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Inorganic Chemistry Year 12 > 3.1.1 Periodicity > Flashcards

3.1.1 Periodicity Flashcards

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1
Q

what order is the periodic table arranged in

A

increasing atomic number

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2
Q

what are the vertical columns called

A

groups (1 to 18)

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3
Q

what are the horizontal rows called

A

periods

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4
Q

what are the blocks

A

correspond to electron found on highest subshell

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5
Q

what is periodicity

A

there are repeating trend in properties across each period

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6
Q

what is ionisation energy

A

how easily an atom loses electrons and forms positive ions

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7
Q

what is first ionisation energy

A

the energy needed to remove one electron from each atom in a mole of gaseous atoms of an element to form one mole of gaseous 1+ ions

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8
Q

what are the ionisation energy equations

A

Na(g) ==> Na(g)1+ + e-
Na(g)1+ ==> Na(g)2+ e-

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9
Q

which 3 factors effect the first ionisation energy of an element

A

1) atomic radius
2) nuclear charge
3) electron shielding

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10
Q

what is atomic radius and how does it effect ionisation energy

A

-greater the distance between outer electrons
-less nuclear attraction
-electrons easily lost

-LOWEST AT TOP RIGHT OF PERIODIC TABLE:
- across period: more protons, means more positive charge of nucleus, means electrons are pulled closer together, reducing atomic radius

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11
Q

explain nuclear charge

A

-more protons in nucleus
- greater nuclear attraction

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12
Q

explain electron shielding

A
  • inner shell electrons repel outer electrons (all negative)
  • causes repulsion called shielding effect
  • reduces nuclear attraction
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13
Q

what is nuclear attraction

A

attraction between positive nucleus and outer electrons

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14
Q

how many IE’s will an atom have

A
  • as many electrons as it has
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15
Q

why does IE increase with successive energies

A
  • less electrons, e- pull closer to nucleus and is nearer it, more IE
  • as size of ion is smaller
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16
Q

what is the definition of 2nd ionisation energy

A

the amount of energy needed to remove one electron from from each 1+ ion in one mole of gaseous ions of an element and form one mole of 2+ gaseous ions

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17
Q

how can the ionisation energy graph tell you which element we are looking at

A
  • number of ionisation energies = number of electrons = number of protons
  • large jump between ionisation energies = moving between shells = tells you number of electrons on outer shell = group number
18
Q

what do the periodic trends of ionisation energy show proof of

A

shells and subshells

19
Q

what is the trend in ionisation energy

A
  • decreases down a group
  • increases across a period
    -(sharp decrease from end of one period to start of the next one)
20
Q

why does ionisation energy decrease down a group

A

1) atomic radius increases, so outer electrons further from nucleus, so lower nuclear attraction
2) more electron shielding, lowering NA
3) nuclear charge increases, but outweighed by other two factors

21
Q

why does ionisation energy increase across a period

A

1) atomic radius decreases
2) electron shielding stays the same
3) nuclear charge increases

  • so NA increases, and harder to lose electron
22
Q

what are the 2 exceptions to the IE rule

A
  • between group 2 and 3, and group 5 and 6
23
Q

why does IE decrease from group 2 to 3

A
  • electron in group 3 is in the p subshell
  • has a higher energy state
  • electron is easier to remove
  • despite increased nuclear charge
24
Q

why does IE decrease from group 5 to 6

A
  • electron in group 5 is unpaired in the p orbital
  • electron in group 6 is paired in a p orbital
  • electron pair repels
  • and is at higher energy state
  • so easier to remove electron
  • despite increased nuclear charge
25
Q

what is metallic bonding

A

strong electrostatic forces of attraction between cations and delocalised electrons

26
Q

what is the position of subatomic particles in metals

A
  • cations are in fixed positions, maintaining structure and shape
  • delocalised electrons are in a mobile pool
27
Q

how are metals held

A

in a giant metallic lattice

28
Q

what and why is the melting point of metals

A

HIGH: (most solid except mercury)
- large amount of energy needed to overcome the strong electrostatic forces of attraction between cations and delocalised electrons
- increases across a period: more delocalised electrons per each ion (e.g. 3- each), higher nuclear charge and smaller atomic radius means stronger electrostatic forces of attraction (referred to as high charge density metals) (still all have same, metallic bonding)

29
Q

explain the conductivity of metals

A

VERY GOOD:
- have delocalised electrons that can move through the structure and act as mobile charge carriers

30
Q

explain the solubility of metals

A

INSOLUBLE: (as atoms)
- metal would react with solvent before chance to dissolve
- also strong metallic bonding

31
Q

which atoms form giant covalent lattices

A

B, Si, C

32
Q

what is the structure of diamond and silicon

A
  • formed of 4/4 available carbon bonds
  • tetrahedral shape
    -109.5 degree bond angle
33
Q

explain melting point of giant covalent lattices

A

HIGH:
very strong covalent bonds need to be overcome

34
Q

explain solubility of giant covalent lattices

A

INSOLUBLE:
- covalent bonds holding the lattice together are too strong to be broken down by interactions with solvent

35
Q

explain the conductivity of giant covalent lattices

A

DO NOT:
- no ions or delocalised electrons to move and carry charge

36
Q

explain the exceptions to the inconductivity of giant covalent lattices

A

GRAPHENE AND GRAPHITE:
- made of hexagonal sheets
- have delocalised electron
- can conduct electricity

GRAPHITE: single layer
GRAPHENE: sheets joined by weak london forces, so slippery

37
Q

which element make up simple molecular lattices

A

non metals

38
Q

explain the boiling point of simple molecular lattices

A

LOW:
- only need to overcome the weak london dispersion forces between molecules, which don’t require a lot of energy to overcome
- more electrons= more LDF = higher boiling point, so generally increases across period (NOT S8)

39
Q

explain the conductivity of simple molecular lattices

A

NONE:
no delocalised electrons or ions
( all electrons are in localised bonds)

40
Q

explain the solubility of simple molecular lattices

A

ONLY IN NON-POLAR:
- not in water, as polar
- LF can form between molecule and non-polar molecules like hydrocarbons

41
Q

explain the particles and forces in types of structures

A

METALLIC= metallic bonding, cations, delocalised electrons
GIANT COVALENT= atoms, covalent bonds
SIMPLE MOLECULAR= molecules, london forces

Inorganic Chemistry Year 12 (5 decks)

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