IB Chemistry - Topics 3&13

• The periodic table arranges elements in order of increasing atomic number (Z) and the group number tells you the number of electrons in the outermost energy level (valence electrons).
• The period number on the periodic table tells you the total number of occupied energy levels.
• As the green colour is removed from the visible light, its complementary colour red is seen.
• The frequency and wavelength of light absorbed to promote the electron in the compound [Ti(H 2 O 6 )] 3+ is in the green region of the visible spectrum.
• The remaining colours of visible light not absorbed are: red, orange, yellow, green, blue, indigo, violet.
• Atomic radius is the distance from the nucleus to the outer electron(s) and is usually measured as half the distance between two bonded nuclei of atoms of the same element.
• Down a group, atomic radius increases because each time we descend a group a new energy level is added.
• Even though nuclear charge increases as we descend a group, so does the number of shielding electrons which cancels this out.
• Negative electrons are held in the atom by the electrostatic attraction between them and the positive nucleus.
• The greater the nuclear charge the more strongly the electrons are “held in”.
• Inner shells of electrons “shield” the outer electrons from some of the attractive force of the nuclear charge.
• The greater the shielding (the more inner shells there are) the less strongly the electrons are attracted to the nucleus.
• The first ionisation energy of an element is “the energy needed to remove one electron from each atom in one mole of gaseous atoms”.
• The amount of energy released when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion is known as the first electron affinity.
• Down a group, the first ionisation energy decreases.
• Across a period, the first ionisation energy increases.
• The metal ion and its surrounding ligands are called a complex ion.
• Halogens form diatomic molecules such as H2, N2, O2, F2, I2, Cl2, Br2.
• Sodium reacts with chlorine to form Sodium Chloride.
• In the Pauli Exclusion Principle, each atomic orbital can hold a maximum of 2 electrons.
• Reactivity decreases as you go down Group 7 because Halogen atoms want to gain one extra electron into their valence shell, which is more difficult as you go down the group due to increasing atomic radius, distance of valence electrons from nucleus, and number of shielding electrons.
• Ionisation energy decreases as you go down Group 1 because less energy is required to remove an electron as you go down the group due to increasing atomic radius, distance of valence electrons from nucleus, and number of shielding electrons.
• Potassium produces lilac flame when reacting with water.
• Potassium reacts with iodine to form Potassium Iodide.
• Scandium and copper are transition metals as they have a partially filled d orbitals.
• In the Aufbau Principle, electrons are placed into the lowest energy orbitals first.
• The attraction between positive metal ions and the sea of delocalised electrons is weaker as you go down Group 1.
• Sodium and Potassium react with water to form a molten sphere and gas respectively.
• Fluorine has the smallest atomic radius and least shielding electrons so can attract electrons more strongly than the other halogens.
• A ligand is a molecule or ion that contains a lone pair of electrons and so can form a coordinate bond to a central metal ion.
• Melting point and boiling point increase as you go down Group 7 because the halogens are all simple molecules with weak intermolecular forces, and as the molecular weight (Mr) increases, the number of electrons increases, and so does the magnitude of the London dispersion forces attracting the molecules to each other.
• Sodium forms a molten sphere when reacting with water.
• The smallest, lowest, and highest ionisation energies are found in Lithium, Sodium, Potassium, Rubidium, Caesium respectively.
• A transition metal is an element that has an atom or ion with a partially filled d sublevel.
• An oxidising agent is a substance which takes electrons from other substances and is itself reduced.
• Electronegativity decreases as you go down Group 1 because it becomes harder to attract a bonding pair of electrons as you go down the group due to increasing atomic radius, distance of valence electrons from nucleus, and number of shielding electrons.
• Electronegativity is the ability of an atom to attract a bonding pair of electrons.
• Down a group, electronegativity decreases.
• Across a period, electronegativity increases.
• Melting point and metallic character are related to the strength of metallic bonds.