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Chemistry topic 3 > Periodicity > Flashcards

Periodicity Flashcards

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1
Q

how is the periodic table arranged?

A
  • by increasing atomic (proton) number
  • in periods showing repeating trends in physical and chemical properties (periodicity)
  • in groups having similar chemical properties
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2
Q

define the term periodicity

A

the periods show repeating trends in physical and chemical properties

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3
Q

what is the definition for the first ionisation energy?

A

ionisation energy is the minimum amount of energy required to remove 1 mole of electrons from 1 mole of gaseous atoms

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4
Q

what is the first ionisation energy for magnesium?

A

Mg(g) ==> Mg+(g) + e-

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5
Q

rules for writing ionisation energy

A

1- always include (g) state symbols
2- ionisation requires energy so they are always an endothermic process and have a positive value

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6
Q

what is shielding?

A

the more electrons shells between the positive nucleus and negative electron that is being removed the LESS ENERGY that is required as there is a weaker attraction

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7
Q

how does atomic radius affect ionisation energy?

A

The bigger the radius, the further away the outer electrons are from the nucleus so the attractive force between nucleus and outer electron reduces and it’s easier to remove electrons.

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8
Q

how does nuclear charge affect ionisation energy?

A

the more protons in the nucleus, the bigger the attraction between the nucleus and outer electrons. This means that more energy is required remove the electron

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9
Q

when there is a high nuclear charge but there is also high shielding, how does this affect ionisation energy?

A

the shielding counteracts the nuclear charge and therefore has no impact on the ionisation energy and it is still easier to remove the electrons as they increase

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10
Q

what is the trend in first ionisation energy down a group?

A

Ionisation energy decreases down a group
- the atomic radius increases down the group, outer electrons are further from the nucleus
- shielding increases down the group as there are more shells between the nucleus and outer shell
==> therefore, attractive forces are weaker and the energy required to remove an electron decreases

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11
Q

what is the general trend in first ionisation energy across a period? (period 2 and 3)

A

The ionisation energy increases across a period
-across the period there is an increasing number of protons in the nucleus which increases the nuclear attraction
-shielding is similar and distance from the nucleus marginally decreases
- more energy is required to remove an outer electron as ionisation energy increases

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12
Q

what is the first exception for trend in first IE in period 3

A

Aluminium:
the outer electron in Al sits in a higher energy sub-shell slightly further from the nucleus than the outer electron in magnesium
==> Mg = 1s2, 2s2, 2p6, 3s2
==> Al = 1s2, 2s2, 2p6, 3s2, 3p1
Al outer electron is in the 3p sub shell whereas Mg outer electron is in the 3s sub shell
(therefore provides evidence for sub-shells)

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13
Q

what is the second exception in first ionisation energy in period 3?

A

Sulfur:
A decrease in 1st IE at sulfur is evidence for electron repulsion in an orbital
- both S and P have outer electrons in the 3p orbital so the shielding is the same
- removing an electron from S involves taking it from an orbital with 2 electrons in unlike in P with 1
- electrons repel each other so less energy is needed to remove an electron from an orbital with 2 in than 1

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14
Q

what is successive ionisation? (second, third, fourth etc IE

A

the removal of more than 1 electron from the same atom

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15
Q

why does each successive ionisation require more energy?

A

you are removing electrons from a shell closer to the nucleus
- general increase in energy as removing an electron from an increasingly more positive ion

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16
Q

what are examples of giant covalent structures

A

graphite and diamond

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17
Q

what is the structure of graphite?

A
  • each carbon bonded 3 times with 4th electron delocalised
  • lots of strong covalent bonds so very high mp and bp
  • layers slide easily as there are weak forces between the layers
  • delocalised electrons mean it can conduct electricity
  • layers are far apart in comparison to covalent bonds so low density
  • graphite is insoluble as the covalent bonds are too strong to break
18
Q

what is the structure of diamond / silicon?

A
  • each carbon bonded four times in tetrahedral shape
  • tightly packed, rigid arrangement means heat is a good conductor
  • very high mp and bp due to many strong covalent bonds and very hard
  • doesn’t conduct electricity
  • insoluble as covalent bonds are too strong to break
19
Q

what is the structure of graphene?

