Chemistry

Core Declarative Knowledge
What should students know?

Distinguish between the properties of elements, compounds and mixtures.
Distinguish the different states of matter.
Interpret observable changes in physical properties and temperature during changes of state.
Isotopes are atoms of the same element with different numbers of neutrons.
Mass spectra are used to determine the relative atomic masses of elements from their isotopic composition.
Emission spectra are produced by atoms emitting photons when electrons in
excited states return to lower energy levels.
The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.
The main energy level is given an integer number, n, and can hold a maximum of 2n2 electrons.
A more detailed model of the atom describes the division of the main energy level into s, p, d and f sublevels of successively higher energies.
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, regions of space where there is a high probability of finding an electron.
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.
Masses of atoms are compared on a scale relative to 12C and are expressed as relative atomic mass Ar and relative formula mass Mr.
Moles calculations.
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.
The molar concentration is determined by the amount of solute and the volume of solution.
Avogadro’s law states that equal volumes of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules.
An ideal gas consists of moving particles with negligible volume and no intermolecular forces. All collisions between particles are considered elastic.
Real gases deviate from the ideal gas model, particularly at low temperature and high pressure.
The molar volume of an ideal gas is a constant at a specific temperature and pressure.
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

Core Procedural Knowledge
What should students be able to do?

Use state symbols (s, , g and aq) in chemical equations.
Different separation techniques like Solvation, filtration, recrystallization, evaporation,
distillation and paper chromatography
Remember different changes of states like melting, freezing, vaporization (evaporation and
boiling), condensation, sublimation and deposition.
Convert between values in the Celsius and Kelvin scales.
Use the nuclear symbol to deduce the number of protons, neutrons and electrons in atoms and ions.
Perform calculations involving non-integer relative atomic masses and abundance of isotopes from given data.
Interpret mass spectra in terms of identity and relative abundance of isotopes.
Qualitatively describe the relationship between colour, wavelength, frequency and energy across the
electromagnetic spectrum.
Distinguish between a continuous and a line spectrum.
Describe the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.
Deduce the maximum number of electrons that can occupy each energy level.
Recognize the shape and orientation of an s atomic orbital and the three p atomic 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.
Convert the amount of substance, n, to the number of specified elementary entities.
Determine relative formula masses Mr from relative atomic masses Ar.
Solve problems involving the relationships between the number of particles, the amount of substance in moles and the mass in grams.
Interconvert the percentage composition by mass and the empirical formula.
Determine the molecular formula of a compound from its empirical formula and molar mass.
Solve problems involving the molar concentration, amount of solute and volume of solution.
Solve problems involving the mole ratio of reactants and/or products and the volume of gases.
Recognize the key assumptions in the ideal gas model.
Investigate the relationship between temperature, pressure and volume for a fixed mass of an ideal gas and analyse graphs relating these variables.
Solve problems relating to the ideal gas equation.

Core Declarative Knowledge
What should students know?

Structure 1.1.1—Elements are the primary constituents of matter, which 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, which are not chemically bonded and so can be separated by physical methods.

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. Names of the changes of state should be covered: melting, freezing, vaporisation (evaporation and boiling), condensation, sublimation and deposition.

Structure 1.1.3—The temperature, T, in Kelvin (K) is a measure of average kinetic energy (Ek) of particles. The kelvin (K) is the SI unit of temperature and has the same incremental value as the Celsius degree (°C).

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. 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.

Structure 1.2.2—Isotopes are atoms of the same element with different numbers of neutrons. Specific examples of isotopes need not be learned. AHL

Structure 1.2.3—Mass spectra are used to determine the relative atomic masses of elements from their isotopic composition. The operational details of the mass spectrometer will not be assessed.

Structure 1.3.1—Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels. Details of the electromagnetic spectrum are given in the data booklet.

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 sublevels of successively higher energies. Recognize the shape and orientation of an s atomic orbital and the three p atomic orbitals.

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. Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron. 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.

AHL Structure 1.3.6—In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization.

