IB Chemistry

Subject Specific Core (SSC) Material 

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Subject Specific Core Material (SSC)

Topic 1 - Stoichiometry

1.1 Mole Concept and Avogadro's constant

1.1.1 Describe the mole concept, and apply it to substances

1.1.2 Perform calculations involving number of particles and moles of a substance


1.2 Formula

1.2.1 Define the term molar mass and calculate the mass of one mole of a species.

1.2.2 Distinguish between atomic mass, molecular mass and formula mass. (g mol-1)

1.2.3 Define the term "relative molecular mass" and "relative atomic mass" (no units)

1.2.4 State the relationship between the amount of a substance (in moles) and mass, and carry out calculations involving amount of substance, mass and molar mass.

1.2.5 Define the terms "empirical formula" and "molecular formula".

1.2.6 determine the empirical and or molecular formula of a given compound.

1.3 Chemical Equations

1.3.1 Balance simple chemical equations when all reactants and products are given.

1.3.2 Identify the mole ratios of any two species in a balanced chemical equation

1.3.3 Apply the state symbols, (s), (l), (g), and (aq).


1.4 Mass and Gaseous Volume Relationships in chemical reactions

1.4.1 Calculate stoichiometric quantities and use these to determine experimental and theoretical yields.

1.4.2 Determine the limiting reagent and the reagent in excess when quantities of two reacting substances are given.

Given a chemical equation and the initial amounts of two or more reactants:

1.4.3 Apply Avagadro's law to calculate the reacting volume of gases


1.5 Solutions

1.5.1 Define the terms 'solute', 'solvent', and concentration (g dm-3 and mole dm-3)

1.5.2 Carry out calculations involving concentration, amount of solute and volume of solution.

1.5.3 Solve simple solution stoichiometry problems.


SSC Topic 2 - Atomic Theory

2.1 The Atom

2.1.1 State the relative mass and relative charge of protons, electrons, and neutrons.

2.1.2 State the relative position of protons, neutrons and electrons in the nucleus and shells

2.1.3 Define the terms 'mass number' (A) and 'atomic number' (Z) and "isotope'.

2.1.4 State the symbol for an isotope given its mass number and atomic number

2.1.5 Explain how the isotopes of an element differ. Isotopes of hydrogen and carbon should be considered.

2.1.6 Calculate and explain non-integer atomic masses from the relative abundance of isotopes.

2.1.7 Calculate the number of protons, electrons, and neutrons in atoms and ions from mass number, atomic number and charge.


2.2 Electron Arrangement

2.2.1 Describe and explain the difference between a continuous spectrum and a line spectrum

2.2.2 Explain how the lines in an emission spectrum are related to the energy levels of electrons.

2.2.3 Describe electron arrangement of atoms in terms of the main levels. Up to Z=18.

2.2.4 Apply the notation for electronic configuration up to Z = 18.


SSC Topic 3 - Periodicity


3.1 The Periodic Table

3.1.1 Describe the arrangement of the elements in the Periodic Table in order of increasing number.

3.1.2 Distinguish between the terms "group"' and "period".

3.1.3 Deduce the relationship between the electron configuration of elements and their position in the Periodic Table.


3.2 Physical Properties

3.2.1 Describe and explain the periodic trends in atomic radii, ionization energies, affinities, electronegativity and melting point for the alkali metals (Li--> Cs), halogens (F-->I) and period 3 (Na-->Ar).

3.3 Chemical Properties

3.3.1 Discuss the similarities in chemical nature of elements in the same group.

The following reactions should be covered:

 3.3.2 Discuss the change in nature, from metallic to non-metallic, of the elements across period 3.


SSC Topic 4 - Bonding

4.1 Ionic bond

4.1.1 Describe the ionic bond as the result of electron transfer leading to attraction between oppositely charges ions.

4.1.2 Determine which ions will be formed when metals in groups 1, 2 and 3 of the Periodic Table lose electrons.

4.1.3 Predict the ions which will be formed when members in groups 6 and 7 gain electrons.

4.1.4 State that transition metals can form more than one ion. Restrict to simple examples such as Fe2+ and Fe3+

4.1.5 Predict whether a compound of two elements would be mainly ionic or mainly covalent from the position of the elements in the periodic table, or from their electronegativity values.

4.1.6 Deduce the formula and state the name of an ionic compound formed from a group 1, 2 or 3 metal and a group 5, 6 or 7 non-metal.


4.2 Covalent Bond

4.2.1 Describe the covalent bond as the result of electron sharing.

4.2.2 Draw the electron distribution of single and multiple bonds in molecules. Draw Lewis diagrams to represent a given molecule or polyatomic ion and use them to predict formula. Restrict to second row elements. Describe electron distributions in simple molecules like O2, N2, CO2 , and ethene and ethyne

4.2.3 State and explain the relationship between the number of bonds, bond length and bond strength.

The comparison should include bond length and bond strength of:

4.2.4 Compare the relative electronegativity values of two or more elements based on their position in the periodic table.

