calorimetry

Thermochemistry

The study of energy changes that accompany physical or chemical changes in matter.

Physical Change

A change affecting the form/state of a substance without altering its chemical composition (e.g., cutting copper, phase changes like ice melting).

Chemical Change

A change in the chemical composition of a substance that produces one or more new substances through reactions (e.g., copper tarnishing).

Energy

The ability to do work, measured in the SI unit Joules (J).

Work

The amount of energy transferred by a force over a distance.

Potential Energy

Energy due to the position or composition of an object (e.g., a ball at the top of a hill, a stretched rubber band).

Kinetic Energy

The energy of motion.

Thermal Energy

The total quantity of kinetic and potential energy in a substance.

Heat

The transfer of thermal energy from a warmer object to a cooler object.

Temperature

A measure of the average kinetic energy of the entities/particles in a substance.

Law of Conservation of Energy (First Law of Thermodynamics)

Energy cannot be created or destroyed; it can only be converted from one form to another during physical or chemical changes.

Chemical System

The specific molecules/substances undergoing a physical change (e.g., melting) or a chemical change (reaction).

Surroundings

All the matter that is not part of the system (e.g., the beaker, thermometer, air).

Open System

A system that can exchange both energy and matter with its surroundings (e.g., an open beaker, operating propane tank).

Closed System

A system that can exchange energy, but NOT matter, with its surroundings (e.g., a glow stick emitting light).

Isolated System

A system that cannot exchange either energy or matter with its surroundings.

Exothermic Change (Physical)

Energy is transferred from the system to the surroundings, causing surroundings to warm up (ΔH < 0). Example: Water freezing.

Endothermic Change (Physical)

Energy is transferred from the surroundings to the system, causing surroundings to cool down (ΔH > 0). Example: Ice melting.

Physical Phase Changes Examples

Dissolving, melting, freezing, evaporation, sublimation, condensation, deposition, and vaporization.

Evaporation vs. Vaporization

Evaporation occurs gradually at the liquid surface; Vaporization is the rapid transition from liquid to gas occurring at the boiling point.

Bond Breaking vs. Bond Formation

Bond breaking ALWAYS requires energy (endothermic); Bond formation ALWAYS releases energy (exothermic).

Exothermic Reaction

A chemical reaction where more energy is released during product bond formation than is required to break reactant bonds (surroundings warm up).

Exothermic Chemical Equation Location

The energy term is written on the PRODUCT side of the chemical equation (e.g., Hydrocarbon + O2 -> CO2 + H2O + energy).

Exothermic Energy Diagram Features

Reactants start with high potential energy (Ep) and products end with lower potential energy, making products more stable.

Endothermic Reaction

A chemical reaction where more energy is required to break reactant bonds than is released during product bond formation (surroundings cool down).

Endothermic Chemical Equation Location

The energy term is written on the REACTANT side of the chemical equation.

Endothermic Energy Diagram Features

Reactants start with low potential energy (Ep) and absorb energy to form products with higher potential energy.

Specific Heat Capacity (c)

The quantity of thermal energy required to raise the temperature of 1 g of a substance by 1 °C.

Specific Heat Capacity of Water

c_water = 4.18 J/(g·°C) (or 4.184 J/(g·°C)). Must be memorized!

Calorimetry

The process of measuring energy changes during a physical or chemical change using a calorimeter.

Calorimeter Components

Consists of an insulated reaction chamber, tight cover, thermometer, and stirrer (e.g., styrofoam cup, bomb calorimeter).

Bomb Calorimeter Usage

Used specifically to study chemical reactions that involve gases.

3 Key Calorimetry Assumptions

1. Negligible heat transfer with outside environment. 2. Negligible heat absorbed by calorimeter itself. 3. Dilute solutions have density (1.00 g/mL) and c (4.18 J/g·°C) of water.

Heat Equation

q = mcΔT (where q = heat energy, m = mass, c = specific heat capacity, ΔT = T_final - T_initial).

Sign of q Meaning

Positive (+q) indicates an endothermic process (energy gained); Negative (-q) indicates an exothermic process (energy lost).

Conservation of Energy in Calorimetry

q_system + q_surroundings = 0, which can also be written as q_system = -q_surroundings.

Calculating Mass (m) for Physical vs. Chemical Changes

Physical changes: solve separately via q = -q without adding masses. Chemical reactions in solution: ADD the masses of both reacting solutions together.

Common Exo/Endo Misconception in Calorimetry

Exothermic reactions release heat into water (water temp increases); Endothermic reactions absorb heat from water (water temp decreases).

Define a system:

the substance undergoing the chang

A student mixes two chemicals in a beaker. She places a thermometer

inside. As the reaction occurs, the temperature increases

chemical

exo

Two chemicals involved - system

Beaker, the student, the thermometer the air - surroundings

A piece of paper catches fire on a stove and burns.

chemical

Endo – catching fire

Exo-actual burning process

Paper and oxygen - system

Stove - surrounding

A chunk of an ionic compound is dropped into

water and dissolves. A thermometer is placed inside a beaker during this

process and the temperature is observed to decrease

Physical

endo

Ionic compound - system

Water, thermometer, beaker - surrounding

You left some change You return to get it and notice your money is extremely

hot.

Physical

Endo -Exo – you touching it,

coins - system

person,Air around outside in the sun. (sun warming) - surroundings

How does the temperature change of the system and the surroundings

when a gas changes into a liquid? Is this process endothermic or exothermic

A gas’s molecules have a higher kinetic energy because of their movement, to

become a liquid, this energy is released to the surroundings decreasing their

movement.

The surroundings would get warmer as a result and this is an

exothermic process

What energy changes need to occur for a substance to undergo

sublimation? Is this process exothermic or endothermic

Sublimation: solid → gas. For this to happen, the solid would need to absorb a

lot of energy which would give more kinetic energy to its particles and allow it

to become a gas.

Since energy is absorbed, this process would therefore be

endothermic

Distinguish between heat, thermal energy, and temperature using examples from your experience.

Temperature =A measure of the average speed (kinetic energy) of the individual particles in a substance. ex. Measuring a single cup of hot tea at 90c versus a cold swimming pool at 20c

Thermal Energy= The total overall energy (kinetic + potential) of all the particles in a substance combined. Because it accounts for every particle, mass matters.

ex. Comparing that same cup of hot tea 90c to the swimming pool 20c, the swimming pool has far more thermal energy. Even though the pool's molecules move slower on average, there are millions times more molecules in total.

Heat: The transfer or flow of thermal energy moving from a warmer object to a cooler object.

ex. Holding a cold metal spoon in your warm hand. Thermal energy naturally flows out of your warm hand and into the cold spoon until both reach the same temperature. That moving energy is heat.

• More Shared Electrons ($\sigma + \pi$):

  • C=O has 2 electron pairs, requiring more total energy to break than the 1 electron pair of C-O

• Shorter Bond Length:

  • The additional bond increases electrostatic attraction between the nuclei and shared electrons, pulling atoms closer C=O is shorter than C-O

• Greater Nuclear Attraction:

  • A shorter interatomic distance creates a stronger bond that requires significantly more energy to pull apart.