AHL Structure 1.3.7—Successive ionization energy (IE) data for an element give information about its electron configuration.

Core Procedural Knowledge
What should students be able to do?

Structure 1.1.1 -Distinguish between the properties of elements, compounds and mixtures. Solvation, filtration, recrystallization, evaporation, distillation and chromatography should be covered. The differences between homogeneous and heterogeneous mixtures should be understood.

Structure 1.1.2 – Distinguish the different states of matter. Use state symbols (s, l, g and aq) in chemical equations.

Structure 1.1.3 – Interpret observable changes in physical properties and temperature during changes of state. Convert between values in the Celsius and Kelvin scales.

Structure 1.2.1— Use the nuclear symbol XZA to deduce the number of protons, neutrons and electrons in atoms and ions.

Structure 1.2.2—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.

AHL Structure 1.2.3—Interpret mass spectra in terms of identity and relative abundance of isotopes.

Structure 1.3.1—Qualitatively describe the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum. Distinguish between a continuous and a line spectrum.

Structure 1.3.2—Describe the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels. The names of the different series in the hydrogen emission spectrum will not be assessed.

Structure 1.3.3—Deduce the maximum number of electrons that can occupy each energy level.

Structure 1.3.4—N/A

Structure 1.3.5—Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron configurations for atoms and ions up to Z=36.

AHL Structure 1.3.6— 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.

AHL Structure 1.3.7—Deduce the group of an element from its successive ionization data.

Core Declarative Knowledge
What should students know?

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. Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix “ide”.

Structure 2.1.3—Ionic compounds exist as three-dimensional lattice structures, represented by empirical formulas.

Structure 2.2.1—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.Lewis formulas (also known as electron dot or 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.

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. Include predicting how non-bonding pairs and multiple bonds affect bond angles.

Structure 2.2.5—Bond polarity results from the difference in electronegativities of the bonded atoms. Bond dipoles can be shown either with partial charges or vectors.

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. Allotropes of the same element have different bonding and structural patterns, and so have different chemical and physical properties.

Structure 2.2.8—The nature of the force that exists between molecules is determined by the size and polarity of the molecules. Intermolecular forces include London (dispersion), dipole-induced dipole, dipole–dipole and hydrogen bonding. The term “van der Waals forces” should be used as an inclusive term to include dipole–dipole, dipole-induced dipole, and London (dispersion) forces. Hydrogen bonds occur when hydrogen, being covalently bonded to an electronegative atom, has an attractive interaction on a neighbouring electronegative atom.

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 2.3.1—A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons.

Structure 2.3.2—The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion.

Core Procedural Knowledge
What should students be able to do?

Structure 2.1.1—Predict the charge of an ion from the electron configuration of the atom.

Structure 2.1.2—Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. Interconvert names and formulas of binary ionic compounds.

Structure 2.1.3— Explain the physical properties of ionic compounds to include volatility, electrical conductivity and solubility.

Structure 2.2.1—Deduce the Lewis formula of molecules and ions for up to four electron pairs on each atom.

Structure 2.2.2—Explain the relationship between the number of bonds, bond length and bond strength.

Structure 2.2.3— Identify coordination bonds in compounds.

Structure 2.2.4— Predict the electron domain geometry and the molecular geometry for species with up to four electron domains.

Structure 2.2.5— Deduce the polar nature of a covalent bond from electronegativity values.

Structure 2.2.6— Deduce the net dipole moment of a molecule or ion by considering bond polarity and molecular geometry.

Structure 2.2.7— Describe the structures and explain the properties of silicon, silicon dioxide and carbon’s allotropes: diamond, graphite, fullerenes and graphene.

Structure 2.2.8— Deduce the types of intermolecular force present from the structural features of covalent molecules.

Structure 2.2.9— Explain the physical properties of covalent substances to include volatility, electrical conductivity and solubility in terms of their structure.

Structure 2.3.1— Explain the electrical conductivity, thermal conductivity and malleability of metals.Relate characteristic properties of metals to their uses.

Structure 2.3.2— Explain trends in melting points of s and p block metals.