Chemistry A Test 1

Element

Pure chemical substance; cannot be broken down into simpler substances by chemical reactions.

Atom

Smallest unit of an element; retains all characteristic chemical properties of that element.

Compound

Substance formed by chemically combining two or more distinct elements in fixed proportions.

Thomson's Plum Pudding Model

Atomic model picturing negatively charged electrons scattered within a uniform sphere of positive charge.

Rutherford's Alpha Scattering Experiment

Alpha particles fired at gold foil; proved atoms are mostly empty space with a dense, positively charged nucleus.

Nucleons

Subatomic particles located in the atomic nucleus; protons and neutrons.

Proton mass and charge

Relative mass = $1$; relative charge = $+1$.

Neutron mass and charge

Relative mass = $1$; relative charge = $0$.

Electron mass and charge

Relative mass = $0.0005$ (negligible); relative charge = $-1$.

Atomic number ($Z$)

Number of protons in an atomic nucleus; uniquely defines the element.

Mass number ($A$)

Total number of nucleons; sum of protons and neutrons in the nucleus.

Isotopes

Atoms of the identical element with the same proton number ($Z$) but different neutron numbers.

Cation

Positively charged ion; formed by losing electrons.

Anion

Negatively charged ion; formed by gaining electrons.

Relative atomic mass ($A_r$)

Weighted average mass of an atom compared to $\frac{1}{12}$ of a carbon-12 atom; accounts for natural isotopic abundances.

Speed of light equation

$c = \nu \lambda$; relates light speed ($c$) to frequency ($\nu$) and wavelength ($\lambda$).

Planck equation

$E = h\nu$; relates photon energy ($E$) to frequency ($\nu$) via Planck's constant ($h$).

Continuous spectrum

Unbroken band displaying all frequencies and wavelengths without gaps; produced by white light through a prism.

Line emission spectrum

Series of discrete, sharp colored lines; emitted when excited electrons fall back to lower energy levels.

Line absorption spectrum

Continuous spectrum interrupted by dark lines; produced when electrons absorb specific frequencies to transition upward.

Hydrogen spectral series

Lyman: UV emissions falling to $n = 1$; Balmer: visible emissions falling to $n = 2$; Paschen: IR emissions falling to $n = 3$.

Quantization of energy

Electron energies are restricted to fixed, discrete levels; intermediate values are forbidden.

Heisenberg's Uncertainty Principle

Impossible to simultaneously determine both the exact position and momentum of an electron.

Atomic orbital

Three-dimensional space around a nucleus with a $90\%$ probability of finding an electron.

Shapes of s and p orbitals

$s$ orbitals are spherical; $p$ orbitals ($p_x$, $p_y$, $p_z$) are dumbbell-shaped along mutually perpendicular axes.

Pauli Exclusion Principle

No two electrons in the same atoms can have the same set of four quantum numbers

Aufbau Principle

Electrons fill orbitals of the lowest available energy levels before higher levels are occupied.

Hund's Rule

Degenerate orbitals are occupied singly with parallel spins before pairing begins.

Chromium ($\text{Cr}$) electron configuration exception

$[\text{Ar}]\,3d^5\,4s^1$; favored due to the enhanced stability of a half-filled $3d$ sub-level.

Copper ($\text{Cu}$) electron configuration exception

$[\text{Ar}]\,3d^{10}\,4s^1$; favored due to the enhanced stability of a completely filled $3d$ sub-level.

Transition metal ionization rule

Electrons are removed from the outer $4s$ sub-level before any $3d$ electrons are lost.

First ionization energy

Energy required to remove one mole of electrons from one mole of gaseous atoms in ground state: $X(g) \rightarrow X^+(g) + e^-$; unit is $kJ\,mol^{-1}$.

Periodic trends in first ionization energy

Increases across a period (higher effective nuclear charge); decreases down a group (increased atomic radius and shielding).

Large Jumps in Ionization Energy

A massive jump occurs when an electron is removed from an inner main energy level closer to the nucleus, revealing the number of outer-shell electrons

Successive Ionization Energies

Each successive ionization energy is higher because removing an electron from an increasingly positive ion requires more energy.