Chapter 5: Chemical Reactions Study Notes

5.1 Thermodynamics

• Definition of Energy:

  • Energy is the capacity to do work, a fundamental concept in both physics and chemistry. It manifests in various forms including kinetic, potential, thermal, and chemical energy.

  • Chemical reactions involve an exchange of energy, dissipating heat in some reactions while transforming potential energy to kinetic energy or vice versa. The energy changes occurring during reactions influence the temperature and state of the substances involved.

• Heat of Reaction:

  • Exothermic Reactions:

  • Definition: Reactions that give off heat (exo = out, thermic = heat).

  • Heat flows out into the surroundings, leading to an increase in temperature of the environment. An example is combustion.

  • Endothermic Reactions:

  • Definition: Reactions that absorb heat energy (endo = in, thermic = heat).

  • Heat flows into the system, resulting in a decrease in temperature of the surroundings. Photosynthesis is a classic example.

  • Change in Enthalpy:

  • Denoted as ΔH.

  • For exothermic reactions, ΔH is negative because the products have lower enthalpy than the reactants, indicating that energy is released.

  • For endothermic reactions, ΔH is positive because products have higher enthalpy than reactants, signifying energy absorption.

5.1.1 Randomness

• Entropy (S):

  • Definition: Measure of change in randomness or disorder in a chemical system, influencing the direction and feasibility of a reaction.

  • Changes in entropy (ΔS) can result from various factors including temperature and concentration gradients.

  • Positive ΔS indicates an increase in randomness (e.g., solid to gas), while negative ΔS indicates a decrease in randomness (e.g., gas to solid).

5.1.2 Free Energy

• Gibbs Free Energy (G):

  • Amount of energy available to do work in a system, crucial for predicting reaction spontaneity.

  • The change in free energy (ΔG) is determined by the equation:
    $<br>ext{ΔG = ΔH - TΔS}<br>$
    where T is the absolute temperature in Kelvin. A negative ΔG indicates a spontaneous reaction.

  • Exergonic Reactions:

  • Definition: Reactions that release energy (ΔG < 0) often used in biological systems to drive other reactions.

5.1.3 Activation Energy

• Definition:

  • Minimum energy required for reactants to collide properly to initiate a reaction, overcoming barriers to reaction progression.

  • Factors such as temperature, concentration, and presence of catalysts can influence the activation energy required.

5.1.4 Energy Content in Food

• Nutritional labels measure energy content in Calories, essential for dietary energy balance.

• Different nutrient molecules have different energy content:

  • Carbohydrates: 4 Cal/g, 17 kJ/g, serving as primary energy sources.

  • Protein: 4 Cal/g, 17 kJ/g, important for tissue repair and growth.

  • Fat: 9 Cal/g, 38 kJ/g, providing essential fatty acids and energy.

• Food Examples:

  • Apple: Contains approximately 21g Carbohydrate, 0.27g Protein, 0.49g Fat, giving a Total of roughly 89 Cal (370 kJ).

5.2 Chemical Reactions: Kinetics

• Rate of Reaction:

  • The rate can be measured through the amount of product formed or reactant used over a specific period, reflecting the speed of chemical transformations.

  • Factors Affecting Reaction Rate:

  • Temperature: Increased temperature raises kinetic energy, resulting in more frequent and energetic molecular collisions, thus accelerating reactions.

  • Amount of Reactant: Higher amounts of reactants increase collision frequency, speeding up the reactions due to more available particles to interact.

  • Presence of a Catalyst:

    • Catalysts lower the activation energy required for a reaction, increasing its rate without undergoing permanent changes themselves.

    • Common catalysts include acids, bases, and metal ions, which can stabilize transition states or improve reactant orientation.

5.2.1 Enzymes

• Definition:

  • Biological catalysts, predominantly proteins, that greatly enhance reaction rates, often by over 10 million times.

  • Enzymes accelerate reactions by attracting and immobilizing reactants at an active site, facilitating the correct orientation necessary for the reaction to occur efficiently.

  • Specificity:

  • Enzymes are highly specific, catalyzing only particular reactions. Defects in enzymes can lead to diseases; for example, tyrosinase deficiency causes albinism due to disrupted melanin production.

5.3 Overview of Chemical Reactions

• Types of Reactions:

  • Synthesis:

  • Definition: Smaller molecules combine to form larger ones. General scheme:
    $<br>A + B <br>ightarrow AB<br>$

  • Decomposition:

  • Definition: Larger molecules break down into smaller ones. General scheme:
    $<br>AB <br>ightarrow A + B<br>$

  • Exchange (Displacement):

  • Involves both synthesis and decomposition (e.g., single and double displacements), relevant in various biological and chemical processes.

  • Reversible vs. Irreversible:

  • Reversible reactions can occur in both directions (e.g.,
    $<br>A + B <br>ightarrow AB <br>ewline AB <br>ightarrow A + B<br>$).

  • Irreversible reactions generally move only in one direction, such as combustion, where reactants cannot reform.

  • Combustion:

  • Organic molecule reacts with oxygen, producing carbon dioxide and water in exothermic reactions. These reactions release large amounts of energy, important for biological energy production.

5.3.1 Chemical Equations

• Organic chemical reactions depict the structure of molecules involved, written with specific connectivity to understand reacting parties.

• Reaction conditions are often displayed (e.g., temperature, catalysts), influencing the overall course of chemical reactions. Coupled reactions may show energy transfer as observed in biochemical pathways such as cellular respiration.

5.4 Oxidation and Reduction

• Redox Reactions:

  • Essential for energy production; oxidation and reduction involve electron transfer crucial in metabolic processes.

  • Inorganic Redox:

  • Involves metal ions; oxidation is the loss of electrons, while reduction is the gain of electrons. This electron flow is vital in various industrial processes.

  • Mnemonic: "OIL RIG" (Oxidation Is Loss, Reduction Is Gain) helps in remembering these concepts.

  • Organic Redox:

  • Gain of oxygen or loss of hydrogen characterizes oxidation; conversely, gain of hydrogen or loss of oxygen signifies reduction. These reactions are crucial in organic synthesis and energy mobilization.

5.4.1 Biological Oxidation and Reduction

• Energy transfer occurs via oxidation reactions of nutrients, e.g., the combustion of glucose during cellular respiration, providing energy for cellular activities.

• Important coenzymes in redox reactions include NAD+ and FAD, which facilitate electron transport in metabolic processes.

5.5 Organic Reactions: Condensation and Hydrolysis

• Condensation Reactions:

  • Two organic molecules combine to form a larger molecule, producing water as a byproduct, crucial in forming complex carbohydrates and proteins.

• Hydrolysis Reactions:

  • Water is consumed to break a molecule into two smaller molecules, reversing condensation reactions and essential for digestion and metabolism.

5.6 Organic Addition Reactions to Alkenes

• Addition of atoms to a double bond in an alkene involves breaking the double bond to form single bonds, a key reaction type in organic synthesis.

• Hydrogenation:

  • Adding H₂ to convert alkene to alkane using catalysts like Pt, Ni, or Pd, a process important in food chemistry; however, it can result in trans fats via partial hydrogenation.

• Hydration:

  • Addition of water to an alkene, typically requiring a catalyst (often an acid or enzyme), forming alcohols, a significant reaction in organic chemistry that expands the utility of alkenes.