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

Common chemical reactions in the laboratory

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Common Chemical Reactions in the Laboratory

Introduction

Chemical reactions are fundamental to laboratory experiments in the IB Chemistry SL curriculum. Understanding common chemical reactions allows students to grasp essential concepts, predict outcomes, and safely conduct practical work. This article explores various types of chemical reactions typically encountered in the laboratory, highlighting their significance and applications within the IB framework.

Key Concepts

Synthesis Reactions

Synthesis reactions, also known as combination reactions, involve the combination of two or more reactants to form a single product. These reactions are crucial for synthesizing compounds from simpler substances.

**General Equation:** $$A + B \rightarrow AB$$

**Example:** The formation of water from hydrogen and oxygen gases is a classic synthesis reaction:

$$2H_2 + O_2 \rightarrow 2H_2O$$

**Applications:** Synthesis reactions are employed in the production of various chemicals, such as ammonia in the Haber process:

$$3H_2 + N_2 \rightarrow 2NH_3$$

**Key Points:**

Reactants combine to form a new compound.
Exothermic or endothermic depending on the reaction.
Important for industrial chemical production.

Decomposition Reactions

Decomposition reactions involve breaking down a compound into two or more simpler substances. These reactions are essential for understanding compound stability and reactivity.

**General Equation:** $$AB \rightarrow A + B$$

**Example:** Decomposition of hydrogen peroxide into water and oxygen:

$$2H_2O_2 \rightarrow 2H_2O + O_2$$

**Applications:** Decomposition reactions are used in the preparation of oxygen and other gases in the laboratory.

**Key Points:**

A single compound breaks down into simpler substances.
Often requires energy input (heat, light).
Used in recycling and waste management processes.

Single Replacement Reactions

Single replacement reactions occur when an element reacts with a compound, displacing another element from it. These reactions are important for understanding reactivity series and displacement.

**General Equation:** $$A + BC \rightarrow AC + B$$

**Example:** Zinc reacting with hydrochloric acid to produce zinc chloride and hydrogen gas:

$$Zn + 2HCl \rightarrow ZnCl_2 + H_2$$

**Applications:** Single replacement reactions are utilized in metal extraction and purification processes.

**Key Points:**

An element replaces another in a compound.
Depends on the reactivity of the elements involved.
Crucial for metallurgy and chemical manufacturing.

Double Replacement Reactions

Double replacement reactions involve the exchange of ions between two compounds, resulting in the formation of two new compounds. These reactions are key to understanding precipitation, neutralization, and gas formation.

**General Equation:** $$AB + CD \rightarrow AD + CB$$

**Example:** Reaction between silver nitrate and sodium chloride to form silver chloride and sodium nitrate:

$$AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$$

**Applications:** Double replacement reactions are fundamental in water purification and wastewater treatment processes.

**Key Points:**

Exchange of ions between two compounds.
Can result in precipitation, gas formation, or neutralization.
Widely used in analytical chemistry for qualitative analysis.

Combustion Reactions

Combustion reactions involve the rapid reaction of a substance with oxygen, releasing energy in the form of heat and light. These exothermic reactions are essential for energy generation and various industrial processes.

**General Equation for Complete Combustion of Hydrocarbons:** $$C_xH_y + O_2 \rightarrow CO_2 + H_2O$$

**Example:** Combustion of methane:

$$CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$$

**Applications:** Combustion reactions are the basis for engines, power generation, and heating systems.

**Key Points:**

Requires oxygen and releases energy.
Produces carbon dioxide and water in complete combustion.
Incomplete combustion can produce carbon monoxide and soot.

Comparison Table

Reaction Type	General Equation	Key Characteristics
Synthesis	$$A + B \rightarrow AB$$	Combines reactants to form a single product
Decomposition	$$AB \rightarrow A + B$$	Breaks down a compound into simpler substances
Single Replacement	$$A + BC \rightarrow AC + B$$	An element displaces another in a compound
Double Replacement	$$AB + CD \rightarrow AD + CB$$	Exchange of ions between two compounds
Combustion	$$C_xH_y + O_2 \rightarrow CO_2 + H_2O$$	Rapid reaction with oxygen, releases energy

Summary and Key Takeaways

Understanding common laboratory chemical reactions is essential for IB Chemistry SL students.
Synthesis and decomposition reactions involve combination and breakdown of compounds, respectively.
Single and double replacement reactions focus on element displacement and ion exchange.
Combustion reactions are critical for energy-related applications.
Mastery of these reactions aids in predicting outcomes and conducting safe experiments.

Examiner Tip

Tips

Use the mnemonic **SCDS** to remember reaction types:
Synthesis, Combustion, Decomposition, Single replacement, and Double replacement.

Balance chemical equations step-by-step, starting with elements that appear once on each side.

Understand the reactivity series to predict outcomes in single replacement reactions effectively.

Did You Know

1. The Haber process, a synthesis reaction, produces over 150 million tons of ammonia annually, essential for fertilizers worldwide.

2. Decomposition reactions like the breakdown of hydrogen peroxide are used in eco-friendly rocket propellants.

3. Combustion reactions not only power engines but also played a pivotal role in the development of early industrial societies.

Common Mistakes

Incorrect: Balancing the equation $H_2 + O_2 \rightarrow H_2O$ as $H_2 + O_2 \rightarrow H_2O$.
Correct: $$2H_2 + O_2 \rightarrow 2H_2O$$

Incorrect: Forgetting to account for all products in a double replacement reaction.
Correct: $$AgNO_3 + NaCl \rightarrow AgCl + NaNO_3$$

Incorrect: Assuming all combustion reactions produce only CO₂ and H₂O.
Correct: Incomplete combustion can produce CO and soot.

FAQ

What is the difference between synthesis and decomposition reactions?

Synthesis reactions involve combining two or more reactants to form a single product, whereas decomposition reactions break down a single compound into two or more simpler substances.

How can you identify a double replacement reaction?

Double replacement reactions involve the exchange of ions between two compounds, typically resulting in the formation of a precipitate, gas, or water.

Why are combustion reactions important in everyday life?

Combustion reactions are essential for energy production in engines, heating systems, and power plants, making them integral to transportation and industrial processes.

What factors affect the rate of a chemical reaction in the laboratory?

Factors include temperature, concentration of reactants, surface area, and the presence of catalysts, all of which can influence the speed and efficiency of reactions.

How do single replacement reactions relate to the reactivity series?

The reactivity series ranks elements based on their ability to displace others in single replacement reactions. A more reactive element can replace a less reactive one from its compound.

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