4.2.5 Identify the relative polarity of bonds based on electronegativity values.

4.2.6 Draw and deduce Lewis structures of molecules and ions for up to four electron pairs on each atom.

4.2.7 Predict shapes of molecules with up to four electron pairs on the central atom. Invoke VSEPR to explain shapes and bond angles for four pairs of electrons Include carbon compounds.

4.2.8 Identify the shape and bond angle for species with two, three, four, five, or six electron pairs on the central atom.

4.2.9 Predict molecular polarity based on bond polarity and molecular shape.


4.3 Intermolecular Forces

4.3.1 Describe the types of intermolecular force (hydrogen bond, dipole-dipole attraction and van der Waals' forces) and explain how they arise from the structural features of molecules.

4.3.2 Describe how the three types of intermolecular force (hydrogen bond, dipole-dipole attraction and Van der Waals' forces) affect the boiling points of substances. Consider physical properties of molecules such as H2O and H2S, ethers vs. alkanols, halogens

4.4 Metallic bonds

4.4.1 Describe metallic bond formation and explain the physical properties of metals


4.5 Physical properties

4.5.1 Compare and explain the following properties of substances resulting from different types of bonding:melting and boiling points, volatility, conductivity and solubility.


 4.5.2 Predict the relative values of melting and boiling points, volatility, conductivity and solubility based on the different types of bonding in substances.


SSC Topic 5 - States of Matter

5.1 States of Matter

5.1.1 Describe and compare solids, liquids, and gases as the three states of matter.

5.1.2 Describe kinetic theory in terms of the movement of particles whose energy is proportional to absolute temperature. Give description of what happens when temperature is increased.

5.1.3 Describe the Maxwell-Boltzmann energy distribution curve.

5.1.4 Draw and explain qualitatively Maxwell-Boltzmann energy distribution curve for different temperatures.

5.1.5 Describe qualitatively, the effects of temperature, pressure and volume changes on a fixed mass of an ideal gas.

5.1.6 State the ideal gas equation, PV = nRT

5.1.7 Apply the ideal gas equation in calculations


SSC Topic 6 - Energetics

6.1 Exothermic and endothermic reactions

6.1.1 define the terms "exothermic reaction", "endothermic reactions" and "standard enthalpy change of a reaction (DH°)

6.1.2 State the relationship between temperature change, enthalpy change and whether a reaction is exothermic or endothermic.

6.1.3 Deduce, from enthalpy level diagram, the relative stabilities of reactants and products and the sign of the enthalpy change.

6.1.4 Describe and explain the changes which take place at the molecular level in chemical reactions.

6.1.5 Suggest suitable experimental procedures for measuring enthalpy changes of reactions in aqueous solution.


6.2 Calculation of enthalpy changes

6.2.1 Calculate the heat change when the temperature of a pure substance.

6.2.2 Explain that enthalpy changes of reaction measure enthalpy changes for specific quantities of either reactants or products.

6.2.3 Analyze experimental data for enthalpy changes of reactions in aqueous solution. Perform a heat of neutralization lab.

6.2.4 Calculate the enthalpy change for a reaction in aqueous solution using experimental data on temperature changes, quantities of reactants and mass of solution.


6.3 Hess' law

6.3.1 Determine the enthalpy change of a reaction which is the sum of two or more reactions with known enthalpy changes of reaction.


6.4 Bond enthalpies

6.4.1 Define" average bond enthalpy". Recognize as for gaseous states and averages.

6.4.2 Calculate the enthalpy change of a reaction using bond enthalpies.


6.5 Entropy

6.5.1 State and explain the factors which increase disorder in a system.(DS)

6.5.2 Predict whether the entropy change for a given reaction or process would be negative or positive.


6.6 Spontaneity

6.6.1 Define "standard entropy change for a reaction" (DG) using values of absolute entropies.

6.6.2 State whether a reaction or process will be spontaneous by using the sign of DG

6.6.2 State and predict the effect of a change in temperature on the spontaneity of a reaction, given standard entropy and enthalpy changes. Use the equation DG° = D H° - TDS° .


SSC Topic 7 - Kinetics


7.1 Rates of reaction

7.1.1 Define the term "rate of reaction" and describe the measurement of reaction rate

7.1.2 Analyze data obtained from rate experiments.


7.2 Collision theory

7.2.1 Describe and explain collision theory. Minimum energy and collision geometry to be considered.

7.2.2 Define "activation energy Ea" and explain that reactions occur when reacting species have E> Ea

7.2.3 Predict and explain, using collision theory, the qualitative effect of particle size, temperature, concentration and catalysts on the rate of a reaction. Increasing the temperature increases the frequency of collisions, but more importantly, the proportion of molecules with E>Ea increases.

