General and Inorganic Chemistry for Physics and TUM-BWL – Vocabulary Flashcards

A set of vocabulary-focused flashcards covering key terms and concepts from the lecture notes on general and inorganic chemistry.

Lewis structure

A diagram that shows how atoms are bonded in a molecule and where the lone pairs of electrons reside, using valence electrons.

Valence electron

An electron located in the outermost shell of an atom that participates in chemical bonding.

Valence Shell Electron Pair Repulsion (VSEPR) theory

A model predicting molecular geometry by minimizing repulsions between electron pairs around a central atom.

Octet rule

A tendency for main-group elements to prefer eight electrons in their valence shell when forming molecules.

Formal charge

A bookkeeping method to estimate charge distribution: FC = valence electrons − nonbonding electrons − 1/2 bonding electrons.

Resonance

The concept that a molecule’s true structure is a hybrid of multiple valid Lewis structures due to delocalized electrons.

Covalent bond

A bonds formed by sharing one or more pairs of electrons between atoms.

Ionic bond

A bond formed by the electrostatic attraction between oppositely charged ions.

Bond order

In MO theory, the bond order is (bonding electrons − antibonding electrons)/2; it reflects bond strength.

Molecular Orbital (MO) theory

A model describing electrons in molecules as occupying molecular orbitals formed from atomic orbitals.

LCAO (Linear Combination of Atomic Orbitals)

A method to construct molecular orbitals by combining atomic orbitals linearly.

Bonding molecular orbital

A MO that lowers energy and stabilizes the molecule when occupied.

Antibonding molecular orbital

A MO that raises energy; occupancy reduces bond stability.

Aufbau principle

Electrons fill orbitals in order of increasing energy, subject to Pauli and Hund’s rules.

Pauli exclusion principle

No two electrons in an atom can have identical quantum numbers; each orbital holds at most two electrons with opposite spins.

Hund’s rule

Electrons fill degenerate orbitals singly with parallel spins before pairing occurs.

Quantum numbers (n, l, ml, ms)

n: principal; l: azimuthal; ml: magnetic; ms: spin quantum numbers describing electron (orbital) states.

Bohr model

Early atomic model with quantized electron orbits and energy levels, primarily for hydrogen.

Rydberg constant (R∞)

A constant used to describe hydrogen spectral lines; R∞ ≈ 1.097×10^7 m⁻¹.

Hydrogen spectral series (Lyman, Balmer, Paschen, Brackett, Pfund)

Groups of lines corresponding to electron transitions to n = 1, 2, 3, 4, 5 respectively.

Avogadro’s number

NA = 6.02214076×10^23 entities per mole; defines the amount of substance in a mole.

Mole (mol)

The SI unit for amount of substance; it contains NA elementary entities.

First law of thermodynamics

Energy is conserved: ΔU = q + w for a closed system.

Enthalpy (H)

Thermodynamic function H = U + pV; useful for heat changes at constant pressure.

Standard enthalpy of formation (ΔH°f)

Enthalpy change when forming 1 mole of a substance from its elements in standard states.

Reaction enthalpy (ΔH°)

Enthalpy change for a reaction at standard conditions; can be computed via Hess’s law.

Hess’s law

The total enthalpy change of a reaction equals the sum of enthalpy changes of its steps.

Gibbs free energy (G)

G = H − T S; determines spontaneity of processes; negative ΔG implies spontaneity at given T and p.

Entropy (S)

A measure of disorder; in statistical form S = k ln W, where W is the number of microstates.

Boltzmann’s formula

S = k ln W; links microscopic states to macroscopic entropy.

Second law of thermodynamics

Entropy of the universe tends toward a maximum in spontaneous processes; processes move toward greater disorder.

Equilibrium and ΔG°

At equilibrium, ΔG = 0 and ΔG° = −RT ln K; K = exp(−ΔG°/RT) relates free energy to equilibrium constant.

Arrhenius equation

k = A exp(−Ea/(R T)); relates rate constant to temperature and activation energy.

Activation energy (Ea)

Energy barrier that must be overcome for a reaction to proceed.

