Energy Changes and Reactions in Chemistry

Chemistry Study Notes

Energy Changes

Exothermic Reactions

  • Definition: An exothermic reaction is characterized by energy being transferred to the surroundings, typically in the form of heat. This is evidenced by a rise in temperature of the surrounding environment.

  • Examples of Exothermic Reactions:

    • Combustion: The process of burning substances to release energy, often producing light and heat.

    • Neutralisation: A reaction wherein an acid and a base react to form water and a salt, releasing energy.

    • Respiration: The biochemical process in which cells convert glucose and oxygen into energy, carbon dioxide, and water.

  • Everyday Applications of Exothermic Reactions:

    • Hand warmers: Utilizing the oxidation of iron to produce warmth.

    • Self-heating cans: Employing exothermic reactions in bases to generate heat.

Endothermic Reactions

  • Definition: An endothermic reaction absorbs energy from the surroundings, resulting in a decrease in temperature of the surrounding environment.

  • Examples of Endothermic Reactions:

    • Photosynthesis: The process by which green plants use sunlight to convert carbon dioxide and water into glucose and oxygen, requiring energy.

    • Thermal Decomposition: A reaction where a compound is broken down into simpler products with the absorption of heat.

  • Everyday Applications of Endothermic Reactions:

    • Sports injury packs: These packs are designed to cool the affected area rapidly, often by using a reaction involving citric acid and sodium hydrogen carbonate.

Practical Demonstration of Chemical Reactions
Example Experiment with Copper Sulfate and Zinc

  1. Set up equipment as depicted in a specific diagram.

  2. Measure the initial temperature of the solution (copper sulfate).

  3. Add the zinc powder to the solution.

  4. Measure the resulting highest or lowest temperatures observed after the reaction.

Energy Profile Diagrams

Components of Energy Profiles

  • Key Terminology:

    • R = Reactants: The starting materials in a chemical reaction.

    • P = Products: The substances formed from the chemical reaction.

    • Ea = Activation Energy: The minimum energy required to initiate a reaction.

    • ΔE = Energy Change: The overall change in energy from reactants to products.

  • Role of Catalysts:

    • A catalyst is a substance that increases the rate of a specific chemical reaction without being consumed in the reaction. It lowers the activation energy (denoted as Ea_cat) by providing an alternative reaction pathway.

Bond Energies and Their Implications

  • Breaking Bonds: Requires energy (endothermic process).

  • Making Bonds: Releases energy (exothermic process).

  • Example Calculation of Bond Energies:

    • For the reaction: 2H + Br_2 \rightarrow 2HBr

    • Energy required to break bonds: 436 kJ/mol (for H-H bonds) + 2(366 kJ/mol) (for two H-Br bonds) + 196 kJ/mol (for Br-Br bond).

    • Energy released when forming bonds: Total bonds formed in products correlating to energy release.

    • Example energy analysis: 629 \text{ kJ/mol (reactants)} - 732 \text{ kJ/mol (products)} = -103 \text{ kJ/mol} which signifies an exothermic reaction due to the higher energy release during bond formation.

Energy Changes - Required Practical Investigations

Practical #1: Temperature Change in Neutralization Reactions

  • Aim: To investigate temperature change in neutralization reactions between hydrochloric acid (HCl) and sodium hydroxide (NaOH).

  • Chemical Reaction:

    • NaOH + HCl \rightarrow NaCl + hago

  • Methodology:

  1. Ensure goggles are worn at all times for safety.

  2. Set up the equipment as per provided diagrams (using polystyrene cups for insulation).

  3. Record the initial temperature of the sodium hydroxide solution.

  4. Measure 20 cm³ of 1 mol/dm³ hydrochloric acid and add it to the sodium hydroxide solution.

  5. Stir the mixture and record the highest temperature reached.

  6. Calculate the temperature change observed.

  7. Repeat this experiment using different concentrations of HCl.

Practical #2: Investigating the Influence of Concentrations

  • Similar methodology as Practical #1 but includes varying volumes and concentrations of sodium hydroxide and hydrochloric acid to analyze the effect on temperature change.

Cells and Batteries

Definition and Functionality of Cells

  • A cell is a fundamental component that consists of two different metals (electrodes) immersed in an electrolyte, enabling the production of electricity.

  • Electrolyte: A substance in liquid form that facilitates the flow of electric current due to the movement of ions.

  • When the cell is connected to a circuit, a potential difference (PD) is established between the electrodes, allowing a charge to flow and producing electricity.

Factors Affecting Cell Voltage

  1. Reactivity of Electrodes: The larger the difference in reactivity between the two metals used as electrodes, the greater the voltage produced.

  2. Electrolytes Used: Various electrolytes yield differing ion reactions that can affect voltage.

  3. A battery is formed when two or more cells are connected in series, increasing the overall voltage output.

Types of Batteries

  • Non-Rechargeable Batteries:

    • Chemical reactions within these cells become irreversible once reactants are depleted, leading to a stop in electricity production.

    • Example: Alkaline batteries.

  • Rechargeable Batteries:

    • These operate reversibly; the chemical reactions can be reversed by supplying an external electrical current.

    • Example: Lithium-ion batteries used in devices such as phones and laptops.

Hydrogen Fuel Cells

Overview of Fuel Cells

  • A fuel cell is an electrochemical cell that converts chemical energy from a fuel into electrical energy.

  • Oxygen and Hydrogen Fuel: The primary components needed are hydrogen (as the fuel) and oxygen (from the air).

  • Electrolysis Reactions: Involves oxidation (hydrogen) and reduction (oxygen) processes across electrodes.

Electrochemical Reactions in Fuel Cells

  1. Oxidation of Hydrogen (Anode - negative electrode):

    • 2H_2 \rightarrow 4H^+ + 4e^-

  2. Reduction of Oxygen (Cathode - positive electrode):

    • 4H^+ + O2 + 4e^- \rightarrow 2H2O

  3. Overall Reaction:

    • 2H2 + O2 \rightarrow 2H_2O where water is the only byproduct.

Advantages and Disadvantages of Hydrogen Fuel Cells

  • Advantages:

    • Produces only water as waste, leading to environmental benefits.

    • Longer operational life compared to conventional batteries and fewer disposal issues.

  • Disadvantages:

    • Hydrogen is a gas, necessitating larger storage space and presenting explosive risks.

    • Economic cost and the requirement for constant supply pose logistical challenges.

Comparing Cells

Key Characteristics and Applications

  1. Non-rechargeable Cells:

    • Example: Alkaline batteries used in remotes and toys.

    • Advantages: More affordable and portable.

    • Disadvantages: Once depleted, they cannot be reused and are toxic post-use.

  2. Rechargeable Cells:

    • Example: Lithium-ion batteries in phones and laptops.

    • Advantages: Can be reused multiple times, more economical than fuel cells in the long run.

    • Disadvantages: More expensive and requires time to recharge.

  3. Fuel Cells:

    • Example: Hydrogen fuel cells used in vehicles like cars and spacecraft.

    • Advantages: Only water vapor as a waste product, continuous energy supplied as long as there is a constant fuel source.

    • Disadvantages: Higher cost and safety concerns around flammability and storage of hydrogen gas.