Chapter 5: Thermodynamics & Kinetics – Chemical Reactions and Energy in GOB Chemistry

Thermodynamics

Gibbs Free Energy and Spontaneity

Thermodynamics studies the energy changes that accompany chemical reactions. The concept of Gibbs free energy ($\Delta G$) is central to predicting whether a reaction will occur spontaneously.

• Gibbs Free Energy Equation: $\Delta G = \Delta H - T\Delta S$ Where $\Delta H$ is enthalpy (heat change), $T$ is temperature in Kelvin, and $\Delta S$ is entropy (randomness).

• Exergonic reactions: Give off energy; $\Delta G$ is negative. These reactions do not require energy input and occur spontaneously.

• Endergonic reactions: Require energy; $\Delta G$ is positive. These reactions are nonspontaneous and need energy input from the surroundings.

• Spontaneous processes: Continue to occur once started and do not require energy from the surroundings.

Example: Cellular respiration is an exergonic process, releasing energy for cellular work.

Visualizing Exergonic and Endergonic Reactions

Energy diagrams help illustrate the difference between exergonic and endergonic reactions.

• In exergonic reactions, the free energy of reactants is higher than that of products ($\Delta G$ is negative).

• In endergonic reactions, the free energy of products is higher than that of reactants ($\Delta G$ is positive).

Example: Photosynthesis is endergonic, requiring energy input from sunlight.

Coupling of Reactions in Living Systems

Living systems often couple exergonic and endergonic reactions to drive necessary processes.

• Free energy released from one reaction can be used to power another reaction that requires energy.

• Adenosine triphosphate (ATP) is the main molecule that transfers energy in cells.

Example: ATP hydrolysis (exergonic) provides energy for muscle contraction (endergonic).

Activation Energy

Even spontaneous reactions may require an initial input of energy, called activation energy, to get started.

• Reactants must collide with sufficient energy and correct orientation for a reaction to occur.

• Activation energy is the energy barrier that must be overcome for reactants to transform into products.

Example: Lighting a match provides activation energy for combustion.

Energy Content in Food

The energy stored in food is released during combustion (reaction with oxygen), producing carbon dioxide, water, and energy.

• Different food molecules contain different amounts of energy.

• Energy is measured in Calories (Cal) or kilojoules (kJ).

Food Molecule	Energy Example	Energy Cal/g	Energy kJ/g
Carbohydrate	Table sugar, potatoes, flour	4	17
Protein	Meats, fish, beans	4	17
Fat	Oil, butter	9	38

Example: Fats provide more than twice the energy per gram compared to carbohydrates or proteins.

Caloric Values of Common Foods

Foods vary in their carbohydrate, protein, and fat content, affecting their total energy value.

Food	Carbohydrate (g)	Protein (g)	Fat (g)	Calories (kJ)
Apple, 1 medium	21	0.27	0.49	89 (372)
Egg white, 1	0.008	2.4	0.9	74 (310)
Peanut butter, 1 Tbsp	3.55	3.98	7.80	100 (420)
Chicken, 1 oz	0	19	27	320 (1300)
Orange juice, 1 cup	27	2	0	116 (484)
Brown rice, 1/2 cup	24.8	2.45	0.78	116 (484)

Example: Peanut butter is high in fat and protein, contributing to its high caloric value.

Chemical Reactions: Kinetics

Reaction Rate

Kinetics studies the speed at which chemical reactions occur and the factors that affect this rate.

• Reaction rate: The amount of product formed (or reactant used up) per unit time.

• Reactant molecules must collide with enough energy and proper orientation for a reaction to occur.

• Factors affecting rate: temperature, concentration of reactants, and presence of a catalyst.

Example: Increasing temperature speeds up the rate of most chemical reactions.

Effect of Temperature

• Increasing temperature increases the kinetic energy of molecules, leading to more frequent and energetic collisions.

• Higher temperature generally increases reaction rate.

Example: Food spoils faster at room temperature than in a refrigerator.

Effect of Reactant Concentration

• Higher concentration of reactants increases the likelihood of collisions, speeding up the reaction.

• As reactants are consumed, the reaction slows down.

Example: More concentrated acid reacts faster with metal than dilute acid.

Role of Catalysts

• Catalyst: A substance that speeds up a reaction by lowering the activation energy, without being consumed.

• Catalysts are not reactants or products and can be reused.

• Acids, bases, and metal ions often act as catalysts.

• Catalysts may immobilize reactants, increasing collision likelihood.

Example: Enzymes are biological catalysts that speed up biochemical reactions.

Enzymes: Biological Catalysts

• Most enzymes are proteins that enhance reaction rates by factors of millions.

• Enzymes immobilize reactants at the active site, orienting them for reaction.

• Enzymes are highly specific; defects can cause disease.

Example: Lactase deficiency leads to lactose intolerance.

Enzyme	Deficiency Symptoms
Tyrosinase	Absence of pigment in skin, hair, eyes
Lactase	Diarrhea, bloating, abdominal cramping
Phenylalanine hydrolase (PKU)	Neurologic symptoms, mental retardation

Overview of Chemical Reactions

Types of Chemical Reactions

Chemical reactions can be classified by how reactants and products change.

• Combination (Synthesis) reactions: Two or more reactants form one product. $A + B \rightarrow AB$

• Decomposition reactions: One reactant breaks into two or more products. $AB \rightarrow A + B$

• Exchange (Displacement) reactions: Reactants swap components to form new products. $AB + CD \rightarrow AD + CB$

Example: Formation of water: $2H_2 + O_2 \rightarrow 2H_2O$ (combination)

Reversible and Irreversible Reactions

• Reversible reactions: Can proceed in both directions; indicated by double arrows ($\rightleftharpoons$).

• Irreversible reactions: Proceed only in one direction, often highly exothermic.

• Chemical equilibrium: The point at which the rates of forward and reverse reactions are equal.

Example: $A + B \rightleftharpoons AB$

Combustion Reactions

• Combustion involves an organic molecule reacting with oxygen to produce carbon dioxide and water.

• Combustion reactions are highly exothermic and considered irreversible.

• They are a type of decomposition and oxidation-reduction reaction.

Example: Burning methane: $CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O$

Writing Chemical Reactions

• Organic reactions show the structure of molecules and reaction conditions.

• Biochemical reactions may be coupled, with energy transfer indicated.

• Notation may include chemical formulas, drawn structures, and energy changes.

Example: ATP hydrolysis: $ATP \rightarrow ADP + P_i$ (energy released)

Summary Table: Types of Chemical Reactions

Reaction Type	General Reaction
Combination	$A + B \rightarrow AB$
Decomposition	$AB \rightarrow A + B$
Exchange	$AB + CD \rightarrow AD + CB$

Additional info: These notes expand on brief points with academic context, definitions, and examples for clarity and completeness.