3.1 the periodic table

Question 1

metallic bond

Answer 1

strong electrostatic attraction between cation (+ ions) and delocalized electrons

Question 2

giant metallic lattice

Answer 2

billions of metal ions held together by metallic bonding (EfA between metal cations and delocalised electrons)

Question 3

giant covalent lattice

Answer 3

3d structure of atoms held together by strong covalent bonds

Question 4

describe the giant covalent structure of diamond?

Answer 4

-C has 4 unpaired valence electrons so can form 4 covalent bonds
-no lone pairs
-tetrahedral structure w bond angles of 109.5*
-C atoms are small > short bonds

Question 5

explain th eproperties of diamond?

Answer 5

-short bonds, very strong covalent bonds, hard, high MP.BP
-all electrons contained in covalent bonds so no free mobile charge carriers, cannot conduct

Question 6

describe the giant covalent structure of graphite?

Answer 6

-each C form 3 covalent bonds and releases 1 delocalised electron
-forms hexagonal layers of C atoms w delcolaized e-s between

Question 7

explain graphites properties?

Answer 7

-each C delocalised one electron, delocalised electrons, mobile charge carriers, can conduct
-soft/slippery (lubricant/pencils) London forces between layers, can slide
-strong covalent bonds, High MP/BP

Question 8

describe the structure of graphene?

Answer 8

-giant covalent structure
-each C atom covalently bonded to 3 other C atoms
-1 atom thick layer of hexagonally bonded C atoms

Question 9

explain the properties of graphene?

Answer 9

-conducts , 1 delocalised electron/ C atom
-high MP/BP, strong covalent bonds

Question 10

describe the structure of silicon?

Answer 10

-each Si has 4 valence electrons and no lone pairs
-forms tetrahedral structure, 109.5*
-atoms larger than C so bonds are longer and weaker, lower MP/BP

Question 11

structure and bonding in silicon dioxide/silica/silicon oxide?
-2 properties?

Answer 11

SiO2
-giant covalent lattice sturcture
-Si atoms form 4 covalent bonds w Oxygen
-tetrahedral

-high MP
-doesnt conduct

Question 12

Boron Nitride

Answer 12

-boron, grp 3, 3 valence electrons
-nitrogen grp 5, 3 to pair and 1 lone pair

-boron is electrodeficient, even when forms 3 bonds outer shell only has 6
-so accepts lone pair from nitrogen and forms dative covalent bond

2 structures:
-4 covalent bonds around each atom- tetrahedral structure (like diamond- “isoelectric”)

-boron form 2 double bonds w N , 1 dative bond w N
-same structure as graphite (no delocalised, cant conduct)

Question 13

define isoelectric

Answer 13

same electronic structure as the Carbons in diamond

Question 14

giant ionic lattice properties explained?

Answer 14

high MP,BP
-large amts energy to overcome strong EfAs between oppositely charged ions in lattice (higher for lattices w greater ionic charges/larger ions)

Dissolve in polar substances
-polar (water) molecules break down the lattice and surround each ion in oslution

Conductivity
-only conduct when (l) or (aq), ionic lattice breaks down, ions free to move and carry a charge
-solid, dont bc ions held in fixed position

Question 15

giant metallic lattice properties?

Answer 15

MP/BP
depends on strength of metallic bond
-MOST high bc high temp to overcome EfA between cations and delocalised electrons

Solubility
-do NOT dissolve, any interaction between solvent and charges in lattice results in reaction NOT dissolving

conductivity
-conduct in (s) or (l), voltage applied across metal, delocalised electrons can move throughout sy=ubstance and carry charge

Question 16

giant covalent lattice properties?

Answer 16

MP/BP
-High, strong cov. bonds between shared e- pair and nucluei of bonded atoms req high amts energy to overcome

Solubility
-Insoluble, cov, bonds too strong to be overcome by interaction w any solvent

Conductivity
-diamond/silicon, NO, all atoms held in rigid structure, no charged particles
-graphene/graphite, delocalised electrons, mobile, carry a charge

Question 17

explain the variation in MP across period 2 and 3 in terms of structure and bonding?

