GACE Chemistry (6–12) (712)

The concept of the atom has evolved significantly over centuries. Ancient Greek philosopher Democritus first proposed that matter consists of indivisible particles called atomos around 400 BCE. However, scientific atomic theory began with John Dalton in the early 1800s.

Dalton's Atomic Theory included several key postulates. All matter is composed of extremely small particles called atoms. Atoms of a given element are identical in mass and properties. Atoms of different elements have different masses and properties. Atoms cannot be created, destroyed, or subdivided in chemical reactions. Compounds form when atoms of different elements combine in simple whole-number ratios.

J.J. Thomson discovered the electron in 1897 through cathode ray experiments. His plum pudding model proposed that atoms consisted of a positively charged sphere with negatively charged electrons embedded throughout, similar to plums in a pudding.

Ernest Rutherford conducted the famous gold foil experiment in 1911. Alpha particles were directed at thin gold foil. Most particles passed straight through, but some deflected at large angles, and a few bounced back. This led to the nuclear model where atoms have a small, dense, positively charged nucleus with electrons orbiting at a distance. Most of the atom is empty space.

Niels Bohr refined the atomic model in 1913 by proposing that electrons orbit the nucleus in specific energy levels or shells. Electrons can jump between energy levels by absorbing or emitting specific amounts of energy. This explained the line spectra of hydrogen.

The modern quantum mechanical model developed through contributions from Heisenberg, Schrödinger, and others. Electrons exist in orbitals, which are probability distributions rather than fixed paths. The Heisenberg uncertainty principle states that we cannot simultaneously know both the exact position and momentum of an electron.

Atoms consist of three fundamental subatomic particles. Protons carry a positive charge of +1 and have a mass of approximately 1.0073 atomic mass units. They are located in the nucleus, and the number of protons defines the element's atomic number.

Neutrons carry no charge and have a mass of approximately 1.0087 atomic mass units. They are also located in the nucleus and contribute to atomic mass. The number of neutrons can vary among atoms of the same element, creating isotopes.

Electrons carry a negative charge of -1 and have a negligible mass of approximately 0.00055 atomic mass units. They exist in orbitals surrounding the nucleus. In a neutral atom, the number of electrons equals the number of protons.

The mass number equals the sum of protons and neutrons. Atomic number equals the number of protons. Isotopes are atoms of the same element with different numbers of neutrons.

Standard atomic notation represents an element as the mass number written as a superscript and the atomic number as a subscript before the element symbol. For example, carbon-12 has 6 protons and 6 neutrons, while carbon-14 has 6 protons and 8 neutrons.

Isotopes have identical chemical properties because they have the same electron configuration. However, they differ in physical properties related to mass. Some isotopes are stable while others are radioactive.

Average atomic mass is calculated using the weighted average of all naturally occurring isotopes. The calculation multiplies each isotope's mass by its fractional abundance and sums the results. This explains why atomic masses on the periodic table are not whole numbers.

Electrons occupy orbitals according to specific rules. The Aufbau principle states that electrons fill orbitals starting with the lowest energy level. The Pauli exclusion principle states that no two electrons in an atom can have the same set of four quantum numbers, meaning each orbital holds a maximum of two electrons with opposite spins. Hund's rule states that electrons occupy orbitals of equal energy singly before pairing.

The four quantum numbers describe each electron's state. The principal quantum number (n) indicates the energy level and has positive integer values. The angular momentum quantum number (l) indicates orbital shape and ranges from 0 to n-1. The magnetic quantum number (ml) indicates orbital orientation and ranges from -l to +l. The spin quantum number (ms) indicates electron spin direction and has values of +1/2 or -1/2.

Orbital shapes correspond to angular momentum quantum numbers. When l equals 0, the orbital is spherical and designated as s. When l equals 1, the orbital is dumbbell-shaped and designated as p. When l equals 2, the orbital has a cloverleaf shape and designated as d. When l equals 3, the orbital has complex shapes and is designated as f.

The order of orbital filling follows this sequence: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p. This order reflects increasing energy levels, with some exceptions due to orbital energy overlap.

Electron configurations are written by listing occupied orbitals with superscripts indicating electron count. For example, oxygen with 8 electrons has the configuration 1s² 2s² 2p⁴. Noble gas notation abbreviates by using the previous noble gas in brackets. Oxygen becomes [He] 2s² 2p⁴.

Exceptions to expected configurations occur with chromium and copper. Chromium is [Ar] 3d⁵ 4s¹ rather than [Ar] 3d⁴ 4s² because half-filled d orbitals provide extra stability. Copper is [Ar] 3d¹⁰ 4s¹ rather than [Ar] 3d⁹ 4s² because fully filled d orbitals are particularly stable.

Valence electrons are those in the outermost principal energy level. For main group elements, the group number indicates the number of valence electrons. Transition metals typically have 2 valence electrons in the s orbital.

Lewis electron dot structures show valence electrons as dots around the element symbol. Electrons are placed one at a time on four sides of the symbol before pairing. This representation helps predict bonding behavior.

Core electrons are all electrons except valence electrons. They are closer to the nucleus and do not typically participate in chemical bonding. Shielding effect describes how core electrons reduce the effective nuclear charge experienced by valence electrons.

When electrons absorb energy, they jump to higher energy levels in a process called excitation. When electrons return to lower energy levels, they emit energy as electromagnetic radiation. The energy emitted corresponds to specific wavelengths of light.

The relationship between energy and wavelength is given by E = hν = hc/λ, where h is Planck's constant, ν is frequency, c is the speed of light, and λ is wavelength. Higher energy transitions produce shorter wavelengths.

Emission spectra show bright lines at specific wavelengths against a dark background. Absorption spectra show dark lines at specific wavelengths against a continuous spectrum background. Each element has a unique spectral fingerprint used for identification.

The hydrogen emission spectrum shows specific series of lines. The Lyman series involves transitions to n=1 and produces ultraviolet radiation. The Balmer series involves transitions to n=2 and produces visible light. The Paschen series involves transitions to n=3 and produces infrared radiation.