3.1.1 Periodicity Flashcards
How many sections is the periodic table divided into
What are they
4 (s,p,d+f block)
Where is s block
Furthest to left and 2 upper sections
Where is p block
Furthest to right
Where is d block
Middle block
Where is f block
Bottom block
Where are the metals + non-metals around the stepped line
Metals are left of stepped line
Non-metals to the right of stepped line
What are borderline elements classed as as they’ve got both metallic + non-metallic properties e.g silicon
Metalloids/semi-metals
What’s a group
Vertical column of elements going down
What’s a period
Horizontal row of elements going across
What does periodic mean
Recurring regularly (a pattern)
What are the 3 trends we look at
1st ionisation energy across a period
Atomic radius
Melting + boiling points across a period
What 3 factors can explain al ionisation energy trends
Atomic radius (atom size) Nuclear charge (proton number) Shielding
What does shielding take priority over
What does this mean
Shielding takes priority over increasing nuclear charge so down a group the ionisation energy increases as shielding increases, despite increased nuclear charge
What’s the radius of an atom defined as
Half the distance between the centres of pairs of identical atoms, covalently bonded
What happens to atomic radius across a period
Why
It decreases as protons are getting added but shielding stays the same so there’s a stronger attraction between electrons + nucleus
What’s happens to atomic radius down a group
Why
It increases as the elements get 1 extra shell of electrons compared to the one before, so the outermost electrons are further from the nucleus
Are atomic radius of metal atoms/ their corresponding ions smaller
Why
Ions
As they lose an electron, so attraction to nucleus is stronger
Why are non-metal ions larger than their corresponding atoms
As they gain an electron , so attraction to nucleus slightly decreases
What’s the only reason a boiling point is low
Only have to break intermolecular forces within simple covalent structures
What have higher melting and boiling points (2)
Why
Metals
Form giant metallic lattices, with strong attraction between metal ions and delocalised electrons
Giant covalent (macromolecular) lattice Have many strong covalent bonds that must be broken
What is ionisation
Removing an electron from an atom, forming positive ion
Define 1st ionisation energy
Eq
Energy needed to remove 1 mole of electrons from 1 mole of atoms in the gaseous state
M(g) -> M+(g) + 1e-
Why are 1st + 2nd ionisation energies both endothermic
As energy must be put in to remove the negative electron from the nucleus’ attractive influence
Define 2nd ionisation energy
Eq
Energy needed to remove 1 mole of electrons from 2 Mike of singly positively charged ions in the gaseous state
M+(g) -> M2-(g) + 1e-
Why is the 2nd IE always greater than the 1st
As the electron is being removed from an already positively charged ion so is more strongly held
How quickly are the electrons close to nucleus removed
Last (1s2)
What does ionisation energy increase with
Why
Removal of each electron
As although charge stays same, remaining electrons feel charge stronger as electrons are being removed from a more positive species
Why is there a big jump between 1st IE and 2nd
As it’s harder to remove the 2nd electron than the first - 2nd e- is closer to nucleus so has a stronger attraction due to less shielding
Why do further out electrons feel less nucleus attraction
As there are more inner shells shielding them
What does a big jump in energy mean
Moving to a shell closer to nucleus and electrons get harder to remove
Why is electron pairing in an orbital mean electrons are easier to remove with less energy
As this creates repulsion between electrons
Why does 1st IE decrease going down a group
Shielding increases so outer electrons feels a weaker nuclear attraction and needs less energy to remove
3 main factors that affect ionisation energy
Atomic radius
Nuclear charge
Shielding
What’s the ionisation energy like if the atomic radius is bigger
Lower IE as weaker attraction to nucleus
How does nuclear charge affect ionisation energy
Stronger nuclear charge, stronger the attraction so greater ionisation energy
How does shielding affect ionisation energy
Greater shielding, weaker attraction to nucleus so lower the ionisation energy
What affects IE going across a group
Increasing nuclear charge
