# Condensed Syllabus

## Structure 1. Models of the particulate nature of matter

### Structure 1.1 Introduction to the particulate nature of matter

**Guiding question**: How can we model 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.**
  * 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.
  * Distinguish between the properties of elements, compounds and mixtures.
  * Solvation, filtration, recrystallization, evaporation, distillation and paper chromatography should be covered. The differences between homogeneous and heterogeneous mixtures should be understood.
    * Tool 1—What factors are considered in choosing a method to separate the components of a mixture?
    * Tool 1—How can the products of a reaction be purified?
    * Structure 2.2—How do intermolecular forces influence the type of mixture that forms between two substances?
    * Structure 2.3—Why are alloys generally considered to be mixtures, even though they often contain metallic bonding?
* **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.**
  * Distinguish the different states of matter.
  * Use state symbols (s, , g and aq) in chemical equations.
  * Names of the changes of state should be covered: melting, freezing, vaporization (evaporation and boiling), condensation, sublimation and deposition.
    * Structure 2.4—Why are some substances solid while others are fluid under standard conditions?
    * Structure 2 (all), Reactivity 1.2—Why are some changes of state endothermic and some exothermic?
* **Structure 1.1.3—The temperature,&#x20;*****T*****, in Kelvin (K) is a measure of average kinetic energy E**<sub>**k**</sub>**&#x20;of**\
  **particles.**
  * Interpret observable changes in physical properties and temperature during changes of state.
  * Convert between values in the Celsius and Kelvin scales.
  * The kelvin (K) is the SI unit of temperature and has the same incremental value as the Celsius degree (°C).
    * Reactivity 2.2—What is the graphical distribution of kinetic energy values of particles in a sample at a fixed temperature?
    * Reactivity 2.2—What must happen to particles for a chemical reaction to occur?

### Structure 1.2 The nuclear atom

**Guiding question**: How do the nuclei of atoms differ?

* **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.**
  * Use the nuclear symbol $${}\_Z^AX$$ 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.
    * Structure 1.3—What determines the different chemical properties of atoms?
    * Structure 3.1—How does the atomic number relate to the position of an element in the periodic table?
* **Structure 1.2.2—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.
    * Nature of science, Reactivity 3.4—How can isotope tracers provide evidence for a reaction mechanism?
* **Structure 1.2.3—Mass spectra are used to determine the relative atomic masses of elements from**\
  **their isotopic composition.**
  * Interpret mass spectra in terms of identity and relative abundance of isotopes.
  * The operational details of the mass spectrometer will not be assessed.
    * Structure 3.2—How does the fragmentation pattern of a compound in the mass spectrometer help in the determination of its structure?

### Structure 1.3 Electron configurations

**Guiding question**: How can we model the energy states of electrons in atoms?

* **Structure 1.3.1—Emission spectra are produced by atoms emitting photons when electrons in**\
  **excited states return to lower energy levels.**
  * Qualitatively describe the relationship between color, wavelength, frequency and energy across the electromagnetic spectrum.
  * Distinguish between a continuous and a line spectrum.
  * 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.**
  * 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.
    * Inquiry 2—In the study of emission spectra from gaseous elements and of light, what qualitative and quantitative data can be collected from instruments such as gas discharge tubes and prisms?
    * Nature of science, Structure 1.2—How do emission spectra provide evidence for the existence of different elements?
* **Structure 1.3.3—The main energy level is given an integer number, n, and can hold a maximum of 2n**<sup>**2**</sup>**&#x20;electrons.**
  * Deduce the maximum number of electrons that can occupy each energy level.
    * Structure 3.1—How does an element’s highest main energy level relate to its period number in the periodic table?
* **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 3.1—What is the relationship between energy sublevels and the block nature of the\
      periodic table?
* **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.
  * Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron\
    configurations for atoms and ions up to Z = 36.
  * 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.
* **Structure 1.3.6—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 value of the Planck constant h and the equations E = hf and c = λf are given in the data booklet.
    * Structure 3.1—How does the trend in IE values across a period and down a group explain the trends in properties of metals and non-metals?
    * Nature of science, Tool 3, Reactivity 3.1—Why are log scales useful when discussing \[H+] and IEs?
* **Structure 1.3.7—Successive ionization energy (IE) data for an element give information about its electron configuration.**
  * Deduce the group of an element from its successive ionization data.
  * Databases are useful for compiling graphs of trends in IEs.
    * AHL Structure 3.1—How do patterns of successive IEs of transition elements help to explain the variable oxidation states of these elements?

