Antoine-Laurent de Lavoisier’s periodic table
- distinguished between metals and non metals
- included some compounds and mixtures
- included terms such as light which he believed to be substances
Berzelius’ periodic table
- published table of atomic weights
- determined the composition by mass of many compounds
- introduced letter based symbols for elements
Dobereiner’s periodic table
- noticed certain groups of 3 elements (triads) ordered by atomic weight would have a middle element with a middle weight and propoerties that were roughly an average of the other two.
Newlands’ periodic table
- elements arranged in order of relative atomic weights
- suggested elements show similar properties to the element 8 places after it (law of octaves)
Mendeleev’s periodic table
- elements ordered by atomic weight
- elements arranged in vertical columns
- gaps left where no element fitted repeating pattern, later the missing elements have been found to match Mendeleev’s predictions
- arrangement of elements in order of atomic weights corresponded to their distinctive chemical properties
Moseley’s input to the periodic table
- determined the atomic number for all known elements
- elements vary periodically with atomic numbers, rather than atomic weight
- this corrected the order of some elements
Seaborg’s input to the periodic table
- discovered the transuranic elements
- placed the actinide series below the lanthanide series at the bottom of the table
How to shorten an electronic configuration (examples for Li, Na and K)
The inner shell configuration is based on the noble gas that comes before the element in the periodic table.v
Li: 1s2 2s1 or [He] 2s1
Na: 1s2 2s2 2p6 3s1 or [Ne] 3s1
K: 1s2 2s2 2p6 3s1 3p6 4s1 or [Ar] 4s1
First ionisation energy
The energy required to remove one electron from each atom in one mole of the gaseous element to form one mole of gaseous 1+ ions.
Equation for the first ionisation energy of sodium
Na(g) –> Na+(g) + e-
Factors affecting ionisation energy
- Atomic radius - the larger the atomic radius the smaller the nuclear attraction on outer electrons
- Nuclear charge - the higher the nuclear charge, the larger the attractive force on outer electrons
- Electron shielding - inner shells of electrons repel the outer shells, the more inner shells the larger the shielding and the smaller the nuclear attraction on the outer electrons
Equation for second ionisation energy of lithium
Li+(g) –> Li2+(g) + e-
Why is each successive ionisation energy higher than the one before?
- As each electron is removed there is less repulsion between remaining electrons so each shell will be drawn closer to the nucleus
- Positive nuclear charge outweighs the negative charge each time an electron is removed
- As distance of outer electrons decreases, nuclear attraction increases so more energy is required to remove each electron
What is the trend in ionisation across a period and why?
The attraction between the nucleus and outer electrons increases across a period so more energy is needed to remove the outer electron.
- number of protons in nucleus increases so there is a higher attraction on electrons as increased nuclear charge pulls the electrons closer to the nucleus. This factor has the biggest effect on ionisation energy
- electrons added to same shell, outer shell is drawn inwards slightly
- same number of shells so shielding doesn’t change
Why is there a decrease in ionisation energy between groups 2 and 13?
Group 13 elements have their outermost electron in a p-orbital whereas group 2 elements have outer electron in s-orbital. p-orbitals have a slightly higher energy than s-orbitals so a slightly further from the nucleus so electrons in p-orbitals are easier to remove.
Why is there a decrease in ionisation energy between group 15 and 16?
In groups 13, 14 and 15, each of the p-orbitals contain a single electron. In group 16, the outer electron is now spin-paired in the first p-orbital. Electrons that are spin paired experience some repulsion which makes the outer electron easier to remove.
Why is there a sharp decrease in ionisation energy between the end of one period and the beginning of the next?
Another shell has been added which is further from the nucleus which leads to an increase in distance of the outer electron from the nucleus and an increase in electron shielding by inner shells.
What is the trend in ionisation energy moving down a group and why?
Moving down a group, first ionisation energy decreases.
- number of shells increases so the distance of outer electrons to nucleus increases so there is a weaker force of attraction on outer electrons
- more inner shells so shielding effect increases, this leads to a weaker attraction on outer electron.
Even though the number of protons increases, the increased attraction is outweighed by the increased nuclear radius and shielding.
Metallic bonding
- Positive ions occupy fixed positions in lattice
- outer shell electrons are delocalised and are shared between all atoms in metallic structure
- metal held together by strong attractions between positive ions and negative electrons
In a giant metallic lattice…
- delocalised electrons spread throughout the structure and can move within the structure
- over the whole structure the charges must balance
A giant metallic lattice is described as a lattice of positive ions in fixed positions surrounded by a sea of delocalised electrons.
Why do metals have high melting and boiling points?
- electrons free to move throughout structure but positive ions remain where they are
- attraction between positive ions and negative delocalised electrons is very strong
- High temperature needed to overcome metallic bonds and dislodge ions from rigid positions
Why are metals good electrical conductors?
- delocalised electrons can move freely anywhere in the metallic lattice
What does it mean to be malleable and ductile and why do metals have these properties?
Ductile - can be drawn out or stretched, allows metals to be drawn into wires
Malleable - can be hammered into different shapes
Delocalised electrons allow the metallic structure to have a degree of ‘give’ which allows atoms and layers to slide past each other. This allows metals to have these properties.
What is the trend in melting points across periods 2 and 3 for groups 1 - 14?
- Between groups 1 and 14, melting points steadily increase because the elements have giant structures. - - The elements with giant metallic lattices (Li and Be for period 2 and Na, Mg and Al for period 3) the nuclear charge and number of outer shell electrons increases. This causes a stronger attraction.
- Elements with a giant covalent lattice (B and C for period 2 and Si for period 3). Each succesive group has more electrons with which to form more covalent bonds