Reaction order

Kinetic order with respect to reactants (unimolecular, bimolecular, trimolecular).

First-order kinetics

Rate law: d[A]/dt = −k[A]; integrated form: ln[A] = −kt + ln[A]0.

Lambert–Beer's law

I = I0 10^(−ε c l); absorbance A = ε c l; relates light transmission to concentration and path length.

Molar absorptivity (ε)

Material-specific constant describing how strongly a species absorbs light at a given wavelength.

Electrochemistry (Q and F)

Q = N e; F ≈ 96485 C/mol (Faraday’s constant); I = dQ/dt; relates charge to moles of electrons.

Electrolysis

Non-spontaneous redox process driven by external electrical energy; e.g., molten NaCl or CuSO4 solutions.

Electrolysis of molten NaCl

2 NaCl(l) → 2 Na(l) + Cl2(g) at the respective electrodes.

Electrolysis of CuSO4

Cathode: Cu2+ + 2 e− → Cu(s); Anode: Cu(s) → Cu2+ + 2 e−.

Autoionization of water

Water self-ionizes: Kw = [H3O+][OH−] ≈ 1.0×10^−14 at 25°C.

pH and pOH

pH = −log10[H+] and pOH = −log10[OH−]; at 25°C, pH + pOH ≈ 14.

Millikan experiment

Experiment to measure the elementary charge e by balancing gravitational and electric forces on droplets.

Electron charge (e)

Elementary charge: e ≈ 1.602×10^−19 coulomb.

Proton and neutron

Nucleons in the nucleus; protons carry positive charge, neutrons are neutral; electrons are far lighter.

Atomic mass number (A) and atomic number (Z)

A = total number of protons and neutrons; Z = number of protons (atomic number).

Electron affinity (EA)

Energy change when adding an electron to a gaseous atom; generally releases energy.

van der Waals forces

Intermolecular forces including dipole–dipole (Keesom), dipole–induced dipole (Debye), and London dispersion forces.

Electronegativity (Pauling scale)

Ability of an atom in a bond to attract shared electrons; differences in EN lead to bond polarity.

Periodic table blocks

s-block, p-block, d-block, f-block; classifications reflecting electron configurations and properties.

Hydrogen bond

A strong type of intermolecular force between a hydrogen atom bound to a highly electronegative atom (N, O, F) and a lone pair on another electronegative atom.

Pf, Balmer, Lyman series (spectroscopy)

Hydrogen spectral series corresponding to transitions to n = 1 (Lyman), 2 (Balmer), 3 (Paschen) etc.

Rutherford scattering

Experiment showing alpha particles scattered by a dense nucleus, leading to nuclear model of the atom.

Stern–Gerlach experiment

Demonstrated spatial quantization of angular momentum (electron spin) using a magnetic field.

Democritus and Dalton

Historical atomistic theories: Democritus proposed indivisible atoms; Dalton proposed atomic theory with conservation of mass and fixed atomic ratios.

Rutherford-Bohr synthesis

Early concepts connecting nucleus-centric atoms with quantized energy levels and spectral observations.

Electron affinity vs ionization energy

EA measures energy change when gaining an electron; ionization energy measures energy to remove an electron.

Chemistry blocks and bonding types

Bonding types depend on electronegativity and how electrons are shared/transferred between atoms.

Schrödinger equation (orbitals)

Mathematical description of how the quantum state of a physical system changes; used to derive atomic and molecular orbitals.

Term symbols and orbital notation

Notation describing electron configurations and resulting angular momentum of atoms and molecules.

Spectroscopic units (eV, nm, cm−1)

Common units to express energy (eV), wavelength (nm), and wavenumber (cm−1) in spectroscopy.

Aufbau principle without exceptions

Idealized filling order of electron orbitals; real systems may show deviations (e.g., Cu, Cr).

Electrolyte vs electrode

Electrolyte is a medium allowing ion conduction; electrode is a solid used as an electrical contact in electrochemical cells.

LCAO-MO energy relation

Energy of molecular orbitals arises from the linear combination of atomic orbitals and their interactions.