Periodicity
- Created by: Lotto65
- Created on: 03-05-18 17:20
View mindmap
- Periodicity
- Periodic Table
- Elements arranged in order of atomic number
- Groups = Same number of valence electrons
- Bromine and mercury only liquids
- Nobles gases (inert gases), nitrogen, oxygen, chlorine, hydrogen and fluorine only gases
- Carbon is a solid at room temperature
- Periods = Same number of energy levels
- Which is main energy level
- Lanthanoids and actinoids omitted - actinoids start with actinium
- Metals formed from elements with lower ionisation energies (on the left and bottom of groups) to readily form positive ions and a sea of delocalised electrons
- Metals = Low ionisation energies, less exothermic elecron affinities, large atomic radii and low electronegativity
- Periodic table arranged according to the highest energy subshell occupied by electrons
- Physical Properties
- Atomic radius
- Increases down group
- Number of electron shells increases
- Nuclear charge increases down a group but so does electron repulsion to counter
- Decreases across period
- Nuclear charge increases to attract outer electrons closer to nucleus
- Increase in nuclear charge but no real change in shielding because same number of shells
- Nuclear charge increases to attract outer electrons closer to nucleus
- Half the internuclear distance
- Increases down group
- Ionic Radii
- Positive ions are smaller than element
- More electrons repelling same nuclear charge in elemental form
- Negative ions are bigger than element
- More electrons for same nuclear charge expands electron cloud as more repulsion
- Metals and non-metals decrease across period due to increase in nuclear charge
- Positive ions are smaller than element
- Ionisation Energy
- Decreases down group
- Increase in electron shells so outer electrons further from nucleus and less attraction. More shielding
- Increase in nuclear charge balanced by shielding
- Increase in electron shells so outer electrons further from nucleus and less attraction. More shielding
- Increases across period
- Nuclear charge increases so electrons more strongly attracted to the nucleus
- Shielding the same. Electrons in same electron shell do not shield each other very well
- Decreases down group
- Electron Affinity
- First electron affinity is exothermic
- Data incomplete as hard to measure experimentally
- Affinity decreases down the group (mainly from chlorine down)
- Less exothermic down the group
- Increase in number of electron shells and atomic radii means attraction is weaker to added electrons
- Across a period
- More exothermic across period
- Increases across a period
- Increase in nuclear charge and smaller atomic radius means stronger attraction to added electron
- Electronegativity
- Decreases down a group
- Increase atomic radius due to more electron shells and more shielding so less attraction
- Increases across a period
- Nuclear charge increases across a period so atomic radius is smaller so more attraction
- No significant change in shielding due to same number of inner shells
- Nuclear charge increases across a period so atomic radius is smaller so more attraction
- The attraction of an atom in a molecule for the electron pair in the covalent bond for which it is part of
- Decreases down a group
- Properties of elements in group 1 and group 17
- Group 1
- Soft, reactive, low melting points
- Reactivity increases down the group
- Atomic radius and number of electron shells increases and more shielding so less attraction for outer electron
- Ionisation energy decreases down group
- React vigorously in oxygen and tarnish rapidly in air
- Reactions with water
- Sodium fizzes, turns into a ball (melts) and moves on the surface
- Potassium bursts into a lilac flame
- Caesium explodes on contact
- Strong bases and ionise completely
- Melting point decreases down the group
- Increases in atomic radius and larger atoms means less strong attraction between positive nuclei and delocalised electrons so less energy to break apart the lattice
- Liquid sodium used as a coolant in some nuclear reactors
- Same group = SIMILAR chemical properties because all have the same number of valence electrons
- Group 17
- Reactivity decreases down group
- Atomic radius increases so larger atom so less strong attraction to outer electron to be added
- Melting point increases down the group
- London forces increase as atomic mass increases so more energy required to break intermolecular forces between molecules
- React with alkali metals to form white or colourless salts
- Alkali metal chlorides, bromides and iodides are soluble and produce colourless, neutral solutions
- Displacement reactions
- Bromine displaced gives orange solution
- Iodine displaced produces dark red-brown solution
- Chlorine is pale yellow-green solution
- Redox reaction and more reactive halogen oxidises less reactive halide ion
- Stronger affinity for electrons so removes them
- Reactivity decreases down group
- Group 1
- Oxides
- Metallic oxides = Basic
- Non-metallic oxides = acidic
- Alkali metal + water --> metal hydroxide
