The periodic table Flashcards
Periodicity
a regular periodic variation of properties of elements with atomic number and position in the periodic table
First ionisation energy-
energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous 1+ ions
Factors affecting ionisation energy
-atomic radius- larger atomic radius, the lower nuclear attraction, therefore lower ionisation energy
-nuclear charge- higher nuclear charge, larger nuclear attraction
-electron shielding- the more shells, the larger the shielding, so less nuclear attraction
1st, 2nd and 3rd ionisation energy equations for Li
Li-> e- + Li+
Li+ -> e- + Li2+
Li2+-> e- + Li3+
Successive ionisation energies
-electrons removed, less repulsion between remaining electrons, so drawn closer nucleus
-positive nuclear charger outweighs the negative charge every time an electron is removed
-as the distance of each electron to the nucleus decreases, nuclear attraction increases, causing ionisation energy to increase
Draw ionisation graph for nitrogen (7)
graph
Ionisation energy trend across period
-general increase because:
-higher nuclear charge as more protons
-electrons added to same shell, which is drawn closer to nucleus
-same number of inner shells, so no change in shielding
8 elements that don’t follow ionisation energy trend
Be-> B and Mg-> Al
-both decrease in ionisation energy even though they’re further along period
-this is because B and Al have their outer electron in a P orbital whereas Be and Mg outer electron is in an S orbital
-P orbitals have higher energy and therefore further from the nucleus
-this therefore decreases nuclear attraction, meaning lower ionisation energy
N->O P->S decrease ionisation energies even though they’re along the period
-because in N and P, each p orbital only contains one electron
-in O and S, the p orbital has 2 electron in = spin paired
-spin paired electrons experience some repulsion and therefore are easier to remove
Ionisation energy trends down a group
-first ionisation energies decrease
-more shells, so more shielding
-larger atomic radius
-nuclear charge increases, but its effects are outweighed by atomic radius and shielding
Metallic bonding
electrostatic forces of attraction between positive ions and delocalised electrons
Giant metallic lattice structure
-delocalised electrons spread throughout
-electrons free to move
-over whole structure, charges balance
Properties of giant metallic lattices
-high MP and BP- attraction between ions and electrons are very strong so high temp required to overcome electrostatic forces
-electrical conductors- delocalised electrons free to move and can carry charge
-ductile- can be drawn out/ stretched eg. wires
-malleable- can be hammered into different shapes
Types bonding across elements in period 2 + 3
Giant metallic:
-Li, Be
-Na, Mg, Al
Giant covalent
-B, C
- Si
Simple molecular
-N2, O2, F2, Ne
-P4, S8, CL2, Ar
Trend melting points in metals
-across the period, MP increases
-this is because ionic charge increases, ionic size decreases
-therefore attraction between ions and electrons increase, so more energy required to disrupt the lattice
Group 2 properties
-high MP + BP
-light metals, low densities
-white compounds
Reaction group 2 element and O2 eg. Ca
2M + O2 -> 2MO
2Ca + O2 -> 2CaO
Reaction group 2 elements and H2O eg. Ca
M + 2H2O -> M(OH)2 + H2
Ca + 2H2O -> Ca(OH)2 + H2
Reaction group 2 elements and dilute acid eg. Ca + HCl
M + 2HCl-> MCl2 + H2
Ca + 2HCl -> CaCl2 + H2
Group 2 oxide + H2O
MO + H2O -> M(OH)2
CaO + H2O -> Ca(OH)2
Solubility of group 2 hydroxides
-solubility of hydroxides increase down the group
-this is because they release more OH-
-therefore alkalinity increases down the group
-Be(OH)2 = insoluble
-more soluble= less likely to form precipitate
Uses group 2 compound
Ca(OH)2
- neutralise soils by reducing acidity
Mg(OH)2
-used in indigestion tablets by neutralising excess stomach acid
CaCO3
-used in limestone and marble
-drawback as it readily reacts with acid, rainwater is slightly acidic so therefore erodes over time
Trend BP of halogens
-Bp increases down the group
-down group physical state changes from gas to liquid to solid (F2 + Cl2= gas, Br2= liquid, I2 + At2= solid)
-this is because each elment has an extra shell of electrons leading to stronger london forces between molecules
Halogen REDOX reactions
-Cl2 can oxidise Br- and I- ions
-Br can only oxides I-
Colour of halogen in water
Cl2= pale green
Br2= orange
I2= brown
Colour of halogen in cyclohexane
Cl2= pale green
Br2= orange
I2= violet
Disproportionation-
reaction where the same element is both oxidised and reduced
eg. Cl2 + H2O -> HClO + HCl
Formation bleach
Cl2 + 2NaOH -> NaCl + NaClO + H2O
CO3 2- test
-add dilute strong acid eg. H2SO4
-collect any gas and bubble through limewater
-should fizz as CO2 produced
-CO2 turns limewater cloudy
SO4 2- test
-add dilute HCl and BaCl2
-white ppt
Halide ions test
-dissolve suspected halide in water
-add silver nitrate
-add ammonia if can’t distinguish colour
-Cl- = white ppts, soluble in dilute NH3
-Br- = cream ppt, soluble in conc NH3
-I- = yellow ppt, insoluble in NH3
NH4+ test
-add NaOH and heat gently
-test any gas with red litmus paper
-ammonia turns red litmus paper blue
