when an electron is gained.

Overlooking Ionic Sizes: Neglecting the effect of ion radii on lattice enthalpy can result in incomplete analyses. Incorrect: Ignoring how smaller ions increase lattice enthalpy. Correct: Considering both charge and ionic size when evaluating lattice enthalpy.

FAQ

What is the Born-Haber cycle?

The Born-Haber cycle is a thermochemical cycle that uses Hess's Law to calculate the lattice enthalpy of an ionic compound by breaking down its formation into a series of steps with known enthalpy changes.

Why is lattice enthalpy important?

Lattice enthalpy measures the strength of ionic bonds in a crystal lattice, influencing properties like melting point, solubility, and stability of ionic compounds.

How does ion charge affect lattice enthalpy?

Higher ion charges increase lattice enthalpy due to stronger electrostatic attractions between ions, resulting in more stable and tightly bound ionic compounds.

Can the Born-Haber cycle be used for covalent compounds?

No, the Born-Haber cycle is specifically designed for ionic compounds. Covalent compounds involve different bonding mechanisms that the cycle does not account for.

What factors influence the accuracy of lattice enthalpy calculations?

Accurate lattice enthalpy calculations depend on precise values for each step in the Born-Haber cycle, including sublimation energy, dissociation enthalpy, ionization energy, and electron affinity.

How is Coulomb's Law related to lattice enthalpy?

Coulomb's Law describes the electrostatic forces between ions, which are directly related to lattice enthalpy. It helps estimate lattice enthalpy based on ion charges and distances.