Structure
Structure refers to the nature of matter from simple to more complex forms.
Structure 1 - Models of the particulate nature of matter
Structure 1.1 - Introduction to the particulate nature of matter
- Structure 1.1.1 - Elements are the primary constituents of matter, which cannot be chemically broken down into simpler substances.
- Structure 1.1.2 - The kinetic molecular theory is a model to explain physical properties of matter (solids, liquids and gases) and changes of state.
- Structure 1.1.3 - The temperature, T, in Kelvin (K) is a measure of average kinetic energy Ek of particles.
Structure 1.2 - The nuclear atom
- Structure 1.2.1 - Atoms contain a positively charged, dense nucleus composed of protons and neutrons (nucleons). Negatively charged electrons occupy the space outside the nucleus.
- Structure 1.2.2 - Isotopes are atoms of the same element with different numbers of neutrons.
- Structure 1.2.3 - Mass spectra are used to determine the relative atomic masses of elements from their isoindex composition.
Structure 1.3 - Electron configurations
- Structure 1.3.1 - Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels.
- Structure 1.3.2 - The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.
- Structure 1.3.3 - The main energy level is given an integer number, n, and can hold a maximum of 2n² electrons.
- Structure 1.3.4 - A more detailed model of the atom describes the division of the main energy level into s, p, d and f sub-levels of successively higher energies.
- Structure 1.3.5 - Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin.
- Structure 1.3.6 - In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization.
- Structure 1.3.7 - Successive ionization energy (IE) data for an element give information about its electron configuration.
Structure 1.4 - Counting particles by mass: The mole
- Structure 1.4.1 - The mole (mol) is the SI unit of amount of substance. One mole contains exactly the number of elementary entities given by the Avogadro constant.
- Structure 1.4.2 - Masses of atoms are compared on a scale relative to 12C and are expressed as relative atomic mass Ar and relative formula mass Mr.
- Structure 1.4.3 - Molar mass M has the units g mol–1.
- Structure 1.4.4 - The empirical formula of a compound gives the simplest ratio of atoms of each element present in that compound. The molecular formula gives the actual number of atoms of each element present in a molecule.
- Structure 1.4.5 - The molar concentration is determined by the amount of solute and the volume of solution.
- Structure 1.4.6 - Avogadro’s law states that equal volumes of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules.
Structure 1.5 - Ideal gases
- Structure 1.5.1 - An ideal gas consists of moving particles with negligible volume and no intermolecular forces.
- Structure 1.5.2 - Real gases deviate from the ideal gas model, particularly at low temperature and high pressure.
- Structure 1.5.3 - The molar volume of an ideal gas is a constant at a specific temperature and pressure.
- Structure 1.5.4 - The relationship between the pressure, volume, temperature and amount of an ideal gas is shown in the ideal gas equation PV = nRT and the combined gas law P₁V₁/T₁ = P₂V₂/T₂.
Structure 2.1 - The ionic model
- Structure 2.1.1 - When metal atoms lose electrons, they form positive ions called cations. When non-metal atoms gain electrons, they form negative ions called anions.
- Structure 2.1.2 - The ionic bond is formed by electrostatic attractions between oppositely charged ions.
- Structure 2.1.3 - Ionic compounds exist as three-dimensional lattice structures, represented by empirical formulas.
Structure 2.2 - The covalent model
- Structure 2.2.1 - A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.
- Structure 2.2.2 - Single, double and triple bonds involve one, two and three shared pairs of electrons respectively.
- Structure 2.2.3 - A coordination bond is a covalent bond in which both the electrons of the shared pair originate from the same atom.
- Structure 2.2.4 The valence shell electron pair repulsion (VSEPR) model enables the shapes of molecules to be predicted from the repulsion of electron domains around a central atom
- Structure 2.2.5 Bond polarity results from the difference in electronegativities of the bonded atoms.
- Structure 2.2.6 Molecular polarity depends on both bond polarity and molecular geometry
- Structure 2.2.7 Carbon and silicon form covalent network structures
- Structure 2.2.8 The nature of the force that exists between molecules is determined by the size and polarity of the molecules.
- Structure 2.2.9 Given comparable molar mass, the relative strengths of intermolecular forces are generally: London (dispersion) forces < dipole–dipole forces < hydrogen bonding
- Structure 2.2.10 Chromatography is a technique used to separate the components of a mixture based on their relative attractions involving intermolecular forces to mobile and stationary phases.
- Structure 2.2.11 Resonance structures occur when there is more than one possible position for a double bond in a molecule.
- Structure 2.2.12 Benzene, C6H6, is an important example of a molecule that has resonance.
- Structure 2.2.13 Some atoms can form molecules in which they have an expanded octet of electrons.
- Structure 2.2.14 Formal charge values can be calculated for each atom in a species and used to determine which of several possible Lewis formulas is preferred
- Structure 2.2.15 Sigma bonds σ form by the head-on combination of atomic orbitals where the electron density is concentrated along the bond axis.
- Structure 2.2.16 Hybridization is the concept of mixing atomic orbitals to form new hybrid orbitals for bonding.
Structure 2.3 - The metallic model
- Reactivity 1.1.1 - Chemical reactions involve a transfer of energy between the system and the surroundings, while total energy is conserved.
- Reactivity 1.1.2 - Reactions are described as endothermic or exothermic, depending on the direction of energy transfer between the system and the surroundings.
- Reactivity 1.1.3 - The relative stability of reactants and products determines whether reactions are endothermic or exothermic.
Structure 2.4 - From models to materials
- Reactivity 1.2.1 - Bond-breaking absorbs and bond-forming releases energy.
- Reactivity 1.2.2 - Hess’s law states that the enthalpy change for a reaction is independent of the pathway between the initial and final states.
- Reactivity 1.2.3 - Standard enthalpy changes of combustion and formation are tabulated for a range of compounds.
- Reactivity 1.2.4 - Calculations involving standard enthalpy changes of reaction (ΔHθ) can be carried out using the relationship ΔHθ = Σ ΔHfθ(products) - Σ ΔHfθ(reactants).
Reactivity 1.5 - Ideal gases
- Reactivity 1.3.1 - Collision theory is used to explain how factors such as concentration, pressure and temperature affect the rate of chemical reactions.
- Reactivity 1.3.2 - The activation energy is the minimum energy required for a reaction to occur.
- Reactivity 1.3.3 - Catalysts provide an alternative reaction pathway with a lower activation energy, increasing the rate of reaction without being consumed.