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a level chemistry > Module 3.1 The perodic table > Flashcards

Module 3.1 The perodic table Flashcards

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1
Q

in the early 1800s what where the only two ways to categorise elements

A
  • by their physical and chemical properties

- by their relative atomic mass

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2
Q

explain Dobereiner’s triads

A

he group elements with similar characteristics in threes and called them triads.

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3
Q

explain the law of octaves

A

Newlands order elements by mass and found similar elements appeared at regular intervals every 8th element was similar.

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4
Q

Mendeleev produced the first accepted version. what did his table contain

A

arranged elements by atomic masses.
left gaps where the next element didn’t seemed to fit.
he predicted the properties of undiscovered elements.

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5
Q

what do all elements in a period have in common

A

they all have the same number of electron shells (although not sub-shells)
period 1 has 1 electron shell and so on.
repeating trends in chemical and physical properties across the period due to this.

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6
Q

the trends that elements share in periods are known as

A

periodicity

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7
Q

what do all elements in the same group have in common

A

they all have the same number of electrons in their outer shell.
this causes them to also have similar CHEMICAL properties.

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8
Q

what does the group number tell you?

A

the number of electrons in the outer shell

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9
Q

s block elements contain which groups

A

groups 1 and 2

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10
Q

what is the outer shell electron configuration of elements in the s block

A

S(1) or S(2)
e.g Mg is 1s(2) 2s(2) 2p(6) 3s(2)
Li is 1s(2) 2s(2)

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11
Q

what part of the perodic table is the d block

A

the transition metals

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12
Q

what is the electron configoration of d block elements

A

outer shell is always a d sub shell
e.g cobalt: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6) 3d(7) 4s(2)
(even though 4s fills first and so the last sub shell is 3d it is not written this way)

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13
Q

the p block contain which groups of elements

A

groups 3-7 and 0

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14
Q

what is the outer shell electron configurations of elements in the p block

A 
from s(1)p(1) to s(2)p(6)
e.g Cl: 1s(2) 2s(2) 2p(6) 3s(2) 3p(6)
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15
Q

period 1 can have what outer sub shell configeration

A

up to 1s(2)

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16
Q

period 2 can have what outer sub shell configeration

A

up to 2s(2) 2p(6)

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17
Q

period 3 can have what outer sub shell configeration

A

up to 3s(2) 3p(6)

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18
Q

ionisation is

A

the removal of one or more electrons

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19
Q

the first ionisation energy is

A

the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms.

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20
Q

you put energy in to ionise an atom so it is an

A

endothermic process

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21
Q

give the equation of the first ionisation of oxygen

A

O(g) –> O+(g) + e-

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22
Q

the lower the ionisation energy

A

the easier it is to form an ion

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23
Q

factors affecting ionisation energy are (3)

A
  • NUCLEAR CHARGE
  • ATOMIC RADIUS
  • SHIELDING
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24
Q

nuclear charge affects ionisation energy how?

A

the more protons there are in nucleus the more ve+ charged the nucleus is and the STRONGER THE FORCE OF ATTRACTION for the electrons

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25
Q

atomic radius affects ionisation energy how?

A

attraction falls off very rapidly with distance. an electron close to the nucleus will be much more strongly attracted than one further away.

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26
Q

shielding affects ionisation energy how?

A

as the no. of e- beteween outer e- and the nucleus increases the outer e- feel less attraction towards the nuclear charge.

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27
Q

a high ionisation energy means there is a

A

strong attraction between the e- and the nucleus so more energy is needed to overcome the attraction and remove the electron.

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28
Q

how does ionisation energy change as you go down a group

A

ionisation energy usually falls i.e it gets esier to remove outer electrons

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29
Q

ionisation energy decreases down a group because (2)

A
  • ATOMIC RADIUS elements further down have more
    electron shells. increasing atomic radius so outer
    electrons are further away from nucleus
  • SHIELDING the extra inner shells shield the outer
    electrons from the attraction of nucleus
    -although protons ve+ charge increases so does the
    electrons ve- charge so there is no effect
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30
Q

how does ionisation energy change as you go across a period

A

ionisation energy increases as you go across a period

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31
Q

why does ionisation generally increase as you go across a period (2)

A
  • NUCLEAR CHARGE number of protons is increasing. as
    ve+ of nucleus increases electrons are pulled closer
    reducing ATOMIC RADIUS
  • extra electrons gained are added to the outer energy
    level so don’t provide extra shielding effect
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32
Q

what are the 2 exceptions to the trends in ionisation energy on the period table?