### Structure 1.4 Counting particles by the mass: The mole

**Guiding question**: How do we quantify matter on the atomic scale?

* **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.**
  * Convert the amount of substance, n, to the number of specified elementary entities.
  * An elementary entity may be an atom, a molecule, an ion, an electron, any other particle or a specified group of particles.
  * The Avogadro constant N<sub>A</sub> is given in the data booklet. It has the units mol<sup>–1</sup>.
* **Structure 1.4.2—Masses of atoms are compared on a scale relative to&#x20;**<sup>**12**</sup>**C and are expressed as**\
  **relative atomic mass (A**<sub>**r**</sub>**) and relative formula mass (M**<sub>**r**</sub>**).**
  * Determine relative formula masses (M<sub>r</sub>) from relative atomic masses (A<sub>r</sub>).
  * Relative atomic mass and relative formula mass have no units.&#x20;
  * The values of relative atomic masses given to two decimal places in the data booklet should be used in calculations.
    * Structure 3.1—Atoms increase in mass as their position descends in the periodic table. What properties might be related to this trend?
* **Structure 1.4.3—Molar mass (M) has the units g mol**<sup>**–1**</sup>**.**
  * Solve problems involving the relationships between the number of particles, the amount of substance in moles and the mass in grams.
  * The relationship $$n=\frac{m}{M}$$ is given in the data booklet.
    * Reactivity 2.1—How can molar masses be used with chemical equations to determine the masses of the products of a reaction?
* **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.**
  * Interconvert the percentage composition by mass and the empirical formula.
  * Determine the molecular formula of a compound from its empirical formula and molar mass.
    * Tool 1—How can experimental data on mass changes in combustion reactions be used to derive empirical formulas?
    * Nature of science, Tool 3, Structure 3.2—What is the importance of approximation in the determination of an empirical formula?
* **Structure 1.4.5—The molar concentration is determined by the amount of solute and the volume of solution.**
  * Solve problems involving the molar concentration, amount of solute and volume of solution.
  * The use of square brackets to represent molar concentration is required.
  * Units of concentration should include g dm<sup>–3</sup> and mol dm<sup>–3</sup> and conversion between these.
  * The relationship n = CV is given in the data booklet.
    * Tool 1—What are the considerations in the choice of glassware used in preparing a standard solution and a serial dilution?
    * Tool 1, Inquiry 2—How can a calibration curve be used to determine the concentration of a 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.**
  * Solve problems involving the mole ratio of reactants and/or products and the volume of gases.
    * Structure 1.5—Avogadro’s law applies to ideal gases.
    * Under what conditions might the behaviour of a real gas deviate most from an ideal gas?

### Structure 1.5 Ideal gases

**Guiding question**: How does the model of ideal gas behavior help us to predict the behavior of real\
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.**
  * Recognize the key assumptions in the ideal gas model.
* **Structure 1.5.2—Real gases deviate from the ideal gas model, particularly at low temperature and high pressure.**
  * Explain the limitations of the ideal gas model.
  * No mathematical coverage is required
    * Structure 2.2—Under comparable conditions, why do some gases deviate more from ideal behavior than others?
* **Structure 1.5.3—The molar volume of an ideal gas is a constant at a specific temperature and**\
  **pressure.**
  * Investigate the relationship between temperature, pressure and volume for a fixed mass of an ideal gas and analyze graphs relating these variables.
  * The names of specific gas laws will not be assessed.
  * The value for the molar volume of an ideal gas under standard temperature and pressure (STP) is given in the data booklet.
    * Nature of science, Tools 2 and 3, Reactivity 2.2—Graphs can be presented as sketches or as accurately plotted data points. What are the advantages and limitations of each representation?
* **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** $$\frac{P\_1V\_1}{T\_1}=\frac{P\_2V\_2}{T\_2}.$$
  * Solve problems relating to the ideal gas equation.
  * Units of volume and pressure should be SI only. The value of the gas constant R, the ideal gas equation, and the combined gas law, are given in the data booklet.
    * Tool 1, Inquiry 2—How can the ideal gas law be used to calculate the molar mass of a gas from experimental data?

## Structure 2. Models of bonding and structure

### Structure 2.1 The ionic model

**Guiding question**: What determines the ionic nature and properties of a compound?

* **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.
  * 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.
    * Structure 3.1—How does the position of an element in the periodic table relate to the charge of its ion(s)?
    * AHL Structure 1.3—How does the trend in successive ionization energies of transition elements explain their variable oxidation states?
* **Structure 2.1.2—The ionic bond is formed by electrostatic attractions between oppositely charged ions.**
  * Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions.
  * Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix “ide”.
  * Interconvert names and formulas of binary ionic compounds.
  * The following polyatomic ions should be known by name and formula: ammonium NH<sup><sub>4<sub></sup><sup>+</sup>, hydroxide OH<sup>–</sup>, nitrate NO<sub>3</sub><sup>–</sup>, hydrogencarbonate HCO<sub>3</sub><sup>–</sup>, carbonate CO<sub>3</sub><sup>2–</sup>, sulfate SO<sub>4</sub><sup>2–</sup>, phosphate PO<sub>4</sub><sup>3–</sup>.
    * Reactivity 3.2—Why is the formation of an ionic compound from its elements a redox reaction?
    * AHL Structure 2.2—How is formal charge used to predict the preferred structure of sulfate?
    * AHL Reactivity 3.1—Polyatomic anions are conjugate bases of common acids. What is the relationship between their stability and the conjugate acid’s dissociation constant, K<sub>a</sub>?
* **Structure 2.1.3—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.
    * Tool 1, Inquiry 2—What experimental data demonstrate the physical properties of ionic compounds?
    * Structure 3.1—How can lattice enthalpies and the bonding continuum explain the trend in melting points of metal chlorides across period 3?

### Structure 2.2 The covalent model

**Guiding question**: What determines the covalent nature and properties of a substance?

* **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. Deduce the Lewis formula of molecules and ions for up to four electron pairs on each atom.
  * 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. Molecules containing atoms with fewer than an octet of electrons should be covered.
  * Organic and inorganic examples should be used.
    * Nature of science—What are some of the limitations of the octet rule?
    * Structure 1.3—Why do noble gases form covalent bonds less readily than other elements?
    * Structure 2.1—Why do ionic bonds only form between different elements while covalent bonds can form between atoms of the same element?
* **Structure 2.2.2—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.
    * Reactivity 2.2—How does the presence of double and triple bonds in molecules influence their reactivity?
