Thermochemistry

Energy and Internal Energy

Thermochemistry is the study of heat released or absorbed during chemical reactions. The internal energy of a system, denoted as E, is the sum of all kinetic and potential energies of its components.

• Change in Internal Energy:

• Heat (q): Energy transferred due to temperature difference between system and surroundings.

• Work (w): Energy transferred when a force moves an object.

• Relationship:

Types of Energy

• Kinetic Energy: Energy of motion (e.g., thermal energy from molecular movement).

• Potential Energy: Energy due to position or composition (e.g., chemical energy in bonds).

Exothermic and Endothermic Processes

• Exothermic: Heat is released by the system ().

• Endothermic: Heat is absorbed by the system ().

Units of Energy

• Joule (J): SI unit of energy.

• Calorie (cal): Amount of energy required to raise 1 g of water by 1°C.

• Nutritional Calorie: 1 Cal = 1000 cal

State Functions

State functions depend only on the current state, not the path taken. Examples: internal energy, enthalpy.

• Pathway Independence: The value is the same regardless of the process.

First Law of Thermodynamics

• Law: Energy can neither be created nor destroyed; it can only be transferred or transformed.

• Equation:

Enthalpy

Enthalpy (H) is the heat content of a substance. The change in enthalpy () is the heat given off or absorbed at constant pressure.

• Enthalpy of Reaction:

• Exothermic: (heat released)

• Endothermic: (heat absorbed)

Calorimetry

Calorimetry measures the heat transferred in a reaction or process.

• Heat Capacity (C): Amount of heat required to raise the temperature of a substance by 1°C.

• Specific Heat (c): Amount of heat required to raise 1 g of a substance by 1°C.

• Formula:

• Example: To heat 35.0 g of Al from 55.0°C to 75.0°C with :

Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the sum of the enthalpy changes for individual steps.

• Equation:

• Manipulation: Reactions can be added, multiplied, or reversed to find overall .

Standard Enthalpy of Formation

• Definition: Enthalpy change for forming 1 mole of a compound from its elements in their standard states.

• Formula:

Atomic Structure and Periodic Properties

Electron Configuration and Orbital Diagrams

Electron configuration describes the arrangement of electrons in an atom. Orbital diagrams visually represent electron placement in orbitals.

• Subshells: s, p, d, f

• Energy Levels: 7 principal energy levels

• Valence Electrons: Electrons in the outermost shell, involved in bonding.

Sublevel Energy Splitting and Shielding

• Shielding: Inner electrons reduce the effective nuclear charge felt by valence electrons.

• Effective Nuclear Charge (): The net positive charge experienced by valence electrons.

Periodic Trends

• Atomic Radius: Increases down a group, decreases across a period.

• Ion Size: Cations are smaller, anions are larger than their parent atoms.

• Ionization Energy: Energy required to remove an electron. Increases across a period, decreases down a group.

• Electron Affinity: Energy change when an electron is added. Increases across a period, decreases down a group.

• Metallic Character: Increases down a group, decreases across a period.

Isoelectronic Particles

• Definition: Species with the same number of electrons.

• Example: Ne, Na+, Mg2+ are isoelectronic.

Chemical Bonding

Types of Bonds

• Ionic Bond: Attraction between cations and anions via electrostatic force.

• Covalent Bond: Sharing of electron pairs between atoms.

• Metallic Bond: "Sea" of electrons shared among metal atoms.

• Macromolecular Crystals: Large networks of covalently bonded atoms.

Valence Electrons and Lewis Structures

• Valence Electrons: Outermost electrons involved in bonding.

• Lewis Structures: Electron dot diagrams showing bonding and lone pairs.

• Steps: Count valence electrons, choose central atom, arrange atoms, assign bonds and lone pairs, check octet rule.

Exceptions to the Octet Rule

• Electron Deficient: Central atom has less than 8 electrons (e.g., BCl3).

• Expanded Octet: Central atom has more than 8 electrons (period 3 or higher, e.g., SF6).

• Odd Number of Electrons: Molecules with unpaired electrons (e.g., NO2).

Electronegativity and Bond Polarity

• Electronegativity: Atom's ability to attract shared electrons. Increases left to right, bottom to top.

• Bond Polarity: Difference in electronegativity determines bond type:

  • Nonpolar Covalent:

  • Polar Covalent:

  • Ionic:

States of Matter and Intermolecular Forces

Properties of Gases, Liquids, and Solids

• Gases: Far apart particles, low density, compressible, indefinite shape.

• Liquids: Close particles, high density, not compressible, indefinite shape.

• Solids: Very close particles, high density, not compressible, definite shape.

Intermolecular Forces

• Dispersion (London) Forces: Weak attractions between nonpolar molecules.

• Dipole-Dipole: Attractions between polar molecules.

• Hydrogen Bonds: Strong dipole-dipole between H and F, O, or N.

• Ion-Dipole: Attraction between ions and polar molecules (e.g., salt in water).

Surface Tension, Viscosity, and Capillary Action

• Surface Tension: Intermolecular forces at the surface create tension.

• Viscosity: Resistance to flow; increases with stronger intermolecular forces.

• Capillary Action: Movement of liquid in narrow spaces due to adhesion and cohesion.

Evaporation and Vapor Pressure

• Evaporation: Liquid to gas; rate increases with temperature and surface area.

• Vapor Pressure: Pressure exerted by vapor above a liquid at equilibrium.

• Boiling Point: Temperature at which vapor pressure equals atmospheric pressure.

Changes of State

• Specific Heat Capacity (c): Heat required to raise 1 g of substance by 1°C.

• Heat of Fusion (): Heat needed to convert solid to liquid.

• Heat of Vaporization (): Heat needed to convert liquid to gas.

• Formula:

Heating Curve and Phase Changes

• Heating Curve: Shows temperature changes and phase transitions as heat is added.

• Phase Changes: Melting, freezing, vaporization, condensation, sublimation, deposition.

Phase Diagrams

A phase diagram plots pressure vs. temperature, showing regions of stability for solid, liquid, and gas.

• Vapor Pressure Curve: Separates gas and liquid.

• Fusion Curve: Separates solid and liquid.

• Sublimation Curve: Separates solid and gas.

• Triple Point: All three phases coexist.

• Critical Point: Above this, only supercritical fluid exists.

Appendix: Periodic Table

• Periodic Table: Fundamental tool for organizing elements by atomic number, electron configuration, and recurring chemical properties.

*Additional info: The periodic table image is included for reference and practice in identifying element positions and periodic trends.*