Periodicity Flashcards
How is atomic number arranged in the periodic table
From left to right atomic number increases.
How are electrons arranged in groups for the periodic table
Elctrons are arranged into vertical columns called groups.
Each element in the a group has the same properties due as it has the same amount of electrons in its outer shell.
Periodic trend across periods 2 and 3
For each period the S and P orbitals are filled in the same way. The S orbital first requiring two electrons then the P orbital requiring 6 electrons.
How are elemennts classified into blocks
There are 4 blocks (s, p, d, f), corresponding to the highest energy sub-shell of each element.
Describe ionisation energy
The measure of how easily an atom losses electrons to form positive ions
Define the first ionisation energy
The amount of energy required to remove one mole of electrons from one mole of gasseous atoms
what three factors affect ionisation energy (nuclear attraction between the nucleus and electrons)
Nuclear charge
Atomic radius
Electron sheilding
Define Nuclear charge
The greater the number of protons in the nucleus the greater the nuclear attration between the nucleus and electrons
Define atomic radius
The larger the distance between the nucleus and electrons the less the nuclear attraction
Define elctron sheilding
Electrons in the inner most shells repel outer shell electrons, reducing the outer shell electrons nuclear attractive force between them and the nucleus.
Define the second ionisation energy
The energy required to remove one mole of electrons from one mole of gaseous 1+ ions to form one mole of gaseous 2+ ions.
What evidence is there for different energy levels in an atom?
Succesive ionistaion energies show a large difference in (n and n+1) ionisation energies suggests that an electron must of been removed from a different shell.
What does succesize ionisation energy show for atoms
Their group and their identity.
What overall trends are there in first ionisation energies.
There is a general increase in first ionisatin energies across a period.
There is a general decrease in first ionisation energies down a group.
There is a sharp decrease in first ionisation enrgies as a new period starts.
Why does first ionisation energy increase across a period.
Electron sheilding remains the same, however nuclear charge increases causing atomic radius to also decrease. Increasing the nuclear attraction between the nucleus and outer shell electrons, increasing the first ionisation energy.
Why does first ionistaion enery decrease down a group?
Atomic radius increases aswell as sheilding, outweighing the overall increase in nuclear charge. Causing a decrease in nuclear attraction between the nucleus and outer shell electrons, decreasing the first ionisation energy.
what subshell trends occur in first ionisation energy?
(Only for period 1 and 2)
Ionisation energy decreases between:
- Beryllium to Boron
- Nitrogen to Oxygen
Why does frist ionisation energy decrease from Beryllium to Boron?
Berylliums highest subshell is 2s whilst boron’s is 2p. Boron’s 2p subshell has a further electron radius and more sheiling than Berylliums 2s subshell and therefore has less nuclear attraction requiring a lower first ionistaion energy.
why does first ionisation energy decrease from Nitrogen to Oxygen?
Nitrogen only has one elctron in all 3 of its p orbitals, whilst oxygen has one of its 3 p orbitals containing two electrons. These electrons repel eachother requiring less energy to remove an electron from oxygen, giving oxygen a lower first ionistion energy.
Define metallic bonding
Metallic bonding is the strong electrostatic attraction between cations (positvie ions) and delocalised electrons.
What structure do all metals have
Giant metallic lattice
Describe the giant metallic lattice
All cations are in a fixed position (maintaining the structure and shape of the metal)
Delocalised electrons are mobile and can move through the structure.
Describe the physical properties of metals
High melting and boiling points
High electrical conductivity
What separates metals conductivity to ionic compounds conductivity
Metals can conduct in any state.
Why do metals have such high melting and boiling points?
Large amounts of energy is required to overcome the strong elctrostatic attraction between the cations and electrons.
Solubility of metals
Metals are not soluble
What elements have giant covalent structures
Carbon, Boron, Silicone
what structures are fromed by silicone and carbon and what type of carbon structure?
Tetrahedral structure and diamond structure for carbon.
How are Graphene and Graphite different
Graphine is a single layer, Graphite is multiple layers
Describe Graphene/ Graphites structure
planar hexagonal
why can Graphite/ Graphene conduct electricity?
Only three of carbons 4 outer shell electrons are used in covalent bonds the remaining electron is delocalised and can freely move about the structure
Why are Giant covalent structures not soluble
Polar solvents are not strong enough to break covalent bonds just by their interactions.
Do giant covalent structrures conduct electricity?
Non except for graphite/ graphene
Describe periodic trends in melting points (period 2 and 3)
increasing melting points from Li to Be (giant metallic) to B and C (giant covalent) then rapid decrease to N through Ne (simple molecular) which all have low melting points .
The process is then repeated increasing with Na, Mg and Al (giant metallic) to Si (giant covalent) rapid decreasing in meltong point once P, S, Cl and Ar are reached (simple molecular).