Answer 17

Increase from group 1-3
-ions becoming smaller (decreased atomic radius) and more charged [ions have a higher charge-density moving across the period]
-metallic lattice contains more electrons
-attractions are stronger, MP Increases

Increase from grp3-4
(in period 3) move from metallic structure to giant covalent structure

Question 18

why does atomic radius decrease across a period?

Answer 18

same shielding, increasing nuclear charge pulling outer shell inwards

Question 19

define first ionisation enrgy

Answer 19

energy needed to remove 1mol of electrons from 1mol of gaseous atoms, to form 1mol of gaseous 1+ ions

Question 20

general formula for 1st ionisation energy

Answer 20

X(g) > X^+ (g) + e-

Question 21

define the term second ionisation energy

Answer 21

enrgy required to remove one electron from each ion in 1mol of gaseous 1+ ions of an element, to form 1mol of gaseous 2+ ions

Question 22

what factors effect ionisation energy?

Answer 22

-distance (nucleus to outermost electrons[atomic radius])
-nuclear charge (on the nucleus)
-electron shielding

Question 23

how does distance/atomic radius affect ionisation energy?

Answer 23

increase atomic radius>decreases Ie

-increase atomic radius> decrease force of attraction from nucleus to outermost electrons (nuclear attraction)>decreases Ie

Question 24

how does nuclear charge affect ionisation enegy?

Answer 24

increase nuclear charge> increase Ie

-greater no. protons in nucleus> increases nucluear charge> increases force of attraction nucleus:outermost electrons (nuclear attraction)>increases Ie

Question 25

how does electron shielding effect ionisation energy?

Answer 25

Increase shielding>decrease Ie

-electrons in outer shell repelled by electrons in inner shells, shielding effect decreases nuclear attraction> decreases Ie

Question 26

how do ionisation energies provide evidence for existence of shells?

Answer 26

-very large drops in ionisation enrgy between elements 2-3 (He-Li), 10-11 (Ne-Na) and 18-19 (Ar-K)
-the second element in each of the pairs has its outermost electron in a new shell
-new shell is further from the nucleus and experiences more shielding, both decrease nuclear attraction to outermost electron
-lowering of attraction outweighs increased attraction from extra electron

Question 27

how does ionisation energy provide evidence for subshells?

Answer 27

-all periods general trend of increasing ionisation energies
-increasing nuclear charge (more protons added)
- no extra shielding (same no. sjhells)
- decreasing atomic radius (same shell is pulled inwards
by extra protons)

-however drop between 2nd and 3rd elements 1st ionisation energies
-electron being removed higher energy orbital, easier to remove

Question 28

trend in ionisation energy down a group?

Answer 28

Ie, depends on attraction nucleus to outermost electrons, affected by nuclear charge/shielding/atomic radius

1st ionisation energy decreases down the group
-down a group ATOMIC RADIUS INCREASES, outer shell electron further from nucleus, decreases nuclear attraction
-down a grp, no. internal energy levels increases, SHIELDING EFFECT INCREASES, decreases nuclear attraction
-down grp, nuclear charge increases but effect offset by increased shielding and aromic radius
-overall down a group attraction from nucleus to outermost electrons decreases, less energy to remove electron, 1st Ie decreases

Question 29

general trend in 1st Ionisation energy across a period?

Answer 29

across a period:
-nuclear charge increases as no. protons increases> increases nuclear attraction to outermost electrosn
-increased nuclear charge>decreases atomic radius (outershell pulled closer)> increases nuclear attraction
-same period so shielding effects constant
- increased nuclear charge and decreased atomic radius increases attraction nuclueas to outermost electrons, more energy to overcome, increases 1st Ie.

Question 30

exceptions to trend in 1st Ie across a period?

Answer 30

Be>B
slight decrease
-intro of 2p^1 subshell means that outermost electron in B experiences more shielding which decreases nuclear attraction

N>O, P>S
slight decrease
-moving fromp^3 to p^4 means electron being removed is now paired, electrons repel, decreases nuclear attraction, decreases 1st Ie if

Question 31

explain why successive ionisation energies always increase?

Answer 31

when each electron is removed the outer hell is drawn closer to the nucleus. this decreases atomic radius which increases nuclear attraction, meaning more energy is required to remove the next electron