Structure 1: Models of the Particulate Nature of Matter
| Subtopic | Subtopic Number | IB Points to Understand |
| Particulate matter | 1.1 (SL) | - Elements– Pure substance made of only one type of atom.
- Compounds– Substance formed when two or more different elements are chemically bonded together in fixed portions.
- Mixtures– Combination of two or more substances (elements, compounds, or both) that are physically, not chemically, bonded.
- Differentiate between the states of matter (Solid, liquid, gas).
- Matter is defined as anything that has mass and takes up space (it has volume).
- Properties of three states of matter (Solids, Liquids and Gasses)
- Temperature is a measure of the average kinetic energy of the particles of a substance.
- Sublimation– When solid turns into a gas without first becoming liquid.
- Condensation– When gaseous state changes into liquid state.
- Use symbols of states of matter in chemical equations (s, l, g).
- Interconversion between the states of the matter.
- Observable changes in physical properties and temperature during changes in state.
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| The Atom | 1.2 (SL) | - Structure of the atom
- Use of nuclear notation
- Deduce the number of protons, neutrons and electrons in atoms and ions.
- Isotopes- atoms of the same element that have the same number of protons, but different number of neutrons in the nuclei.
- Isotope abundance calculation.
- Non-integer relative atomic masses.
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| Electronic configuration | 1.3 (SL) | - Emission spectra- spectrum of light emitted by an atom or molecule when electrons from an excited state move to a ground state.
- The hydrogen line emission spectrum consists of a series of lines of different colors in the visible region of the spectrum.
- Maximum number of electrons that can occupy each energy level.
- Shapes and orientations of S and P orbitals.
- Aufbau Principle- filling out the electronic configuration of an atom by adding electrons to lowest energy orbitals before filing higher energy orbitals.
- Hund’s rule- All orbitals will be singly occupied before any orbital is doubly occupied.
- Pauli’s exclusion principle- no two electrons can be identical within an atom. Simply stated, it means that only two electrons can fit into each atomic orbital and they must have opposite spins.
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| Moles (SL) | 1.4 (SL) | - The mole is an SI unit, symbol mol, defined as a fixed amount, n, of a substance
- The relative atomic mass Ar of an atom is a weighted average of the atomic masses of its isotopes and their relative abundances
- The empirical formula of a compound is the simplest whole-number ratio of atoms or amount (in mol) of each element present in a compound.
- The molecular formula is the actual number of atoms or amount (in mol) of elements in one structural unit or one mole of the compound, respectively.
- The reactant that determines the quantity of product is known as the limiting reactant product.
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| Ideal Gases | 1.5 (SL) | - The kinetic theory of gasses is a model used to explain and predict the behavior of gasses at a microscopic level
- Standard temperature (0 °C/273 K) and Standard pressure (100kPa)
- Boyle’s Law: Relationship between volume and pressure
- Charles’s Law: Relationship between volume and temperature
- Gay-Lussac’s Law: Relationship between pressure and temperature
- Ideal Gas Equation: PV = nRT
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| Further Atoms | 11.1 (HL) | - Mass spectra
- Relative abundance of isotopes.
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| Further electron configurations | 11.2 (HL) | - Deduce groups of elements from ionization energy data.
- Trends in ionization energy across a period and down the groups.
- Convergence limit and its relationship to ionization energy.
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Structure 2: Models of bonding and Structure:
| Subtopic | Subtopic Number | IB Points to Understand |
| Ionic bonds & structure | 2.1 (SL) | - An ionic Bond refers to the electrostatic attraction experienced between the electric charges of a cation (positive ion) and an anion (negative ion).
- Electrons positioned outside the nucleus are less tightly held and outer electrons, known as valence electrons, can be transferred when atoms react together.
- Predict the charge of ions using electronic configuration, including transition elements.
- Ionic compounds- These forces are known as an ionic bond, and ions held together in this way.
- Electrostatic forces- The oppositely charged ions resulting from electron transfer are attracted to each other and are held together by this force.
- Physical properties of ionic compounds.
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| Covalent bonding and forces | 2.2 (SL) | - Covalent Bonding- atoms share electrons with each other in order to attain a noble gas electron configuration.
