Chemisty - Unit 2 ✨

Connections Academy 🥰

bond energy

the energy that is required to break chemical bonds

chemical bond

the force that holds atoms, molecules, or ions together to form chemical compounds

endothermic reaction

a chemical reaction that absorbs energy from the surrounding environment

enthalpy (H)

the sum of a system’s internal energy and the product of its pressure and volume

exothermic reaction

a chemical reaction that releases energy into the environment

law of conservation of energy

energy can be neither created nor destroyed, only transformed between forms (e.g., bond ↔ thermal)

ΔH (enthalpy change)

Hproducts − Hreactants; negative for exothermic, positive for endothermic

Hess’s law

total enthalpy change is the same whether a reaction occurs in one step or multiple steps

standard enthalpy of formation

ΔH when one mole of a compound forms from its elements in standard states

bond dissociation enthalpy

the energy required to homolytically cleave a bond, yielding radicals

system

the reactants and products in a chemical reaction being studied

surroundings

everything outside of that system

Additional/Internal

open vs. closed vs. isolated system

can exchange energy and/or matter differently

state function

property (e.g., enthalpy, internal energy) that depends only on the system’s state, not on how it got there

spontaneous process

occurs without outside intervention; indicated by ΔG < 0 under standard conditions

Gibbs free energy (G)

G

calorimeter

apparatus used to measure heat flow between the system and surroundings

chemical bond

an electrical force that holds the atoms of molecules together

bond energy

the energy required to break apart a chemical bond

activation energy (Eₐ)

minimum energy needed to initiate bond breaking in reactants

transition state

high‑energy, fleeting configuration at the top of the Eₐ hill

collision theory

molecules must collide with proper orientation and sufficient energy to react

reaction coordinate diagram

plots potential energy versus reaction progress, showing Eₐ and ΔH

catalyst

provides an alternative pathway with lower activation energy without being consumed

electronegativity

a measure of the ability of an atom to attract electrons to itself

chemical reaction system

the reactants and products in a particular reaction being studied

covalent bond

sharing of electron pairs between atoms

ionic bond

electrostatic attraction between oppositely charged ions

polar covalent bond

unequal sharing of electrons due to electronegativity difference

Lewis structure

diagram showing the arrangement of electrons in molecules

resonance

delocalization of electrons across adjacent bonds, lowering overall energy

conduction

the transfer of energy by direct contact

convection

the transfer of energy by currents of moving fluids

radiation

the transfer of energy as electromagnetic waves

endothermic reaction

a chemical reaction in which there is a net input of energy from the surroundings into the system

exothermic reaction

a chemical reaction in which there is a net output of energy from the system into the surroundings

kinetic energy (KE)

energy of motion

potential energy (PE)

energy due to position or composition

activation energy

the energy that must be added to a system to break the bonds of the reactants

thermochemistry

the study of energy changes in chemical and physical changes

Maxwell–Boltzmann distribution

statistical spread of molecular kinetic energies at a given T

degrees of freedom

translational, rotational, vibrational modes that contribute to KE

enthalpy vs. internal energy (U)

U = q + w; under constant pressure, ΔH = ΔU + PΔV

endotherm/exotherm on PE diagram

uphill vs. downhill net change in PE from reactants to products

heat (q) vs. work (w)

energy transfer as heat versus mechanical work

relative potential energy

the difference in potential energies of reactants and products that determines net energy change

ionic bond

the bond between oppositely charged ions in a compound

catalyst

a substance that speeds up a reaction by lowering activation energy

lattice energy

energy released when gaseous ions form an ionic solid; related to ionic bond strength

Born–Haber cycle

thermochemical cycle for calculating lattice energies via Hess’s law

bond order

number of shared electron pairs between two atoms; higher bond order → greater bond energy

Madelung constant

geometric factor in calculating lattice energy for ionic crystals

thermodynamic stability

lower PE species are more stable; exothermic formation → more stable products