Reactions Unit 1

Endothermic

A reaction that absorbs energy from the surroundings, resulting in a positive enthalpy change (ΔH).

Exothermic

A reaction that releases energy to the surroundings, resulting in a negative enthalpy change (ΔH).

Enthalpy Changes (ΔH)

The heat energy exchange in a reaction at constant pressure.

ΔH (negative)

Indicates an exothermic reaction where energy is released.

ΔH (positive)

Indicates an endothermic reaction where energy is absorbed.

Calorimetry Assumptions

All heat was transferred entirely to/from the 'surroundings' measured, with no heat loss to the environment.

In solutions, water's exact density and specific heat must be used.

Complete combustion occurred.

Standard Enthalpy Changes

The enthalpy change when a reaction occurs under standard conditions (298 K, 1 atm pressure).

Hess's Law

The total enthalpy change for a reaction is independent of the path taken, depending only on initial and final states.

Bond Enthalpies

The average energy required to break one mole of a specific bond in a gaseous molecule.

Enthalpy of Combustion (ΔHc)

The enthalpy change when one mole of a substance completely combusts (reacts with excess oxygen) under standard conditions.

Enthalpy of Formation (ΔHf)

The enthalpy change when one mole of a substance is formed from its elements in their standard states.

Covalent Bond Enthalpy

The energy required to break one mole of a covalent bond between two atoms in the gaseous state.

Atomization Energy

The energy required to convert one mole of a substance in its standard state into gaseous atoms, always endothermic.

Ionization Energy (IE)

The energy required to remove one mole of electrons from one mole of gaseous atoms, forming gaseous ions, always endothermic.

Electron Affinity (EA)

The energy change when one mole of electrons is added to one mole of gaseous atoms, forming gaseous ions; 1st is exothermic, 2nd is endothermic.

Lattice Enthalpy

The energy change when one mole of an ionic lattice is formed from its gaseous ions, always very exothermic.

Magnitude is based on ionic bond strengths: smaller and higher charge = stronger.

Born-Haber Cycle

An enthalpy cycle representing the formation of an ionic compound from its elements in their standard states.

Combustion

A chemical reaction in which a substance reacts rapidly with O2 in excess; forms oxides. Can occur from hydrocarbons (most common), nonmetals, and reactive metals.

Incomplete Combustion

Combustion that occurs with a limited oxygen supply, producing less energy than complete combustion, but same amounts of CO and/or CO2

Fossil Fuels

Carbon-based fuels formed from the remains of ancient organisms, examples include coal, petroleum, and natural gas.

Biofuels

Fuels derived from recently living organisms, such as ethanol and biodiesel.

Fuel Cells

Electrochemical devices that convert chemical energy of H2 directly into electrical energy with high efficiency and minimal pollutants. Only product is H2O.

Greenhouse Gases

Gases that trap heat in the atmosphere, including carbon dioxide (CO2), methane (CH4), and water vapor (H2O). Allow sunlight to enter but prevent some of the heat from escaping back into space, contributing to the greenhouse effect.

Greenhouse gases absorb infrared radiation emitted from Earth's surface b/c molecular structure. Bonds stretch and bend, leading to increase in vibrations & warming of atmosphere.

Increases in greenhouse gases from human activities, such as burning fossil fuels and biofuels, are major factor in global warming and climate change.

Where does oxidation occur in fuel cells? What is the half equation for the reactant?

Oxidation occurs at the anode, where hydrogen molecules are split into protons and electrons. The half equation for the reactant is: H2 → 2H⁺ + 2e⁻.

Coal

Crude Oil

Natural Gases

Ethanol

Incomplete combustion increases with…

Length of carbon chain

pros of coal

pros of crude oil

pros of natural gas

cons of coal

cons of crude oil

cons of natural gas