- In a Lewis symbol representation, each element is surrounded by a number of dots (or crosses), which represent the valence electrons of the element.
- Bond strength, bond length, bond order.
- Bonding pairs of electrons (showing the covalent bond as single, double, or triple bonds) and
- Non-bonding pairs of electrons, often called the lone pairs, which are pairs of electrons not involved in the bonding
- Valence shell electron pair repulsion (VSEPR) theory can be used to deduce the shapes of covalent molecules
- Molecules with two electron domains will position them at 180° to each other
- Molecules with three electron domains will position them at 120° to each other
- Molecules with four electron domains will position them at 109.5° to each other
- Resonance involves using two or more Lewis structures to represent a particular molecule or ion
- Difference between graphite, diamond, fullerene and graphene
- Compare the strengths of different intermolecular forces.
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| Metallic bonds | 2.3 (SL) | - Delocalized: when there is no other element present to accept the electrons and form an ionic compound, the outer electrons are held only loosely by the metal atom’s nucleus.
- Strength of a metallic bond is directly proportional to the charge on the ion, and inversely proportional to the radius of the ion.
- Properties of metals, and experimental uses of metals.
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| Models to materials | 2.4 (SL) | - Bonding Triangle.
- Bonding model to explain the properties of materials.
- An alloy is a mixture that consists either of two or more metals, or of a metal (or metals) combined with an alloying element composed of one or more nonmetals.
- Polymers- are macromolecules made from repeating subunits known as monomers.
- Formation of addition polymers.
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| Further covalent bonding | 12.1 (HL) | - Resonance structures.
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- Lewis structures with 5 and 6 domains around the central atom.
- Sigma bonds (σ) and pi bonds (π).
- Define hybridization as the concept of mixing atomic orbitals to form new hybrid orbitals for bonding. Analyze the hybridization and bond formation in molecules and ions. Identify the relationships between Lewis formulas, electron domains, molecular geometry and type of hybridization. Predict the geometry around an atom from its hybridization, and vice versa.
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| Further metallic bonding | 12.2 (HL) | - Transition elements have delocalised d-electrons.
- Explain why transition elements have high melting points and high electrical conductivity.
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| Condensation polymers | 12.3 (HL) | - Reactions involved in the formation of condensation polymers.
- Representations of monomer structures.
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Structure 3: Classification of Matter
| Subtopic | Subtopic Number | IB Points to Understand |
| The Periodic Table & Trends | 3.1 (SL) | - Vertical columns: Groups
- The Horizontal rows of elements numbered from 1 to 7 are termed period
- Period number is equal to the principal quantum number, n
- Effective nuclear charge
- Atomic radius is the distance from the center of the nucleus to the outermost shell containing electrons.
- The formation of positive ions involves the loss of the outer shell. The formation of negative ions involves the addition of electrons into the outer shell.
- The ionization energy, IE, is the minimum energy required to remove an electron from a neutral gaseous atom in its ground-state.
- The first electron affinity of an element (∆Hea) is the energy change when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous ions.
- Electronegativity, symbol χ, is defined as the relative attraction that an atom has for the shared pair of electrons in a covalent bond.
- Group 1, 17 and 18 chemical properties.
- Reactions of Group 1 metals and Group 17 halides.
- Understanding oxidation states.
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| Functional Groups | 3.2 (SL) | - A homologous series is a series of compounds that can be grouped together based on similarities in their structure and reactions. Members of a homologous series can be represented by the same general formula.
- Trends in melting and boiling points in the homologous series.
- Full structural formula, condensed structural formula and skeletal formula.
- The alkenes and alkynes are two more hydrocarbon homologous series that contain carbon– carbon double and triple bonds, respectively.
- Identify the following functional groups by name and structure: halogeno, hydroxyl, carbonyl, carboxyl, alkoxy, amino, amido, ester, phenyl. Understand and use the terms “saturated” and “unsaturated”.
- IUPAC nomenclature.
- Molecules, having the same molecular formula but different arrangements of the atoms, are known as structural isomers.
- Identify straight-chain, position and functional group isomers, and primary, secondary and tertiary alcohols.
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| Further periodic table | 13.1 (HL) | - Explain the discontinuities in the trend of increasing first ionization energy across a period with reference to the energy of the electron removed.
