General Chemistry and Thermodynamics Flashcards

Vocabulary practice flashcards covering basic chemistry terminology, states of matter, bonding, thermodynamics, and acid-base theory based on the lecture transcript.

Heterogeneous

A mixture containing distinct regions made of separate substances.

Homogeneous

A mixture where substances mix evenly; also referred to as a solution.

Entropy

The amount of disorder within a system.

Solid

A state of matter with low entropy and a fixed volume where particles are bonded together and wiggle.

Liquid

A state of matter where particles have enough energy to move around but forces are strong enough to keep them together.

Gas

A state of matter with high entropy where particles move freely.

Plasma

An ionized gas that exists at very high temperatures or high electric potential, such as in stars.

Metallic bonds

A grid of positively charged nuclei surrounded by freely moving, delocalized electrons.

Isomers

Compounds that have the same molecular formula but are not the same structure.

Electronegativity

The strength of a covalent bond where a nucleus tugs on the electrons of another atom.

Polar covalent bond

A bond where one element pulls on the electrons but does not steal them, creating an electric dipole with partial charges ($\delta+$ and $\delta-$).

Nonpolar covalent bond

A bond where electrons are shared equally, typically occurring when the electronegativity difference is less than $0.5$.

Ionic bond

A bond where one element needs an electron and another has an extra, resulting in the transfer of electrons and the formation of a cation and an anion.

Catalyst

A substance that reduces the activation energy needed for a chemical reaction to occur.

Enthalpy ($H$)

The internal energy or heat content of a system.

Exothermic

A reaction where heat is given off because the energy at the beginning is greater than at the end ($\Delta H$ is negative).

Endothermic

A reaction where heat is absorbed because the energy at the beginning is less than at the end ($\Delta H$ is positive).

Gibbs Free Energy ($G$)

The energy available to do work, calculated by the equation \Delta G = \text{\Delta H} - T\text{\Delta S}, used to determine if a reaction is spontaneous.

Exergonic

A spontaneous reaction where energy is released and $\Delta G < 0$.

Endergonic

A non-spontaneous reaction where energy is absorbed and $\Delta G > 0$.

Aufbau principle

The rule used to determine electron configuration by filling orbitals from the lowest energy level to the highest.

Stoichiometry

The study of numerical relationships and ratios between reactants and products in chemical reactions.

Mole

A unit used to measure the amount of a substance.

pH

A measure of the concentration of Hydronium ($[H_3O^+]$) calculated as $pH = -\log_{10}([H_3O^+])$; values below $7$ are acidic.

pOH

A measure of basicity calculated as $pOH = -\log_{10}([OH^-])$; it is related to pH by the equation $pH + pOH = 14$.

Br%nsted-Lowry acid

A molecule that donates protons ($H^+$).

Br%nsted-Lowry base

A molecule that accepts protons ($H^+$).

Amphoteric

A substance, such as water, that can act as both an acid and a base.

Redox Reactions

Reduction-oxidation reactions that involve the transfer of electrons and changes in the oxidation numbers of elements.

Chemical equilibriums

A state that exists when reversible reactions take place at the same speed in both the forward and reverse directions ($\Delta G = 0$).

Electronic Quantum Numbers

Four values used to describe electrons: $n$ (shell), $l$ (subshell shape), $m_l$ (orbital orientation), and $m_s$ (intrinsic spin).