The Periodic Table consists of blocks of elements s block d block p block
There is a clear link between the Periodic Table and the electronic configuration of an element 1s 2s 2p 3s 3p 4s 3d 4p 1s
ATOMIC RADIUS Atomic radius decreases from Na to Ar Nuclear charge (number of protons) increases Electrons are added to the same main energy level Shielding is constant The electrons experience a greater electrostatic force of attraction to the nucleus
FIRST IONISATION ENERGY First ionisation energy increases from Na to Ar Nuclear charge (number of protons) increases Atomic radius decreases Shielding is constant The outermost electron experiences a greater electrostatic force of attraction to the nucleus
ELECTRONEGATIVITY Electronegativity increases from Na to Cl Nuclear charge (number of protons) increases Atomic radius decreases Shielding is constant It becomes easier to attract the electrons of a covalent bond.
MELTING POINT / K Period 2 5000 4000 3000 2000 1000 0 Giant lattice Simple molecular Atomic
MELTING POINT / K Period 3 2000 1500 1000 500 0 Giant lattice Simple molecular Atomic
MELTING POINT Metals have giant lattice structures Strong electrostatic forces of attraction between positive metal ions and their delocalised electrons Melting point increases from Na to Al Size of ionic charge increases Number of delocalised electrons increases Size of ion (ionic radius) decreases
MELTING POINT The Group IV element has the highest melting point in each period Giant lattice structure Strong covalent bonds between atoms
MELTING POINT Group V, VI, and VII elements have very low melting points Simple molecular structure Covalent bonding Weak van der Waals forces of attraction between molecules
OXIDATION Loss of electrons (OILRIG) Increase in oxidation state REDUCTION Gain of electrons (OILRIG) Decrease in oxidation state
OXIDISING AGENT An electron acceptor The oxidising agent is always reduced. REDUCING AGENT An electron donor The reducing agent is always oxidised.
Oxidation states in a compound add up to zero Oxidation states in an ion add up to the charge The most electronegative element has a negative oxidation state and all the other elements have positive oxidation states Element Oxidation state Uncombined element 0 Hydrogen 1 or +1 Oxygen 2 Group I +1 Group II +2 Group VII 1 to +7
Used to represent oxidation or reduction alone Note the reagent and product species given. Balance the element oxidised or reduced. Balance charge by adding electrons to one side. (OILRIG)
Make the number of electrons in each half equation the same. Add the two half equations together.
Atomic radius increases from Be to Ba Atomic number increases so there are more electrons More main energy levels are needed to accommodate these Each additional main energy level is further from the nucleus
For the elements of Group II, ionic radius is always smaller than the atomic radius. On ionising, the Group II atom loses both the electrons in its outermost energy level It loses its outermost energy level
First ionisation energy decreases from Be to Ba The outermost electron is easier to remove Despite more protons in the nucleus Because atomic radius increases Because there are more occupied energy levels shielding it from the nucleus It experiences a weaker electrostatic force of attraction
Electronegativity decreases from Be to Ba Nuclear charge (number of protons) increases Atomic radius increases Shielding increases Thus it becomes more difficult to attract the electrons of the bond
Melting point decreases from Be to Ba Size of ionic charge stays the same Number of delocalised electrons stays the same Size of ion (ionic radius) increases When the electrons are further from the centres of positive charge in the ions, the electrostatic forces of attraction are weaker
Atoms react by losing two electrons to form an M 2+ ion The oxidation state of the elements increases from 0 to +2 The elements are oxidised Group II elements are reducing agents (electron donors)