- Existence of energy levels.
- A transition element is an element that has an atom with an incomplete d-sublevel or that gives rise to cations with an incomplete d-sublevel.
- Characteristics of transition elements.
- Transition metals are often found with different oxidation states.
- Transition metals appear coloured because they absorb visible light.
- Transition metals absorb light because the d orbitals split into two sub-levels.
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| Further functional groups | 13.2 (HL) | - Stereoisomers have an identical molecular formula and bond multiplicity but show different spatial arrangements of the atoms.
- Stereoisomers can be subdivided into two major classes, conformational isomers and configurational isomers.
- Conformational isomers therefore differ from one another in the arrangement of atoms around a single bond.
- Configurational isomers can be interconverted only by the breaking of bonds or through the rearrangement of the stereocenters.
- A high-resolution 1HNMR spectrum can show further splitting of some absorptions.
- Splitting patterns result from spin–spin coupling.
- Single X-ray Crystallography is a scientific method used to determine the arrangement of atoms of a crystalline solid in three dimensional space.
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Reactivity 1: What drives chemical reactions
| Subtopic | Subtopic Number | IB Points to Understand |
| Enthalpy | 4.1 (SL) | - Enthalpy (H) is a measure of the amount of heat energy contained in a substance. It is stored in the chemical bonds and intermolecular forces as potential energy.
- Enthalpy (H) is a measure of the amount of heat energy contained in a substance. It is stored in the chemical bonds and intermolecular forces as potential energy.
- Exothermic reactions give out heat and result in a transfer of enthalpy from the chemicals to the surroundings and ∆H reaction is negative.
- A few reactions are endothermic as they result in an energy transfer from the surroundings to the system. In this case the products have more enthalpy than the reactants and ∆H is positive.
- Energy profile diagrams for endothermic and exothermic reactions.
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| Cycles of Energy | 4.2 (SL) | - Exothermic (bond formic) reactions release energy.
- Endothermic (bond breaking) reactions absorb energy.
- Calculate enthalpy change of reaction from average bond enthalpy data.
- The bond enthalpy is the energy needed to break one mole of bonds in gaseous molecules under standard condition.
- Hess’s Law states that the enthalpy change for a chemical reaction is independent of the route taken.
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| Fuels | 4.3 (SL) | - Equations for reactions of combustion, including hydrocarbons and alcohols.
- Equations for reactions of incomplete combustion, including hydrocarbons and alcohols.
- Fossil fuels and the greenhouse effect.
- Production of biofuels through biological fixation.
- Pros and cons of biofuels.
- Non-renewable and renewable energy sources.
- Deduce half-equations for the electrode reactions in a fuel cell, including Hydrogen and methanol as fuels.
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| Further energy cycles | 14.4 (HL) | |
| Entropy and spontaneity | 14.2 (HL) | - The combination of enthalpy, entropy, and temperature of the system can be used to define a new state function called Gibbs free energy.
- The Gibbs free energy change of formation, ΔGf, represents the free energy change when 1 mol of a compound is formed from its elements under standard conditions of 298 K and a pressure of 100 kPa.
- Define entropy, S, as a measure of the dispersal or distribution of matter and/or energy in a system. Predict whether a physical or chemical change will result in an increase or decrease in entropy of a system. Calculate standard entropy changes, ΔS⦵, from standard entropy values, S⦵.
- Perform calculations using the equation ΔG = ΔG⦵ + RT lnQ and its application to a system at equilibrium ΔG⦵ = −RT lnK.
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Reactivity 2: How much, how fast and how far?
| Subtopic | Subtopic Number | IB Points to Understand |
| Amount of change | 5.1 (SL) | - Chemical reactions using symbols states of matter.
- The reactant that determines the quantity of product is known as the limiting reactant.
- The theoretical yield refers to the maximum amount of product obtainable, assuming 100% of the limiting reactant is converted to product.
- Atom economy is a measure of efficiency in green chemistry.
- Use the mole ratio of an equation to calculate: the masses, volumes, and concentrations of reactants and products. Use Avogadro’s law and definitions of molar concentration.
- The mole is an SI unit, symbol mol, defined as a fixed amount, n, of a substance.