A
  • the drop between groups 2 and 3

- the drop between groups 5 and 6

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33
Q

why is there a drop in ionisation energy between groups 2 and 3

A

SUB-SHELL STRUCTURE

  • outer electron in group 3 is in a p orbital rather than s
  • p orbital has slightly higher energy than s in the same shell
  • so electron is found further away from the nucleus
  • these factors override the effects of the increasing nuclear charge resulting in ionisation energy dropping slightly
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34
Q

why is there a drop in ionisation energy between groups 5 and 6?

A

p ORBITAL REPULSION

  • in group 5 elements the electron is being removed from a singly occupied orbital
  • in group 6 elements the electron is being removed from an orbital containing 2 e-
  • the repulsion between the 2 electrons means that the e- is easier to remove from shared orbitals.
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35
Q

what does successive ionisation energies involve

A

involves the removal of additional electrons

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36
Q

the equation for the 2nd ionisation of oxygen is:

A

O+(g) –> O2+(g) + e-

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37
Q

successive ionisation energies can show

A

the shell structure of an element

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38
Q

how does successive ionisation energies show shell structure

A
  • within each shell, successive ionisation energies will increase as you are removing e- from increasingly +ve ion and there is less repulsion amongst remaining electrons
  • big jumps in ionisation energy happen when a new shell is broken into as electron is being removed in a shell closer to the nucleus.
  • this shows shells.
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39
Q

successive ionisation energies can also predict

A

electronic structure of an element.

by counting the no, of electrons removed before each big jump tells u how many e- are in each shell.

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40
Q

diamond, graphite and graphene are…

A

giant covalent lattices

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41
Q

what are giant covalent lattices

A
  • huge networks of covalently bonded atoms.

- carbon atoms can form this type of structure because they can each form 4 strong covalent bonds

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42
Q

what are different forms of the same element in the same state called

A

allotropes

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43
Q

name and explain the properties of diamond (5)

A
  • very high melting point due to strong covalent bonds
  • diamond is the strongest known substance due to giant covalent lattice tetrahedral shape each carbon bonded to 4 others
  • good thermal conductor. (vibrations travel easily)
  • can’t conduct electricity (all outer electrons are held in localised bonds.
  • it wont solved in any solvent
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44
Q

name a element with very similar properties to carbon and what it can form

A

silicon

  • in same periodic group
  • also forms crystal lattice with similar properties to carbon
  • can form 4 strong covalent bonds
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45
Q

name and explain the properties of graphite (4)

A
  • feels slippery:
    weak forces between layers in graphite are easily
    broken so that sheets can slide over each other.
  • conduct electricity:
    delocalised electrons aren’t attached to any particular
    carbon and are free to move along sheets.
  • very high melting point:
    because of strong covalent bonds in hexagonal sheets
  • insoluble:
    covalent bonds are too strong to break
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46
Q

describe the structure of graphene

A

is a sheet of carbon atoms joined in hexagons.

the sheet is one atom thick making it 2-dimesonal

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47
Q

describe and explain the properties of graphene (3)

A
  • conducts electricity:
    de-localised electrons free to move along the sheet. the single layer allows quicker movement above and below sheet making it the best known conductor
    -extremely strong:
    de-localised electrons strength covalent bonds
    -transparent and light
    due to it only being a single layer
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48
Q

metal can have giant structures called

A

giant metallic lattices structures.

