C7 Periodicity

Question 1

Define periodicity

Answer 1

A repeating trend in properties of the elements across each period of the periodic table.

Question 2

How is the periodic table arranged now?

Answer 2

Arranged in increasing atomic number

Groups (vertical columns)

Periods (horizontal rows)

Question 3

What’s the trend in electron configuration across period 2?

Answer 3

2s sub-shell fill with 2 electrons

Followed by 2p sub-shell with 6 electrons

Question 4

How was the periodic table arranged back then and by who?

Answer 4

Mendeleev

Order of atomic mass (didn’t know about sub atomic particles)

Lined up elements in groups with similar properties

Left gaps (and predicted their properties)

Question 5

What’s the trend in electron configuration across period 3

Answer 5

3s sub shell fill with 2 electrons

Then 2p sub shell with 6 electrons

Question 6

What’s the history of the atoms?

Answer 6

early 1800s, Dalton, tiny hard spheres and couldn’t be divided or splitted

end 1800s, Thomson (discovered tiny neg charged particle ‘electron’) plum pudding model, electrons embedded in ball of pos charge.

10 yrs later, Geiger & Marsden, experiment with alpha particles fired at thin piece of gold (some alpha particles repel, must be tiny spot of pos charge in centre of atom), Rutherford - nuclear model (electrons orbit nucleus which contains protons)

Bohr, suggested electrons orbit nucleus in shells

1932, Chadwick, discovered neutrons (uncharged particles) in nucleus

Question 7

Why is the 1st ionisation energy of oxygen less than 1st ionisation energy of nitrogen

Answer 7

because of electron pairing)

in N & O highest energy e-s in 2p sub shell

in O, paired e-s in 1 of 2p orbitals repel each other, so its easier to remove e- form O than N

Question 8

What’s the trend in 1st ionisation energy down a group?

Answer 8

• atomic radius increases
• more inner shells (shielding increases)
• nuclear attraction on outer electrons decreases
• 1st IE decrease

Question 9

What’s the trend in 1st ionisation energy across a period?

Answer 9

• nuclear charge increases
• similar shielding (same shell)
• Nuclear attraction increases
• atomic radius decreases
• 1st IE increases

Question 10

why is the 1st IE of boron less than the 1st IE of berylium?

Answer 10

(because of 2s and 2p sub shells)

• fall in 1st IE from Be to B marks filling of 2p sub shell
• 2p sub shell in B has higher energy than 2s sub shell in Be
• in B 2p e- easier to remove than 1 of 2s e-s in Be

so 1st IE of B less than Be

Question 11

define first ionisation energy

Answer 11

energy required to remove 1 electron from each atom in 1 mole of gaseous atoms of an element to form 1 mole of gaseous 1+ ions.

Na(g) -> Na+ (g) + e-

Question 12

What factors affect ionisation energy?

Answer 12

Atomic radius

Nuclear charge

Electron shielding

Question 13

Define 2nd ionisation energy

Answer 13

energy required to remove 1 electron form each ion in 1 mole of gaseous 1+ ions of an element to form 1 mole of gaseous 2+ ions.

He+(g) -> He 2+(g) + e-

Question 14

why is the 2nd ionisation energy of helium greater than the 1st?

Answer 14

• 2 protons attract 2 electrons in 1s sub shell
• 1st electron lost, 2nd electron pulled closer to nucleus
• greater nuclear attraction on 2nd electron increases
• more IE needed to remove 2nd electron

Question 15

Define metallic bonding

Answer 15

Strong electrostatic attraction between positive ions and delocalised electrons

Question 16

why can metals conduct electricity in solid and liquid states?

Answer 16

delocalised electrons can move through giant metallic lattice and carry charge

Question 17

what is electrical conductivity of giant covalent lattices?

Answer 17

non conductors of electricity

except graphene and graphite (forms of C)

Question 18

what is the structure and bonding of B, C and Si?

Answer 18

Giant covalent lattice

many billions of atoms are held together by a network of strong covalent bonds

Question 19

what is the solubility of Giant Covalent Lattices?

Answer 19

• insoluble in nearly all solvents
• CBs holding atoms together in lattice are too strong to be broken by interaction with solvents

Question 20

Do giant covalent lattices have high/low MPs & BPs?

Answer 20

high MPs & BPs

high temperatures required to provide large quantity of energy required to break strong CBs

Question 21

why do metals have high melting and boiling points?

Answer 21

high temperatures are necessary to provide the large quantity of energy required to overcome strong metallic bonds (electrostatic attraction between cations and delocalised electrons)

Question 22

what is the solubility of metals?

Answer 22

metals do not dissolve

( there maybe some interaction between polar solvents and charges in a metallic lattice )

Question 23

Carbon and Silicon

Answer 23

group 14 (4)

4 electrons in their outer shell

C (in diamond form) and Si use these 4 electrons to form CBs to other C/Si atoms

Question 24

what is the periodic trend in MPs across period 2 & 3?

Answer 24

MP:
- increase from group 1 to 4
- sharp decrease from group 4 to 5 marks change from giant to simple molecular structures
- are comparatively low from group 5 to 0 giant

metallic structure => giant covalent structure => simple molecular structure

Question 25

Graphene and graphite

Answer 25

• 3 of 4 outer shell e-s used in CB
• remaining e- released into pool of delocalised e-s shred by all atoms in the structure
• good electrical conductors
• giant covalent structures of C
• planar hexagonal layers
• bond angle = 120* by electron-pair repulsiongraphene =
• single layer of graphite
• hexagonally arranged C atoms linked by strong CBs

graphite =
- parallel layers of hexagonally arranged C atoms
- layers bonded by weak LFs
- 3 of 4 outer shell e-s used in CB
- remaining e- released into pool of delocalised e-s shred by all atoms in the structure

Question 26

Metallic bonding and structure

Answer 26

electrons delocalise and can move

cations are fixed in place and maintain structure and shape of the metal

Question 27

what is the name of the structure that contains metallic bonding?

Answer 27

Giant metallic lattice

Question 28

what is the name given to elements near to the metal/non-metal divide on the periodic table?

Answer 28

semi metals or metalloids

e.g. B, Si, Ge, As, and Sb

Question 29

what are the properties of most metals?

Answer 29

• strong MBs
• high electrical conductivity
• high MPs & BPs

Question 30

why can’t silicon and carbon (diamond) conduct electricity?

Answer 30

in both structures, all 4 outer shell electrons are involved in CB, so none are available for conducting electricity.

Question 31

what do many non-metallic elements exist as?

Answer 31

simple covalently bonded molecules

in solid state, they form simple molecular lattice structures which are held together by weak intermolecular forces - theses structures have low MPs and BPs

Question 32

what is the bond angle and structure of (carbon) Diamond?

Answer 32

• tetrahedral arrangement
• bond angle = 109.5*

because of electron pair repulsion