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| Rate of change | 5.2 (SL) | - The rate of a reaction depends on how quickly the concentration of either reactant or product changes with respect to time.
- Collision theory.
- Factors affecting rate of reaction.
- In order for a collision to lead to reaction, the particles must have a certain minimum value for their kinetic energy, known as the activation energy, Ea.
- Maxwell Boltzman curve on the effect of temperature on successful collisions of particles.
- Construct Maxwell–Boltzmann energy distribution curves to explain the effect of different values for Ea on the probability of successful collisions.
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| Extent of change | 5.3 (SL) | - A state of equilibrium is when both the forward and reverse reactions occur simultaneously with products and reactants constantly being interconverted.
- Dynamic equilibrium occurs in a closed system.
- The law of chemical equilibrium states that at a given temperature the ratio of the concentration of products to the concentration of products to the concentration of reactants is a constant, Kc.
- The effects of changing experimental conditions on the equilibrium constant.
- Le Châtelier’s principle.
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| Further rates of change | 15.1 (HL) | - Reactions may occur by more than one step and the slow step determines the rate of the reaction. The slow step is termed the rate-determining step (RDS).
- Zero to third order reaction.
- Evaluate proposed reaction mechanisms and recognize reaction intermediates. Distinguish between intermediates and transition states, and recognize both in energy profiles of reactions.
- Construct and interpret energy profiles from kinetic data.
- Rate equation from experimental data.
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| Further extent of change | 15.2 (HL) | - Reaction quotient (Q).
- Solve problems involving values of K and initial and equilibrium concentrations of the components of an equilibrium mixture. Understand the approximation that [reactant]initial ≈ [reactant]eqm when K is very small.
- Use of ΔG⦵ = −RT lnK
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Reactivity 3: What are the Mechanics of Chemical Change
| Subtopic | Subtopic Number | IB Points to Understand |
| Proton Transfer | 6.1 (SL) | - Bronsted-Lowry- the theory defines “an acid as a proton donor and a base as a proton acceptor”
- If the acid dissociates fully, it will exist entirely as ions in solution. It is said to be a strong acid.
- If the acid dissociates only partially, it produces an equilibrium mixture in which the undissociated form dominates. It is said to be a weak acid.
- Distinguishing between strong and weak acids and bases using electrical conductivity, rate of reaction and their pH.
- The stronger the acid, the larger the Ka
- The weaker the acid, the larger the pKa
- The stronger the base, the larger the Kb.
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| Electron transfer | 6.2 (SL) | - Oxidation describes a process in which the oxidation state increases and reduction describes a process in which the oxidation state decreases.
- Oxidation and reduction will always occur together and reactions of this type are known as redox reactions.
- Oxidation state is the apparent charge of an atom in a free element, a molecule, or an ion.
- Half-equations can be very useful in balancing complex redox reactions. Each half-equation represents the separate oxidation and reduction processes.
- The reactant that accepts electrons is called the oxidizing agent as it brings about oxidation of the other reactant. In the process it becomes reduced.
- Likewise the reactant that supplies the electrons is known as the reducing agent, because it brings about reduction and itself becomes oxidized.
- Voltaic (or galvanic) cells – these convert chemical energy to electrical energy. Voltaic cells convert energy from spontaneous, exothermic chemical processes to electrical energy.
- Electrolytic cells – these convert electrical energy to chemical energy, bringing about a
- non- spontaneous process.
- An electrode is a conductor of electricity used to make contact with a non-metallic part of a circuit, such as the solution in a cell
- The electrode where oxidation occurs is called the anode. The electrode where reduction occurs is called the cathode.
- Deduce equations to show reduction of carboxylic acids to primary alcohols via the aldehyde, and reduction of ketones to secondary alcohols. Include the role of hydride ions in the reduction reaction. The names and formulas of specific reducing agents, and the mechanisms of reduction, will not be assessed.
- Recall that the reduction of unsaturated compounds by the addition of hydrogen lowers the degree of unsaturation. Deduce the products of the reactions of hydrogen with alkenes and alkynes.
