Chemical Reactions and Stoichiometry

Law of Conservation of Matter

The law of conservation of matter states that atoms are neither created nor destroyed in a chemical reaction. This principle requires that the number of atoms of each element be the same on both sides of a chemical equation.

• Balanced Chemical Equation: Uses coefficients to ensure equal numbers of each atom on both sides.

• Reactants: Substances that undergo change.

• Products: Substances formed as a result of the reaction.

• Example:

Important: Do not balance equations by changing subscripts, as this alters the identity of the compound.

The Mole and Avogadro's Number

A mole (mol) is a unit representing 6.022 x 1023 particles (atoms or molecules), known as Avogadro's number. It allows chemists to relate atomic scale quantities to measurable amounts.

• 1 mole = 6.022 x 1023 atoms or molecules

• Example: 1 mole of NaCl = 58.4 g, 1 mole of Cu = 63.55 g

• Key Point: The mass of 1 mole varies by substance, but the number of particles is always the same.

Stoichiometry

Stoichiometry uses the coefficients in balanced equations to determine the ratios of reactants and products, usually in moles.

• Stoichiometric Ratio: Derived from coefficients in the equation.

• Example: means 2 mol H2 react with 1 mol O2 to produce 2 mol H2O.

• Mass Relationships: Can convert moles to grams using molecular weights.

Redox Reactions

Oxidation and Reduction

Redox reactions involve the transfer of electrons between substances. Oxidation is the loss of electrons, while reduction is the gain of electrons.

• Oxidation: Loss of electrons, increase in oxidation number.

• Reduction: Gain of electrons, decrease in oxidation number.

• Mnemonic: LEO says GER (Lose Electrons Oxidation, Gain Electrons Reduction); OIL RIG (Oxidation Is Loss, Reduction Is Gain).

Oxidation Numbers

Oxidation numbers are assigned to atoms to track electron transfer in reactions.

• Rules:

  • Uncombined atom or homoatomic molecule: 0

  • Monatomic ion: equals its charge

  • Group IA: +1, Group IIA: +2

  • H: +1, O: -2 (except peroxides)

  • Sum of oxidation numbers equals overall charge

• Example: In CO2, C = +4, O = -2

Redox in Metabolism

Metabolic reactions often involve oxidation of carbon compounds, releasing energy.

• Example: Combustion of glucose:

Enthalpy of Reactions

Enthalpy (ΔH)

Enthalpy is the thermal energy released or absorbed in a reaction per mole.

• Exothermic Reaction: Releases energy, ΔH is negative.

• Endothermic Reaction: Absorbs energy, ΔH is positive.

• Formula:

Comparison Table: Exothermic vs. Endothermic Reactions

Spontaneity and Thermodynamics

Free Energy (ΔG)

Free energy determines whether a process is spontaneous. It is the energy available to do work.

• Formula:

• Exergonic: Releases free energy, ΔG is negative, spontaneous.

• Endergonic: Requires free energy input, ΔG is positive, non-spontaneous.

Entropy (ΔS)

Entropy is a measure of disorder in a system. Processes that increase disorder have positive ΔS.

• Examples: Solid → liquid → gas, breaking apart molecules, increasing number of gas molecules.

Relationship Between ΔG, ΔH, and ΔS

• Formula:

• Negative ΔH and positive ΔS favor spontaneity.

• Temperature (T) increases the impact of entropy.

Reaction Kinetics

Activation Energy and Reaction Rate

Activation energy (Eact) is the energy required to start a reaction. Reaction rate measures how quickly reactants are converted to products.

• Factors Affecting Rate:

  • Temperature: Higher temperature increases rate.

  • Concentration: Higher concentration increases rate.

  • Catalyst: Lowers activation energy, increases rate, unchanged by reaction.

  • Inhibitor: Slows down reaction rate.

Equilibrium

Equilibrium and Reversible Reactions

Equilibrium occurs when the rate of the forward reaction equals the rate of the reverse reaction. Reversible reactions can proceed in either direction.

• Equilibrium Constant (Keq): Ratio of product concentrations to reactant concentrations at equilibrium.

• Formula:

• Keq > 1: Forward reaction favored; Keq < 1: Reverse reaction favored.

Le Châtelier’s Principle

When a system at equilibrium is disturbed, it shifts to relieve the stress and restore equilibrium.

• Changing Concentration: Adding reactant or removing product shifts equilibrium to the right (forward).

• Changing Temperature:

  • Exothermic: Adding heat shifts equilibrium to the left (reverse).

  • Endothermic: Adding heat shifts equilibrium to the right (forward).

• Adding Catalyst: Speeds up reaching equilibrium but does not change Keq.

Summary Table: Key Definitions