Chemistry 30 Diploma flashcards

32 cards. Read the term and the definition here. The deck at the top of this page is optional.

Enthalpy change (ΔH)

The heat exchanged with the surroundings at constant pressure; negative for exothermic reactions that release energy and positive for endothermic reactions that absorb it.

Calorimetry equation q = mcΔT

Heat gained or lost equals mass times specific heat capacity times temperature change; it is applied to the water or solution to find the heat of a reaction.

Molar enthalpy

The enthalpy change divided by the moles of a specified reactant or product; it is quoted in kJ/mol so that different reactions and fuels can be compared.

Hess's law

The enthalpy change of an overall reaction equals the sum of the enthalpy changes of its individual steps, because enthalpy is a state function independent of pathway.

Standard enthalpy of formation

The enthalpy change when one mole of a compound forms from its elements in their standard states; it is zero for an element in its standard state.

Bond energy estimate for ΔH

An approximate enthalpy change equal to the sum of the average bond energies of bonds broken minus the sum of the bond energies of bonds formed, for gaseous species.

Activation energy

The minimum energy colliding particles need to form the activated complex; on a potential energy diagram it is the height from the reactants to the peak.

Enthalpy of fusion and vaporization

The heat absorbed when one mole of a substance melts or boils; the temperature stays constant during the change and q = nΔH gives the energy.

Oxidation

Loss of electrons by a species, shown by an increase in oxidation number; oxidation always happens at the anode of an electrochemical cell.

Cathode

The electrode where reduction occurs; it is positive in a galvanic cell and negative in an electrolytic cell, and electrons arrive there in both cases.

Standard cell potential

E°cell = E°cathode − E°anode using standard reduction potentials; the cell reaction is spontaneous when the calculated value of E°cell is positive.

Salt bridge

A tube of electrolyte that completes the circuit and keeps both half-cells electrically neutral by letting ions migrate without mixing the two solutions.

Gibbs energy from cell potential

ΔG° = −nFE°, where n is the moles of electrons transferred and F is 96 500 C/mol; a positive cell potential gives a negative Gibbs energy change.

Electrolytic cell

A cell in which an external power supply drives a nonspontaneous reaction; the anode is positive, the cathode is negative, and electrical energy becomes chemical energy.

Faraday's law of electrolysis

The moles of electrons passed equal the current in amperes times the time in seconds divided by 96 500; half-reaction stoichiometry then gives moles of product.

Strong oxidizing agent

A species with a high standard reduction potential, which is easily reduced and therefore removes electrons from, and oxidizes, other species.

Alkane general formula

CnH2n+2; saturated hydrocarbons containing only single carbon–carbon bonds, named with the suffix -ane and reacting mainly by substitution reactions.

Alkene general formula

CnH2n for a non-cyclic alkene; unsaturated hydrocarbons with one carbon–carbon double bond, named with the suffix -ene and reacting by addition.

Markovnikov's rule

When a hydrogen halide adds to an alkene, the hydrogen attaches to the carbon holding more hydrogens, leaving the halogen on the more substituted carbon.

Oxidation of alcohols

Primary alcohols form aldehydes and then carboxylic acids; secondary alcohols form ketones; tertiary alcohols are not oxidized under these conditions.

Esterification

A carboxylic acid heated with an alcohol and a strong acid catalyst forms an ester and water; the reaction is reversible and esters are often fragrant.

Cis-trans isomerism

A form of geometric isomerism in alkenes: identical groups on the same side of the double bond are cis, and on opposite sides they are trans.

Branching and boiling point

More branched isomers have lower boiling points than straight-chain isomers of the same formula because their molecules have less surface contact and weaker dispersion forces.

Dehydration of an alcohol

Heating an alcohol with concentrated sulfuric acid removes the OH group and a hydrogen from an adjacent carbon, producing an alkene and water.

Equilibrium constant K

The ratio of product concentrations to reactant concentrations, each raised to its coefficient in the balanced equation; pure solids and liquids are excluded and K depends only on temperature.

Le Châtelier's principle

When a system at equilibrium is disturbed by a change in concentration, temperature or pressure, it shifts in the direction that partially offsets that change.

Ion product constant of water

Kw = [H3O+][OH−] equals 1.0 × 10^−14 at 25 °C; it allows conversion between hydronium and hydroxide concentrations in aqueous solution.

pH and pOH

pH = −log[H3O+] and pOH = −log[OH−]; in aqueous solution at 25 °C, pH + pOH = 14, so one can be found from the other.

Conjugate acid-base pair

Two species that differ by one proton; the stronger the acid, the weaker its conjugate base, and Ka multiplied by Kb equals Kw for the pair.

Buffer solution

A mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid, that resists pH change when small amounts of strong acid or base are added.

Henderson-Hasselbalch equation

pH = pKa + log([conjugate base]/[weak acid]); it gives the pH of a buffer and the component ratio needed to reach a target pH.

Equivalence point in titration

The point where the moles of added acid or base exactly match the moles required by the balanced equation; the pH is 7 only for strong acid–strong base titrations.