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49
Q

describe how giant metallic lattices structures form (metallic bonding):

A
  • e- in outer shell are delocalised leaving a +ve metal cation
  • metal cations are electrostatically attracted to the delocalised e-.. they form a lattice of closely packed cations in a sea of delocalised electrons.
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50
Q

describe and explain the properties of metals in terms of metallic bonding (5)

A
  • MELTING POINT effect by no. of delocalised e- per atom
    the more there are the stronger the bonding and higher the melting point. the size of metal ion and lattice structure also have an effect.
  • MALLEABLE AND DUCTILE:
    no bonds holding specific ions together they can slide past each other.
    -GOOD THERMAL CONDUCTORS
    delocalised e- can pass kinetic energy to each other
    -GOOD ELECTRICAL CONDUCTORS
    delocalised electrons can move and carry current
    -INSOLUBLE
    except in liquid metals, because of the strength of metallic bonds
51
Q

what factors affect melting and boiling points of simple molecular structures

A

Melting and boiling points depend on strength of induced dipole dipole forces
the more atoms their are the stronger the induced dipole- dipole forces.
noble gases have very low melting and boiling points as they exist as individual atoms resulting in very weak induced dipole-dipole interactions.

52
Q

summaries properties ionic bonding result in:

A 
- high melting and boiling points
solid at STP
does not conduct electricity as solid
conducts electricity as liquid 
not soluble in water
53
Q

summaries properties simple molecular (covalent) bonding results in:

A

low melting and boiling points (only have london forces
at STP usually liquid or gas occasionally solid
does not conduct electricity as a solid or liquid
solubility depends on how polarised the molecule is

54
Q

summaries properties of giant covalent lattices

A

high melting and boiling points
solid at STP
does not conduct electricity when solid apart from graphite and graphene
does not conduct as liquid (not normally in this state)
is not soluble in water

55
Q

summaries properties of metalic bonding

A

high melting and boiling points
solid at STP
conducts electricity as solid and liquid (delocalised e-)
are not soluble in water

56
Q

going across the period how does the melting and boiling points change for metals (Li, Be, Na, Mg, Al)eland why

A

melting and boiling points increases across a period

because the metallic bonds get stronger as the ionic radius decreases and the number of delocalised electrons increases.

57
Q

elements with giant covalent lattice structure (C and Si) have strong ________ so have _____ melting and boiling points.

A

elements withe giant covalent lattice structure have strong covalent bonds so have high melting and boiling points.

58
Q

elements that form simple molecular structures only have weak ____________ to overcome between molecules so have ___ melting and boiling points.

A

elements that form simple molecular structures only have weak intermolecular forces to overcome between molecules so have low melting and boiling points

59
Q

noble gases have the ________ melting and boiling points in their periods because ___

A

noble gases have the lowest melting and boiling points in their periods because they are held together by the weakest forces

60
Q

group 2 elements form __ ions because

A

group 2 elements form 2+ ions because they have two e- in their outer shell s(2)
so they lose two outer e- to form 2+ ions

61
Q

how does reactivity change as you go down group 2 and why

A

reactivity increases down group 2

  • ionisation energies decreases because:
    1. increasing atomic radius
    2. shielding effect
  • easier to lose electrons to form +ve cations
62
Q

when group 2 elements react they are

A

oxidised
from state of 0 to +2 forming M2+ ions
M –> M2+ +2e-
0 +2

63
Q

group 2 metals + water —>

A

group 2 metals + water –> metal hydroxide + hydrogen

M + 2H2O —> M(OH)2 + H2

64
Q

group 2 metals + oxygen —>

A 
group 2 metals + oxygen --> metaloxides 
(they burn in oxygen)
2M + O2 ---> 2MO
0                     +2 
           0             -2
65
Q

group 2 metals + acid –>

A

group 2 metals + acid(dilute) —-> salt + hydrogen

M + 2HCl –> MCl2 + H2

66
Q

the group 2 metals are known as

A

the alkaline earth metals

67
Q

group 2 oxides and hydroxides are what:

A

they are bases

most are also soluble in water, so are also alkalis

68
Q

group 2 oxides + water —>

A

metal hydroxides
which dissolve.
hydroxide ion OH- make the solution strongly alkaline

69
Q

what group 2 oxide is the exepetion

A

magnesium oxide

only reacts slowly and the hydroxide it forms is not very soluble.

