Module 3 (Periodicity) Flashcards
What is the periodic table?
An arrangement of all the elements in order of increasing atomic number. Elements with similar chemical properties are found in the same group
What are the 4 blocks in the periodic table?
- S block
- D block
- P block
- F block
What is Periodicity?
Periodicity is a repeating pattern across different periods
Explain Atomic Radius
Atomic radii decrease as you move from left to right across a period because the increased number of protons creates more positive charge attraction for electrons which are in the same shell similar shielding.
What is Ionisation energy?
The energy required to remove an outer electron from an element
What is the equation for first ionisation?
H(g) -> H+(g) + e-
What factors affect ionisation energy?
There are three main factors
1. The attraction of the nucleus
(The more protons in the nucleus the greater the attraction)
2. The distance of the electrons from the nucleus
(The bigger the atom the further the outer electrons are from the nucleus and the
weaker the attraction to the nucleus)
3. Shielding of the attraction of the nucleus
(An electron in an outer shell is repelled by electrons in complete inner shells,
weakening the attraction of the nucleus)
Explain why there is a large increase from the 3rd and 4th ionisation energy in Aluminium
There is a large increase in ionisation energy as it is from an inner shell, so is under a stronger force of attraction from the nucleus and has less shielding
Why are successive ionisation energies always larger?
The second ionisation energy of an element is always bigger than the first ionisation energy. This is because the ion formed, is smaller than the atom and the proton-to-electron ratio in the 2+ ion is greater than in the 1+ ion. The attraction between the nucleus and electron is therefore stronger
Explain the trend in the first ionisation energy down group 1
The ionisation energy decreases down the group as the outer electron is further from the nucleus and experiences greater shielding from inner electrons so the attraction from the nucleus is weaker and therefore easier to overcome
Why does the ionisation energy increase across periods 2 and 3?
- The number of protons in the nucleus increases
- The outer electron is in the same shell with similar shielding
- The electrostatic attraction between the outer electron and the nucleus is greater
What are examples of an exception to the trend in ionisation energy?
- Boron and aluminium have lower first ionisation energies than beryllium and magnesium. This is because the outer electrons of boron and aluminium are in a p-orbital higher energy.
- Oxygen and sulfur have lower first ionisation energies than nitrogen and phosphorus, this is because the outer electrons of oxygen and sulfur have [aired electrons in a p-orbital, these repel each other, so one electron is more easily removed
Describe the bonding in metals
Metallic bonding is the electrostatic force of attraction between the positive metal ions and the delocalised electrons
What factors affect the strength of metallic bonds?
The three main factors that affect the strength of metallic bonding are:
1. Number of protons/ Strength of nuclear attraction.
(The more protons the stronger the bond)
2. Number of delocalised electrons per atom (the outer shell electrons are
delocalised)
The more delocalised electrons the stronger the bond
3. Size of ion.
(The smaller the ion, the stronger the bond.)
Describe the bonding and structure of diamond.
Tetrahedral arrangement of carbon atoms. 4 covalent bonds per atom
Describe the bonding and structure of graphite.
Planar arrangement of carbon atoms in layers. 3 covalent bonds per atom in each layer. 4th outer electron per atom is delocalised. Delocalised electrons between layers.
Properties of Giant covalent strucures:
Boiling and melting points: high- because of many strong covalent bonds in macromolecular structure. Take a lot of energy to break the many strong bonds
Solubility in water: insoluble insoluble
conductivity when solid:
- Diamond and sand: poor, because electrons can’t move (localised)
- graphite: good as free delocalised electrons between layers
conductivity when molten: poor
general description: solids
Properties of Metallic lattices:
Boiling and melting points: high- strong electrostatic forces between positive ions and sea of delocalised electrons
Solubility in water: insoluble
conductivity when solid: good, delocalised electrons can move through the structure
conductivity when molten: (good)
general description:
-shiny metal
-Malleable as the positive ions in the lattice are all identical. So the planes of ions can slide easily over one another
-attractive forces in the lattice are the same whichever ions are adjacent
Trends down Group 2:
Melting points decrease down the group. The metallic bonding weakens as the atomic size increases. The distance between the positive ions and delocalized electrons increases. Therefore the electrostatic attractive forces between the positive ions and the delocalized electrons
Atomic radius increases down the Group. As one goes down the group the atoms have more shells of electrons making the atom bigger
Group 2 metals all have the outer shell s2 electron configuration.
General reactions of Group 2 metals:
When the group 2 metals react, they lose their outer shell s2 electrons in redox reactions to form 2+ ions. The energy to remove these electrons is the first and second ionisation energies.