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| Electron sharing | 6.3 (SL) | - Free-radical refers to a species that is formed when a molecule undergoes homolytic fission: the two electrons of a covalent bond are split evenly between two atoms resulting in two free-radicals that each have a single electron
- Initiation, Propagation and Termination
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| Electron-pair sharing | 6.4 (SL) | - Heterolytic fission of a bond creates a cation and an anion, as the electrons involved in the bond are unevenly split between the two atoms.
- Nucleophilic substitution reactions.
- The electron-deficient carbon is attacked by electron-rich species known as nucleophiles.
- An electrophile is an electron-poor species capable of accepting an electron pair. It acts as a Lewis acid.
- Explain why alkenes are susceptible to electrophilic attack. Deduce equations for the reactions of alkenes with water, halogens, and hydrogen halides.
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| Further proton transfer | 16.1 (HL) | - A buffer is a solution that resists a change in pH upon the addition of small amounts of a strong base or strong acid, or upon the dilution of the buffer through the addition of water
- The change in pH does not show a linear relationship with the volume of base added, partly due to the logarithmic nature of the pH scale
- Titration of a strong acid with a strong base: pH at equivalence = 7
- Titration of a weak acid with a strong base: pH at equivalence > 7
- Titration of a strong acid with a weak base: pH at equivalence < 7
- Titration of a weak acid with a weak base: pH at equivalence – difficult to identify
- An indicator is typically a weak acid or a weak base that displays a different color in acidic or alkaline environments
- Different indicators must be used for different titrations, depending on the pH at the equivalence point.
- Calculating acid dissociation constant Ka and the base dissociation constant Kb.
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| Further electron transfer | 16.2 (HL) | - Electroplating – It is a process of deposition of a thin layer of a metal over another using an electrolytic process.
- A Standard Hydrogen Electrode (SHE) is an electrode that scientists use for reference on all half-cell potential reactions.
- Explain the effects of concentration and the nature of the electrode for electrolysis of NaCl(aq) and CuSO4(aq).
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| Further electron-pair sharing | 16.3 (HL) | - Compounds that contain transition elements and in which the central metal ion is bonded, via coordinate bonding, to a group of molecules or ions (termed the ligands) are termed transition metal complexes.
- Ligands classification: Monodentate and Polydentate.
- SN2 reactions and primary halogenoalkanes: Nucleophilic substitution in primary halogenoalkanes proceed in one step. The rate- determining step involves both the halogenoalkane and the nucleophile.
- Drawing mechanisms for SN2 reactions
- SN1 reactions and tertiary halogenoalkanes: Tertiary halogenoalkanes undergo nucleophilic substitution reactions that involve two steps. The rate-determining step involves only the halogenoalkane.
- Drawing mechanisms for SN1 reactions.
- Reactions involving symmetrical and unsymmetrical alkenes.
- Electrophilic substitution reactions and its drawing mechanism.
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Frequently Asked Questions (FAQs)
Q1: What topics are covered in the IB Chemistry syllabus?
A: The IB Chemistry syllabus covers a wide range of topics, including atomic structure, bonding, energetics, kinetics, equilibrium, acids and bases, oxidation and reduction, organic chemistry, and analytical chemistry.
Q2: How is the IB Chemistry course assessed?
A: The IB Chemistry course is assessed through a combination of internal and external assessments. Internal assessments include lab work, field work, and other practical activities, while external assessments include written exams and a scientific investigation.
Q3: What skills do students need to succeed in IB Chemistry?
A: Students in IB Chemistry need to have strong analytical, problem-solving, and experimental design skills. They should also be able to communicate scientific ideas clearly and accurately, and be comfortable with data analysis and statistical analysis.
Q4: How is the IB Chemistry course different from other high school chemistry courses?
A: The IB Chemistry course is designed to be more rigorous and in-depth than other high school chemistry courses. It emphasizes a conceptual understanding of chemical concepts, as well as practical skills such as experimental design, data analysis, and scientific communication.
Q5: What are the benefits of taking the IB Chemistry course?
A: Taking the IB Chemistry course can provide students with a strong foundation in chemistry that will prepare them for future studies in the field. It can also help students develop critical thinking and problem-solving skills that are valuable in a wide range of careers. Additionally, the IB Chemistry course is recognized by colleges and universities around the world, which can be beneficial for college admissions.
