1-States of matter 1-1 Everythings is made of particles 1-2 Solids, liquids, and gases 1.3 The particles in solids, Liquids, and gases 1.4 Heating and cooling curves 1.5 A closer look at gases

Chemistry IB _ Grade 11 and 12 Topic 2 : Atomic structure

Chemistry IB _ Grade 11 and 12 Topic 3 : Periodicity

Chemistry IB _ Grade 11 and 12 Topic 4 : Chemical bonding and structure

Chemistry IB _ Grade 11 and 12 Topic 5 : Energetics/thermochemistry

Chemistry IB _ Grade 11 and 12 Topic 6 : Chemical kinetics

Chemistry IB _ Grade 11 and 12 Topic 7 : Equilibrium

Chemistry IB _ Grade 11 and 12 Topic 8 : Acids and bases

Chemistry IB _ Grade 11 and 12 Topic 9 : Redox processes

Chemistry IB _ Grade 11 and 12 Topic 10 : Organic chemistry

Chemistry IB _ Grade 11 and 12 Topic 11 : Measurement and data processing

Chemistry IB _ Grade 11 and 12 Topic 1 : Stoichiometric relationships

Structure 2 / IB Chemistry / Structure 2.2 (lesson / Worksheets / Tests/ Tables / Figures) 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. Intermolecular forces include London (dispersion), dipole-induced dipole, dipole–dipole and hydrogen bonding. 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 / IB Chemistry / Structure 2.1 +worksheets +Formulae of common ions / ionic compounds Structure 2. Models of bonding and structure 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 1 / IB Chemistry / Structure 1.5 (Including worksheets) Structure 1.5- Ideal gases Structure 1.5.1 - An ideal gas consists of moving particles with negligible volume and no intermolecular forces. All collisions between particles are considered elastic. 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

Structure 1 / IB Chemistry / Structure 1.4 Structure 1.4—Counting particles by mass: The mole Structure 1.4.1 Structure 1.4.2 Structure 1.4.3 Structure 1.4.4 Structure 1.4.5 Structure 1.4.6

Structure 1.3 : Electron configurations 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. Structure 1.3.3 : Deduce the maximum number of electrons that can occupy each energy level. Structure 1.3.4 : Recognize the shape and orientation of an s atomic orbital and the three p atomic orbitals. 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. 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. Structure 1.3.7 : Deduce the group of an element from its successive ionization data.

Structure 1 / IB Chemistry / Structure 1.1 Structure 1.1 : Introduction to the particulate nature of matter Structure 1.1.1 : Distinguish between the properties of elements, compounds and mixtures. Structure 1.1.2 : Distinguish the different states of matter. Use state symbols (s, , 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 / IB Chemistry / Structure 1.2 Structure 1.2 : The nuclear atom Structure 1.2.1 : Use the nuclear symbol A Z X 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. Structure 1.2.3 : Interpret mass spectra in terms of identity and relative abundance of isotopes.

Reactivity 1.2 - Energy cycles in reactions 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, ΔHc ⦵, and formation, ΔHf ⦵, data are used in thermodynamic calculations. Reactivity 1.2.4 - An application of Hess’s law uses enthalpy of formation data or enthalpy of combustion data to calculate the enthalpy change of a reaction. Reactivity 1.2.5—A Born–Haber cycle is an application of Hess’s law, used to show energy changes in the formation of an ionic compound.