70
Q

how does the strength of the alkali solution group 2 oxides make as you go down the group

A

as you go down the group the more strongly alkaline solution is. because they hydroxides get more soluble.

71
Q

group 2 compounds are often used to

A

neutralise acidity

72
Q

what group 2 hydroxide is used in agriculture to neutralise acidic soils

A

calcium hydroxide (slaked lime, Ca(OH)2 )

73
Q

what 2 group 2 compounds are commonly used in indigestion tablets as antacids

A

magnesium hydroxide Mg(OH)2

calcium carbonate CaCO3

74
Q

what is the equation for neutralisation

A

H+ + OH- —> H2O

75
Q

neutralisation equation for magnesium hydroxide

A

Mg(OH)2 + 2HCl —-> 2H2O + MgCl

76
Q

give fluorine’s: formula, colour, physical state at 20oc and electronic structure

A

formula: F2
colour: pale yellow
sate: gas
1s(2) 2s(2) 2p(5)

77
Q

give chlorine’s: formula, colour, physical stat at 20oc and electronic structure

A

formula: Cl2
colour: green
state: gas
1s(2) 2s(2) 2p(6) 3s(2) 3p(5)

78
Q

give bromine’s: formula, colour, physical state at 20oc and electronic structure

A

formula: I2
colour: red-brown
state: liquid
1s(2) 2s(2) 2p(6) 3s(2) 3p(6) 3d(10) 4s(2) 4p(5)

79
Q

give Iodine’s: formula, colour, physical state at 20oc and electronic structure

A

formula: I2
colour: grey
state: solid
1s(2) 2s(2) 2p(6) 3s(2) 3p(6) 3d(10) 4s(2) 4p(6) 4d(10) 5s(2) 5p(5)

80
Q

halogens exist as _______ molecules

A

halogens exist as diatomic molecules

81
Q

what happens to halogen boiling and melting points as you go down the group

A

their boiling and melting points increase

82
Q

why does halogen boiling and melting points increases as you down the group?

A
  • due to increasing strength of the London forces as the size and relative mass of the atoms increases
83
Q

what happens to the physical state of halogens as you go down the group

A

they halogens become less volatile as you go down the group.

84
Q

what happens to halogen reactivity as you go down the group

A

halogens get less reactive as you go down the group

85
Q

why do halogens get less reactive as you go down the group?

A
  • as you go down atomic radii increases so outer e- are further from the nucleus.
  • shielding increases as there are more inner electrons
  • makes it harder for atoms to attract the electron they need to form an ion, so the larger atoms are less reactive
86
Q

how do halogens react?

A

halogens react by gaining an electron in their outer shell to form 1- ions.
this means they are reduced and are oxidising agents.

87
Q

what is another way to say that halogens get less reactive as you go down the group

A

halogens become less oxidising as you go down the group.

88
Q

how do halogen displacement reactions work

A

halogens with a larger oxidising strength displace halogens with lower oxidising strengths.
e.g fluorine will replace bromine.

89
Q

write the full and ionic equation with oxidation numbers to show what happens when bromine water displaces potassium iodine

A

full: Br2 + 2KI —> 2KBr + I2
ionic: Br2 + 2I- —-> 2Br- + I2
Br 0 -1
I -1 0

90
Q

bromine displaces the iodine ions ________ them

A

bromine displaces the iodine ions oxidising them

91
Q

chlorine will displace what halogens

A

chlorine will displace bromide and iodide

92
Q

bromine will displace what halogens

A

bromine will displace iodide

93
Q

iodine will displace what halogens

A

iodine will not displace any halogens

94
Q

fluorine will displace what halogens

A

fluorine will displace chloride, bromide and iodide

95
Q

what solution can be used to test for halogens

A

silver nitrate solution

96
Q

how do you test for halogens

A
  • first add dilute nitric acid to remove ions that might interfere
  • then add silver nitrate solution (AgNO3)
  • if halogen present a precipitate will form (silverhalide)
  • identify with colour
  • add ammonia solution do see how well the siverhalide dissolves to help identity
97
Q

