Ionisation Energy Flashcards
What is ionisation energy
How easily an atom looses electrons tp form positive ion is
1st ionisation energy
The energy required to remove 1 electron from each atom, is 1 mole of gaseous atoms of an element to form 1 mole of gaseou 1+ ions
Is removing an electron from an atom/ ion EXO or ENDO ther mic
ENDOTHERMIC
2nd ionisation energy
The differnence between 1+ and 2+
The energy required to remove one electron from each ion is 1 mole fo gaseous 1+ ions of an element to form 1 mole gaseous 2+ ions
Write the 1st and 2nd ionisation energy equations for al
- Al(g) — al+ (g) + e-
2. Al+(g)— al2+(g) + e-
Why are successive ionisation energies larger
The 2nd ionisation energy of an element is ALWAYS bigger than the 1st ionisation
Due to the ion formed- smaller than the atom
The proton to the electron ratio in the 2+ ion is greater than in the 1+ ion - so the attraction between nucleus adn electron is stronger
What a re the factors that effect ionisation. Energy
Nuclear charge
Electron shielding
Atomic radius
Nuclear charge - factors effecting ionisation energy
The more protons there are in the nucleus, the stronger the attraction between nucleus and electrons
Electron shielding
Inner shel electrons repel outer shell electrons
- repulsion reduces attraction between the nucleus and outer electons - SHIELDING AFFECT
Atomic radius
Greatest effect
Greater distance between the nucleus and outer electons - less nuclear attraction
Trend in ionisation energy down a group
Atomic radius increases
More inner shells so shielding increases
Nuclear attack on on outer electron decreases
1st ionisation energy decrease
Trend across a period
Nuclear charge increases( no. Of protons) Same no. Of shells- shielding the same Nuclear attraction increases Atomic radius decreases First ionisation energy increases
Subs hell structure - affect on ionisation energy
Drop between group 2 and 3
- the 2p subs hell has a higher energy than the 2s subs hell, so the 2p elecron is EASIER to remove- atom will have a lower 1st ionisation energy than 1 with its electron in its 2s sub sell
- the electron is further away in HGIHER energy level
- p has additional shielding
Pairing
Between group 5 and 6
Having electrons PAIRED in an orbital - easier to remove tha as there’s repulsion between the 2 electrons, spin paired = easier to remove
What are the 5 trends in ionisation energy
High ionisation energy of noblel gases- full outer shell
General increase across each period
Drop in ionisation energy between group 2 and 13( subs hell structure)
Drop in ionisation energy between group5 and 16( pairing)
Sharp decrease in 1st ionisation energy between end of 1 period and start of the next
Describe metallic bonding
The electrons in the outermost she’ll of the atom are DELOCALSED
Metal cation electrostatically attracted to the delocalised electrons, form a lattice of closely packed cations in a sea of DELOCALSED electrons
Features of a giant metallic lattice
The more delocalised electrons- the higher the melting and boiling not point ( mg 2+ has 2 electrons per atom so has a HGIHER melting poitn than na+)
Size of the metal io and lattice structure - smaller ionic radius will hold delocalise electrons closer to the nucleus
No bonds holding specific ions together- metal ions can slide past each other when the structure is pulled- a malleable and ductile
Ductile meaning
Can be pulled into a wire
MALLEALE MEANING
Can be hammered into sheets
Are metals insoluble or soluble
Insoluble due to strength of metallic bonds
Trend in MP and BP across period ( 1-13) metals
Ionic charge increases and size decereases
Number of outer shell electrons increases
Attraction increases - mp + BP increases
Giant COVALENTLY structures - properties
High melting and boiling points- strong covalent bonds
Solubility-insoluble as covalent bonds can’t be broke by interactions with solvent s
Electrical conductivity - 2 exceptions of graphite and graphene( only use 3 of carbons outer electrons- delocalised)
BP AND MP across the periodic table
MELTING poitn increases for group 1-14( metallic bonds get stronger, ionic radius decreases number of delocalised electrons increases)
High melting points of group 14 ( carbon and silicon )
Low melting points of group 15-18(week London dispersion forces)
