Chemistry Study Notes: Atomic Structure, Spectra, Electron Configurations, Bonding

Elements are the primary constituents of matter and cannot be chemically broken down into simpler substances.
Compounds consist of atoms of different elements chemically bonded together in a fixed ratio.
Mixtures contain more than one element or compound in no fixed ratio, are not chemically bonded, and can be separated by physical methods.
Distinguish between the properties of elements, compounds and mixtures.
The differences between homogeneous and heterogeneous mixtures should be understood.

The kinetic molecular theory is a model to explain the physical properties of matter (solids, liquids and gases) and changes of state.
Distinguish the different states of matter.
Names of the changes of state: melting, freezing, vaporization (evaporation and boiling), condensation, sublimation and deposition.

Atoms contain a positively charged, dense nucleus composed of protons and neutrons (nucleons).
Negatively charged electrons occupy the space outside the nucleus.
Use the nuclear symbol to deduce the number of protons, neutrons and electrons in atoms and ions.
Relative masses and charges of the subatomic particles should be known; actual values are given in the data booklet. The mass of the electron can be considered negligible.

Isotopes are atoms of the same element with different numbers of neutrons.
Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data.
Differences in the physical properties of isotopes should be understood.
Specific examples of isotopes need not be learned.

Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels.
Qualitatively describe the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum.
Distinguish between a continuous and a line spectrum.
Full electron configurations and condensed electron configurations using the noble gas core should be covered.
Orbital diagrams, i.e. arrow-in-box diagrams, should be used to represent the filling and relative energy of orbitals.
The electron configurations of Cr and Cu as exceptions should be covered.

The line emission spectrum of hydrogen provides evidence for electrons existing in discrete energy levels that converge at higher energies.
Describe the emission spectrum of hydrogen, including the relationships between the lines and energy transitions to the first, second and third energy levels.
The names of the different hydrogen series will not be assessed.

The main energy level is designated by an integer n and can hold a maximum of $2n^2$ electrons.
Deduce the maximum number of electrons that can occupy each energy level using the formula $N_{max}(n)=2n^2\,.$

A more detailed model divides the main energy level into sublevels s, p, d and f with successively higher energies.
Recognize the shape and orientation of:
- s orbital: spherical symmetry.
- p orbitals: three dumbbell-shaped orbitals oriented along the x, y and z axes (px, py, pz).

Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin.
Sublevels contain a fixed number of orbitals:
- s: 1 orbital
- p: 3 orbitals
- d: 5 orbitals
- f: 7 orbitals
Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z = 36.

In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization.
Explain the trends and discontinuities in first ionization energy (IE) across a period and down a group.
Calculate the value of the first IE from spectral data that gives the wavelength or frequency of the convergence limit.
The data booklet provides the Planck constant $h$ and the relations $E=hf$ and $c=\lambda f$ .

When metal atoms lose electrons, they form positive ions (cations). When non-metal atoms gain electrons, they form negative ions (anions).
Predict the charge of an ion from the electron configuration of the atom.
The formation of ions with different charges from a transition element should be included.

The ionic bond is formed by electrostatic attractions between oppositely charged ions.
Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix -ide. Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. Interconvert names and formulas of binary ionic compounds.
The following polyatomic ions should be known by name and formula: ammonium NH4+, hydroxide OH-, nitrate NO3-, hydrogencarbonate HCO3-, carbonate CO3^{2-}, sulfate SO4^{2-}, phosphate PO4^{3-}.

Ionic compounds exist as three-dimensional lattice structures, represented by empirical formulas.
Explain the physical properties of ionic compounds to include volatility, electrical conductivity and solubility.
Include lattice enthalpy as a measure of the strength of the ionic bond in different compounds, influenced by ion radius and charge.

A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.
The octet rule refers to the tendency of atoms to gain a valence shell with a total of 8 electrons. Deduce the Lewis formulas of molecules and ions for up to four electron pairs on each atom.
Lewis formulas (Lewis structures) show all the valence electrons (bonding and non-bonding pairs) in a covalently bonded species.
Electron pairs in a Lewis formula can be shown as dots, crosses or dashes. Molecules containing atoms with fewer than an octet of electrons should be covered; use organic and inorganic examples.

Single, double and triple bonds involve one, two and three shared pairs of electrons respectively.
Explain the relationship between the number of bonds, bond length and bond strength.

A coordination bond is a covalent bond in which both electrons of the shared pair originate from the same compound.
Identify coordination bonds in compounds (e.g. NH4+, H3O+, O3).