The trend of ionisation energy down the group:
The first and second ionisation energies decrease down the group. The outermost electrons are held more weakly because they are successively further from the nucleus in additional shells. In addition, the outer shell electrons become more shielded from the attraction of the nucleus by the repulsive force of inner shell electrons
The trend of reactivity down group 2:
Reactivity of group 2 metals increases down the group (The reactivity increases down the group. As the atomic radii increase there is more shielding. The nuclear attraction decreases and it is easier to remove outer electrons. Cations form more easily.)
Reactions with oxygen:
The group 2 metals will burn in oxygen.
e.g.
Mg burns with a bright white flame
2Mg + O2 -> 2MgO
MgO is a white solid with a high melting point due to its ionic bonding
Mg will also react slowly with oxygen without a flame.
Mg ribbon will often have a thin layer of magnesium oxide on it formed by reaction with oxygen
2Mg + O2 -> 2MgO
This needs to be cleaned off with emery paper before doing reactions with Mg ribbon
If testing for reaction rates with Mg and acid, an uncleaned Mg ribbon would give a false result because both the Mg and MgO would react but at different rates.
Mg + 2HCl -> MgCl2 + H2
MgO + 2HCl -> MgCl2 + H2O
Reactions with water:
Magnesium reacts in steam to produce magnesium oxide and hydrogen. The Mg would burn with a bright white flame
Mg (s) + H2O (g) -> MgO (s) + H2(g)
Mg will also react with warm water, giving a different magnesium hydroxide product
Mg + 2 H2O -> Mg(OH)2 + H2
This is a much slower reaction than the reaction with steam and there is no flame
The other group 2 metals will react with cold water with increasing vigour down the group to form hydroxides
Ca + 2 H2O (l) -> Ca(OH)2(aq) + H2(g)
Sr + 2 H2O (l) -> Sr(OH)2(aq) + H2(g)
Ba + 2 H2O (l) -> Ba(OH)2(aq) + H2(g)
The hydroxides produced make the water alkaline
One would observe:
fizzing, (more vigorous down group)
the metal dissolving, (faster down group)
the solution heating (more down group)
and with calcium, a white precipitate appears (less precipitate forms down group)
Reactions with acids:
The group 2 metals will react with acids with increasing vigour down the group to form a salt and hydrogen
Ca + 2HCl (aq) -> CaCl2(aq) + H2(g)
Sr + 2 HNO3(aq) -> Sr(NO3)2(aq) + H2(g)
Mg + H2SO4(aq) -> MgSO4(aq) + H2(g)
If barium metal is reacted with sulfuric acid it will only react slowly as the insoluble barium sulfate produced will cover the surface of the metal and act as a barrier to further attack. Ba + H2SO4 -> BaSO4 + H2 The same effect will happen to a lesser extent with metals going up the group as the solubility increases. The same effect does not happen with other acids like hydrochloric or nitric as they form soluble group 2 salts.
The action of water on oxides of elements in Group 2
The group 2 oxides react with water to form hydroxides of varying solubility
CaO (s) + H2O (l) -> Ca(OH)2(aq) pH 12
Group 2 oxides are basic as the oxide ions accept H+ ions to become hydroxide ions in these reactions
MgO (s) + H2O (l) -> Mg(OH)2(s) pH 9Mg(OH)2
is only slightly soluble in water so fewer free OHions are produced and so lower pH
Calcium hydroxide is reasonably soluble in water. It is used in agriculture to neutralise acidic soils. If too much calcium hydroxide is added to the soil, excess will result in soils becoming too alkaline to sustain crop
Magnesium hydroxide is classed as partially soluble in water.
A suspension of magnesium hydroxide in water will appear slightly alkaline (pH 9) so some hydroxide ions must therefore have been produced by a very slight dissolving. Magnesium hydroxide is used in medicine (in suspension as milk of magnesia) to neutralise excess acid in the stomach and to treat constipation.
Mg(OH)2 + 2HCl -> MgCl2 + 2H2O
It is safe to use as it so weakly alkaline.
An aqueous solution of calcium hydroxide is called lime water and can be used as a test for carbon dioxide. The limewater turns cloudy as white calcium carbonate is produced.
Ca(OH)2 (aq) + CO2 (g) -> CaCO3 (s) + H2O(l)
What are Halogens?
Group 7 elements, all of which exist as diatomic molecules:
Fluorine (F2): very pale yellow gas. It is highly reactive
Chlorine (Cl2): greenish, reactive gas, poisonous in high concentrations
Bromine (Br2): red liquid, that gives off dense brown/orange poisonous fumes
Iodine (I2): shiny grey solid sublimes to purple gas.