7.2.4 Explain that reactions can occur by more than one step and that one step can determine the rate of reaction.


SSC Topic 8 - Equilibrium


8.1 Dynamic equilibrium

8.1.1 Outline the characteristics of a system in a state of equilibrium.


8.2 The Position of Equilibrium

8.2.1 State the equilibrium constant expression Kc for a homogeneous reaction.

8.2.2 Deduce the extent of a reaction from the magnitude of the equilibrium constant.

8.2.3 Describe and predict the qualitative effects of changes in temperature, pressure and concentration on the position of equilibrium and the value of the equilibrium constant.

8.2.4 State and explain the effect of a catalyst on an equilibrium reaction.

8.2.5 Describe and explain the application of the equilibrium and kinetics concepts to the Haber process and the Contact process.



SSC Topic 9 - Acids and Bases

9.1 Properties of acids and bases

9.1.1 Outline the characteristic properties of acids and bases in aqueous solution.


9.2 Strong and Weak Acids and Bases

9.2.1 Describe and explain the differences between strong and weak acids and bases in terms of the extent of dissociation, reaction with water and conductivity.

9.2.2 State whether a given acid or base is strong or weak.

9.2.3 Describe and explain data from experiments to distinguish between strong and weak acids and bases, and determine the relative acidities and basicititees of substances.


9.3 The pH scale

9.3.1 Distinguish between aqueous solutions that are acidic, neutral or basic using the pH scale.

9.3.2 Identify which of the two or more aqueous solutions is more acidic or basic, using pH values

9.3.3 State that each change of one pH unit represents a tenfold change in acidity or basicity.

9.3.4 Deduce changes in [H+] when pH of a solution changes by more than one pH unit.


9.4 Buffer Solutions

9.4.1 Describe a buffer solution in terms of its compositions and behaviour.

9.4.2 Describe ways of preparing buffer solutions


9.5 Acid-base Titrations

9.51 Draw and explain a graph showing pH against volume of titrant for titrations involving strong acids and bases.


SSC Topic 10 - Oxidation and Reduction


10.1 Oxidation and reduction

10.1.1 Define 'oxidation' and 'reduction' in terms of electron loss and gain.

10.1.2 Calculate the oxidation number of an element in a compound.

10.1.3 State and explain the relationship between oxidations numbers and the names of compounds.

10.1.4 Identify whether and element is oxidized or reduced in simple redox reactions, using oxidation numbers.

10.1.5 Define the terms 'oxidizing agent' and 'reducing agent'.


 10.2 Reactivity

10.2.1 deduce a reactivity series based upon the chemical behaviour of a group of oxidizing and reducing agents.

10.2.2 Deduce the feasibility of a redox reaction from a given reactivity series.

10.2.3 Describe and explain how a redox reaction is used to produce electricity in a voltaic cell.


10.3 Electrolysis

10.3.1 Draw a diagram showing the essential components of an electrolytic cell.

10.3.2 Describe how current is conducted in an electrolytic cell.

10.3.3 Deduce the product for the electrolysis of a molten salt.

10.3.4 Distinguish between the use of a spontaneous redox reaction to produce electricity in a voltaic cell and the use of electricity to carry out a non-spontaneous redox reaction in an electrolytic cell.

10.3.5 describe and explain the use of electrolysis in electroplating.


SSC Topic 11 - Organic Chemistry

11.1 Homologous series

11.1.1 Describe the features of a homologous series.

11.1.2 Predict and explain the trends in boiling points of members of a homologous series.


11.2 Hydrocarbons

11.2.1 Draw structural formulas for the isomers of the non-cyclic alkanes up to C6.

11.2.2 Name straight-chain alkanes up to C6. Use IUPAC rules

11.2.3 Explain the relative inertness of alkanes

11.2.4 Draw structural formulas for straight-chain alkenes up to C5.Geometric isomers (cis/trans) not required

11.2.5 Describe complete and incomplete combustion of hydrocarbons. Formation of CO and C in incomplete combustion.

11.2.6 State that the combustion of hydrocarbons is an exothermic process.


11.3 Other functional groups.

Interrelationships involving significant functional groups. Addition, substitution, oxidation, condensation, esterification and polymerization to be considered. Reproduce flow chart as in IB syllabus details

11.3.1 Draw and state the names of compounds containing up to five carbon atoms with one of the following functional groups: alkanal, alkanoic acid, alkanol, amide, amine, ester, hologenoalkane. Full and condensed forms required.

11.3.2 Explain that functional groups can exist as isomers.

11.3.3 Discuss the existence of optical isomers Be able to identify chiral center.

11.3.4 Discuss the volatility, solubility in water and acid-base behavior of the seven functional groups.

11.3.5 Outline the the reaction of symmetrical alkenes with hydrogen, bromine, hydrogen halides and water.

11.3.6 Outline the use of reactions of alkenes

11.3.7 Outline the polymerization of alkenes.

11.3.8 Outline the condensation reaction of an alcohol with a carboxylic acid to form an ester, and state the uses of esters.

11.3.9 Describe the partial and complete oxidation of ethanol

11.3.10 Deduce the condensation polymers formed by amines and by carboxylic acids.

11.3.11 Describe the formation of peptides and proteins from 2-amino acids.