Chem 105 middterm three introduction to heat (lecture 1)

High Chemical Potential Energy

Indicates high reactivity; the substance is unstable and has a strong potential to react.

Low Chemical Potential Energy

Indicates high stability; the substance is unreactive or less prone to chemical changes.

Bond Breaking

An endothermic process that always absorbs energy from the surroundings.

Bond Formation

An exothermic process that always releases energy into the surroundings.

First Law of Thermodynamics

Energy cannot be created or destroyed in an isolated system. ΔE_system + ΔE_surroundings = 0.

System vs. Surroundings Relationship

ΔE_system = -ΔE_surroundings. Any energy lost by the system is gained by the surroundings, and vice versa.

Positive ΔE (+ΔE)

Means the system has gained internal energy from its surroundings.

Negative ΔE (-ΔE)

Means the system has lost internal energy to its surroundings.

State Function

A property whose value depends only on the initial and final states of the system, not the path taken between them (e.g., ΔE, ΔH).

Two Modes of Energy Transfer

Heat (q) and Work (w). Chemical energy can only be transferred through these two mechanisms.

Condition for Chemical Work

A chemical reaction performs pressure-volume work only when gases are produced or consumed.

Physics Definition of Work

w = F × D (Work = Force multiplied by Distance).

Pressure-Volume Work Equation

w = -PΔV (Work equals negative external pressure multiplied by change in volume).

Internal Energy Change Equation

ΔE = q + w (where ΔE is total internal energy change, q is heat, and w is work).

Specific Heat Capacity Equation

q = m c ΔT (q = heat energy in Joules, m = mass in grams, c = specific heat, ΔT = change in temperature).

Translational Motion

Kinetic motion where an entire molecule moves from one location in space to another.

Rotational Motion

Kinetic motion where a molecule rotates/spins around an axis.

Vibrational Motion

Kinetic motion where bond lengths and angles oscillate as atoms move toward and away from each other.

Exothermic Reaction

A reaction that releases heat to the surroundings (ΔH < 0). System energy decreases; surroundings feel warmer.

Endothermic Reaction

A reaction that absorbs heat from the surroundings (ΔH > 0). System energy increases; surroundings feel colder.

Endothermic Phase Changes

Phase transitions that require absorbing heat: Melting (Fusion), Vaporization, and Sublimation.

Exothermic Phase Changes

Phase transitions that release heat: Freezing, Condensation, and Deposition.

Phase Change Heat Equation

q = ΔH × n (where q is total heat absorbed/released, ΔH is molar enthalpy of phase change, and n is moles).

Relationship of ΔE and ΔH without Gases

When no gases are involved in a reaction, ΔV ≈ 0 so w ≈ 0, making ΔE ≈ ΔH.

Fusion

The phase change from solid to liquid. Requires absorption of the enthalpy of fusion (ΔH_fus).

Vaporization

The phase change from liquid to gas. Requires absorption of the enthalpy of vaporization (ΔH_vap).

Sublimation

The direct phase transition from solid to gas without passing through the liquid phase.