testing for chlorine with silver nitrate what will form

A
  • a white precipitate

- dissolves in dilute NH3

98
Q

testing for bromine with silver nitrate what will form

A
  • cream precipitate

- dissolves in conc. NH3

99
Q

testing for iodine with silver nitrate what will form

A
  • yellow precipitate

- insoluble in conc. NH3

100
Q

what is the ionic equation when a halide ion reacts with silver nitrate

A

Ag+ + X- —> AgX

101
Q

halogens undergo ___________ with alkalis

A

halogens undergo disproportionation with alkalis

102
Q

in disporportionation reactions what happens

A

something is simultaneously oxidised and reduced

103
Q

give the full and ionic equation with oxidation numbers for when a halogen reacts with sodium hydroxide NaOH

A

X2 + 2NaOH —> NaXO + NaX + H2O
X2 + 2OH —> XO- + X- + H2O
0 +1 -1

104
Q

what are the 3 different oxidation states chlorine can be at

A
  • chloride Cl- -1
  • chlorine Cl2 0
  • chlorate ClO- +1
105
Q

all halogens can undergo disportionation reactions apart from…

A

fluorine

106
Q

chlorine and sodium hydroxide make what

A

chlorine and sodium hydroxide make bleach

107
Q

what is common household bleach

A 
sodium chlorate(I) solution
NaClO
108
Q

the equation for when chlorine and sodium hydroxide react to make bleach as well as oxidation numbers of chlorine

A

2NaOH + Cl2 —-> NaCLO + NaCl + H2O

0 +1 -1

109
Q

write the equation for when chlorine is mixed with water

A

Cl2 + H2O HCl + HClO
0 -1 +1
chloric (I) acid

110
Q

chlorine undergoes _________ when mixed with water

A

chlorine undergoes disproportionation reaction with water

111
Q

why is chlorine / compound with chlorate (I) ions mixed with drinking water

A

to kill bacteria to make water safe to drink / swim in

112
Q

write the equation for when acid ionises to make chlorate(I) ions

A

HClO + H2O ClO- + H3O+

113
Q

why is chlorine an important part of water treatment (3):

A
  • it kills disease causing microorganisms
  • some chlorine remains in the water and prevents reinfection further down the supply
  • it prevents growth of algae eliminating bad tastes, smells and removes discolouration caused by organic compounds
114
Q

what are the risks of using chlorine in water (3)

A
  • chlorine gas irritates respiratory system.
  • liquid chlorine causes severe chemical burns so accidents involving chlorine are serious, fatal.
  • chlorine reacts with organic compounds in water forming chlorinated hydrocarbons many are carcinogenic.
115
Q

what are the alternatives to chlorine (2)

A
  • ozone (O3) a strong oxidsing agent kills microoganisms expensive and has a short half life so treatment is not permanent
  • uv light: damage microorganisms DNA, ineffective in cloudy water wont stop later contamination.
116
Q

how do you test for carbonates (CO32-)

A
  • add dilute acid if present then carbon dioxide will be released
  • test for carbon dioxide using lime water
117
Q

what is the equation for the reaction between carbonate ion and dilute acid

A

CO3(2-) + 2H+ –> CO2 + H2O

118
Q

how do you test for sulfates?

A
  • add dilute HCl followed by barium chloride solution, BaCl2
  • barium sulfate forms a white precipitate
119
Q

what is the equation for sulfate tests

A

Ba2+ + SO4(2-) —> BaSO4

120
Q

how do you test for halides

A
  • add nitric acid, then add silver nitrate

- add ammonia to see solubility

121
Q

how do you test for ammonium compounds

A
  • ammonia gas NH3 is alkaline. damp red litmus paper will turn blue
  • to test for ammonium ions NH4+ add sodium hydroxide NaOH and warm the mixture. test for ammonia given off
122
Q

what is the equation for testing for ammonium ions

A

NH4+ + OH- —-> NH3 + H20

123
Q

what order do you test for substances in

A
  1. test for carbonates
  2. test for sulfates
  3. test for halides

a level chemistry (5 decks)

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