- Metal oxide + acid --> salt + water
- Aluminium oxide is amphoteric
- Reacts with acids and bases to form salts
- Forms Al3+ with acid and aluminium hydroxide with bases
- Does not react with water
- Insoluble
- Reacts with acids and bases to form salts
- Magnesium oxide sparingly soluble and only slightly alkaline
- NO is produced in an internal combustion engine, is insoluble and forms neutral solutions
- Contributes to photochemical smog
- NO2 forms nitric (V) acid with water
- N2O also a neutral oxide
- Atomic radius
- D-block elements
- Transition element: An element that forms at least one stable ion with a partially filled d subshell
- Zinc has a full d subshell, does not exhibit some of the properties of transition elements, has one stable ion and does not form coloured compounds
- Scandium is regarded as a transition element but has no d subshell and nearly always forms a 3+ ion and occasionally a 2+ ion but the bonding is complicated and not always a 2+ ion is involved
- Atomic radius and first ionisation energy minimal change across the period
- Properties
- Formed coloured compounds/ complexes
- Compounds can be magnetic
- Metals and their complexes can act as catalysts
- Finely divided iron acts as a catalyst in the Haber process
- Normally homogeneous catalysts (same phase as reactants) but iron is a heterogeneous catalyst in the Haber process (different phase to reactants)
- Depends on transition metal having varying oxidation numbers and coordinating with other molecules/ ions
- Can form complex ions
- Can exhibit multiple oxidation numbers in compounds/ complexes
- 4s and 3d subshells close in energy so no large jump in ionisation energy when they are removed
- Number of electrons lost depends on lattice enthalpy, hydration enthalpy and ionisation energy
- 4s and 3d subshells close in energy so no large jump in ionisation energy when they are removed
- High melting points and densities - typical metals
- 4s electrons removed before 3d electrons to form ions
- Maximum oxidation number increases as number of electrons in 4s and 3d subshells increases
- Always have oxidation state +2
- 4s electrons
- Magnetism
- Paramagnetism = Unpaired electrons and are attracted by a magnetic field. More dominant
- Diamagnetism = Paired electrons repel a magnetic field slightly
- More unpaired or paired electrons, the greater the magnetic moment
- Compounds of copper (I) are diamagnetic only
- Complex Ions
- Coordinate covalent bond where ligand is Lewis base and ion is Lewis acid
- Ligand: Negative ions or neutral molecules that have a lone pair of electrons
- Hexaaquairon(II) ion
- All transition metals apart from titanium can form octahedral complexes with the water ligand with a 2+ charge
- Can undergo substitution reactions to replace ligands (add concentrated HCl to add Cl- ions)
- Reversible reactions
- In the presence of concentrated HCl, copper (II) sulfate solution goes from blue to green
- CO is a neutral ligand. Br- is a negatively charged ligand (all halogens)
- Coloured Complexes
- In the gaseous state, all d orbitals are degenerate in transition metal
- When surrounded by ligands, 2 orbitals in transition metal raise to higher energy and 3 to lower energy
- Repulsion between ligand lone pairs and electrons in d orbitals
- 2 d orbitals face ligands so are raised. 3 are between ligands so are lowered relative to the other 2 orbitals
- Absorption of light energy promotes electron from lower to higher orbitals so light emitted shows everything but the colour of the light absorbed (shows complementary colour)
- MUST BE IN VISIBLE SPECTRUM
- REQUIRES A PARTIALLY FILLED D SUBSHELL
- Need a partially filled d subshell to allow an electron to be promoted (to have electrons to promote and for space for them to be promoted to)
- Otherwise compounds will be colourless
- Factors affecting colour complexes
- Greater difference in energy between lower and higher orbitals means higher frequency of light absorbed
- Transition metal ion
- Manganese 2+ is pale pink/ colourless but Fe2+ is pale green
- Different electron configurations and different levels of repulsion and therefore splitting
- Higher nuclear charge means ligands pulled in more closely, more repulsion and greater splitting
- Oxidation number
- Greater oxidation number, greater splitting
- Different electron configurations
- Greater charge on ion means ligands closer, more repulsion, more splitting
- Nature of ligand
- Spectrochemicals series of how much ligands split orbitals
- Stronger field ligands = Greater splitting ligands
- Neutral molecules generally cause more splitting than ions apart from CN- which is the strongest with CO
- Charge density = Charge per unit volume
- F- and I- have same charge but F- is smaller so more repulsion and greater splitting of d orbitals
- Pi bonding between lone pair and d orbitals reduces splitting so more pi bonding, less splitting
- Pi electron donation effect
- CO has electron density in cloud from metal ion donated to it, meaning there is a greater chance of splitting
- Periodic Table
Comments
No comments have yet been made