Reactivity of the elements increases from beryllium to barium. First and second ionisation energies decrease from beryllium to barium There is an increase in atomic radius There is an increase in the number of occupied energy levels shielding the nucleus It becomes easier to lose two electrons They experience a weaker electrostatic force of attraction
Reactions of elements (Redox reactions) With oxygen 2Ca(s) + O 2 (g) 2CaO(s) With water Ca(s) + 2H 2 O(l) With acid Ca(s) + 2HCl(aq) Ca(OH) 2 (aq) + H 2 (g) alkaline CaCl 2 (aq) + H 2 (g)
Reactions of oxides With water CaO(s) + H 2 O(l) Ca(OH) 2 (aq) alkaline With acid CaO(s) + 2HCl(aq) CaCl 2 (aq) + H 2 O(l)
Reactions of hydroxides With water Ca(OH) 2 (s) + (aq) Ca(OH) 2 (aq) alkaline With acid Ca(OH) 2 (s) + 2HCl(aq) CaCl 2 (aq) + 2H 2 O(l)
Reactions of carbonates With heat CaCO 3 (s) CaO(s) + CO 2 (g) With acid CaCO 3 (s) + 2HCl(aq) CaCl 2 (aq) + H 2 O(l) + CO 2 (g) Testing for carbon dioxide Ca(OH) 2 (aq) + CO 2 (g) CaCO 3 (s) + H 2 O(l)
Atomic radius increases from fluorine to iodine. Atomic number increases so there are more electrons More main energy levels are needed to accommodate these Each additional main energy level is further from the nucleus
For the elements of Group VII, ionic radius is always larger than the atomic radius. A halide ion always has one more electron than the atom from which it was formed. This electron enters the outermost energy level of the atom and there is an increase in the mutual electrostatic forces of repulsion between the negatively charged electrons.
First Ionisation Energy decreases from fluorine to iodine. The outermost electron is easier to remove Despite more protons in the nucleus Because atomic radius increases Because there are more occupied energy levels shielding it from the nucleus It experiences a weaker electrostatic force of attraction
Electronegativity decreases from fluorine to iodine. Nuclear charge (number of protons) increases Atomic radius increases Shielding increases Thus it becomes more difficult to attract the electrons of the bond
Boiling point increases from fluorine to iodine. Size of molecules increases Number of electrons increases More temporary induced dipoles occur Strength of van der Waals forces increases More heat energy is needed to overcome these forces.
Volatility decreases from fluorine to iodine. Size of molecules increases Number of electrons increases More temporary induced dipoles occur Strength of van der Waals forces increases More heat energy is needed to overcome these forces.
Oxidising agents are electron acceptors Fluorine, chlorine, bromine and iodine are all powerful oxidising agents The ability to behave as oxidising agents decreases from fluorine to iodine There is an increase in atomic radius There is an increase in the number of occupied energy levels of electrons shielding the nucleus It becomes more difficult to attract an electron.
Reducing agents are electron donors Fluoride, chloride, bromide and iodide ions are all powerful reducing agents The ability to behave as reducing agents increases from fluoride to iodide ions There is an increase in ionic radius There is an increase in the number of occupied energy levels of electrons shielding the nucleus It becomes easier to donate an electron.
INE Names of the Halogen elements end in INE Names of the Halide ions end in IDE IDE Cl - (aq) Br - (aq) I - (aq) Cl 2 (aq) x Br 2 (aq) x x I 2 (aq) x x x
Redox reactions (displacement) Cl 2 (aq) + 2Br - (aq) Cl 2 (aq) + 2I - (aq) Br 2 (aq) + 2I - (aq) 2Cl - (aq) + Br 2 (aq) 2Cl - (aq) + I 2 (aq) 2Br - (aq) + I 2 (aq) Orders of reactivity: Chlorine > Bromine > Iodine Iodide > Bromide > Chloride
Redox reactions (disproportionation) Cl 2 (aq) + H 2 O(l) HCl(aq) + HClO(aq) Cl 2(aq) + 2NaOH (aq) NaCl (aq) + NaClO (aq) + H 2 O (l) Cl 2(aq) + 2OH - (aq) Cl - (aq) + ClO - (aq) + H 2 O (l) In a disproportionation reaction, one species is both oxidised and reduced
Testing for halide ions in aqueous solution Ag + (aq) + Cl - (aq) AgCl(s) White precipitate soluble in dil. NH 3 (aq) Ag + (aq) + Br - (aq) AgBr(s) Cream precipitate soluble in conc. NH 3 (aq) Ag + (aq) + I - (aq) AgI(s) Pale yellow precipitate insoluble in NH 3 (aq)