Chemistry: Atomic Structure, Electron Configurations, Periodic Trends & Bonding

Vocabulary flashcards covering key concepts in atomic structure, electron configurations, periodic trends, bonding, isotopes, and basic chemistry quantities.

Electrostatic force

The attraction between opposite charges (positive and negative) that holds atoms together.

Mass number (A)

Total number of protons and neutrons in the nucleus of an atom.

Atomic number (Z)

Number of protons in the nucleus; defines the element.

Neutron number (N)

Number of neutrons in the nucleus.

Ion

Atom that has gained or lost electrons to form a charged species.

Cation

Positive ion; typically named the same as the element.

Anion

Negative ion; named by replacing the ending with -ide.

Electron configuration

Arrangement of electrons in orbitals around the nucleus according to Aufbau principles.

Orbital

3D region around the nucleus that can hold up to 2 electrons.

Subshell

Set of orbitals within a shell with the same energy (s, p, d, f).

s subshell

1 orbital; holds up to 2 electrons.

p subshell

3 orbitals; holds up to 6 electrons.

d subshell

5 orbitals; holds up to 10 electrons.

f subshell

7 orbitals; holds up to 14 electrons.

Aufbau principle

Electrons fill the lowest-energy orbitals first.

Pauli Exclusion Principle

An orbital can hold at most two electrons with opposite spins.

Hund's Rule

Orbitals of the same energy are filled singly before pairing begins.

Copper electron configuration exception

Cu often favors [Ar] 3d10 4s1 for stability.

Ground state

The lowest-energy arrangement of electrons in an atom.

First ionisation energy

Energy required to remove one mole of electrons to form X+(g).

Ionisation energy trend across a period

Increases across a period as nuclear charge rises and radius falls.

Ionisation energy trend down a group

Decreases down a group due to greater distance and shielding.

Atomic radius

Half the distance between the nuclei in a diatomic molecule.

Across a period: atomic radius decreases

Radius gets smaller with more protons and less shielding.

Down a group: atomic radius increases

Radius gets larger as additional electron shells are added.

Electronegativity

Attraction of an atom in a covalent bond for bonding electrons; increases across a period, decreases down a group.

Metallic character

Tendency of an atom to lose electrons; increases down a group, decreases across a period.

Ionic bonding

Electrostatic attraction between oppositely charged ions in a lattice (e.g., NaCl).

Covalent bonding

Sharing of electron pairs between atoms.

Covalent network

Giant 3D network of covalent bonds (e.g., diamond).

Covalent molecular

Molecules held together by covalent bonds with weaker intermolecular forces.

Diamond

Giant covalent network with each carbon bonded to four others; very high MP, non-conductive.

Graphite

Layered covalent network with delocalized electrons; conducts electricity and heat; layers slide.

Metallic bonding

Attraction between positive metal ions and a sea of delocalized electrons; strong, malleable, good conductors.

Isotope

Atoms of the same element with different numbers of neutrons.

Relative isotopic mass

Mass of a specific isotope relative to 1/12 the mass of carbon-12.

Relative atomic mass (Ar)

Weighted average mass of an element’s isotopes relative to 1/12 of carbon-12.

Mass spectrometry

Technique to determine isotopic composition and relative abundances by mass/charge (m/z).

Homogeneous mixture

Uniform composition throughout; single phase.

Heterogeneous mixture

Non-uniform composition; distinct phases or layers.

Pure substance

One kind of matter; either an element or compound.

Element

Pure substance made of one type of atom.

Compound

Substance formed from two or more elements bonded chemically.

Mole

Amount of substance containing Avogadro’s number of particles (6.02×10^23).

Avogadro’s number

6.02×10^23 particles per mole.

Molar mass

Mass per mole of a substance (g/mol); equals relative molecular/formula mass.

Endothermic

Reactions that absorb heat from surroundings; feel cold.

Exothermic

Reactions that release heat to surroundings; feel hot.

Enthalpy change (delta H)

Heat absorbed or released per mole during a reaction.

Activation energy

Minimum energy required for collisions to lead to a reaction.

Bond enthalpy

Average energy required to break a specific bond per mole.

Theoretical yield

Maximum amount of product that could be formed from given reactants (no losses).

Experimental yield

Actual amount of product obtained from an experiment.

Percent yield

(Experimental yield / Theoretical yield) × 100.

Limiting reagent

Reactant completely consumed, limiting product formation.

Excess reagent

Reactant not completely consumed after reaction.

Percent composition

Mass percent of each element in a compound.

Alkane

Saturated hydrocarbons with only single C–C bonds; prefixes meth, eth, prop, but, etc.

Alkene

Unsaturated hydrocarbons with at least one C=C double bond.

Alkyne

Hydrocarbons with a C≡C triple bond.

Aromatic (arenes)

Hydrocarbons containing at least one benzene ring.

Isomer

Compounds with the same molecular formula but different arrangements.

Homologous series

Family of organic compounds with the same functional pattern and CH2 increments.

Ionic compound

Compound held together by ionic bonds; forms a lattice and is often soluble in water.

Crystal lattice

3D repeating arrangement of ions in an ionic solid.

Absorption spectrum

Spectrum showing wavelengths absorbed by a substance (dark lines).

Emission spectrum

Spectrum showing wavelengths emitted by a substance (bright lines).

Wavelength and energy relationship

Longer wavelength = lower energy; shorter wavelength = higher energy.