A
  • 1 layer of graphite so 1 atom thick
  • lightweight and transparent
  • delocalised electrons so conductor of electricity
  • delocalised strengthen the covalent bonds which gives high strength
20
Q

what is metallic bonding?

A

strong electrostatic attraction between cations and delocalised electrons

21
Q

describe metallic bonding

A
  • positive metal ions form as metals donate electrons to form a ‘sea’ of delocalised electrons
  • electrostatic attraction between cations and negative delocalised electrons
22
Q

how does the melting point increase in giant metallic lattice structures (metals)

A

the more electrons an atom can donate to the delocalised system, the higher the melting point
==> eg, Mg has higher melting point than sodium as it can donate 2 electrons (group 2) whereas Na only donates 1 per atom (group 1)

23
Q

what are the features of metals

A
  • metals are good thermal conductors as the delocalised electrons can transfer kinetic energy
  • metals are good electrical conductors as the delocalised electrons are mobile and can carry a charge
  • high mp and bp because strong electrostatic attractions
  • solid metals are insoluble as the metallic bond is too strong to break
24
Q

what is the trend in melting points across period 3 and period 2?

A
  • first three elements are metals so they have metallic bonding
  • general increase in melting points as metal ions have increasing positive charge so an increasing number of delocalised electrons and a smaller ionic radius which causes a stronger metallic bond
25
Q

in period 3, why does silicons and in period 2, why does carbons melting point increase so much more than all other elements?

A

highest melting point because it has a giant covalent (macromolecular) structure
- networks of atoms bonded by strong covalent bonds so lots of energy required to overcome the strong bonds

26
Q

why is there a huge decrease in melting points after silicon and carbon?

A

phosphorus and nitrogen have a weaker simple molecular structure
- melting point is determined by weaker induced dipole-dipole forces

Then sulphur and oxygen slightly increase in melting because they are a larger simple molecular structure
- so has larger induced dipole-dipole forces and hence higher melting point

27
Q

what happens to the melting point in the last two elements in period 2 and 3?

A

Chlorine and Fluorine have a lower melting point because they are smaller simple molecular structure
- has smaller induced dipole-dipole forces and hence a lower melting point

Argon and Neon have lower mp than the rest due to them only existing as individual atoms
- smallest induced dipole-dipole forces and hence lowest melting point

28
Q

what charge do group 2 ions form and how?

A

outer shell is in the s2 orbital and the loss of electrons in redox reactions form 2+ ions

29
Q

what is the trend in physical properties in group 2?

A
  • the atomic radius increases down the group as more shells are added
30
Q

what is the trend in ionisation energy in group 2?

A

1st IE decreases and reactivity increases down group 2
- extra shells added meaning more shielding hence weaker electrostatic attraction between nucleus and outer electrons
- outer electrons are further from the nucleus which weakens the attractions
Therefore, these make it easier to remove the outer electron and so less energy is needed to remove the outer electron
==> increase in nuclear charge down the group but the shielding counteracts this

31
Q

what’s formed when group 2 elements react with water?

A

form bases
- reaction forms a metal hydroxide and hydrogen

32
Q

what’s the reactivity of group 2 when they react with water and why?

A

reactivity increases down the group with water
==> there is no reaction with Be
- atom gets larger and the electron is further from nucleus so easier to remove and so more reactive and more shielding

33
Q

how does magnesium react with water?

A

reacts slowly with cold water but more vigorously with steam which produces MgO instead of a hydroxide

34
Q

what do group 2 elements form when they react with oxygen?

A

form bases to form metal oxides
==> eg - 2Mg + O2 ==> 2MgO

35
Q

what are group 2 oxides?

A

white solids

36
Q

what do group 2 oxides form when they react with water?

A

form bases to form alkaline solutions (hydroxide is formed) when added to water

37
Q

how do group 2 oxides react?

A

react readily with water to make hydroxides which dissociates to form OH- ions
==> exception is magnesium oxide as it reacts very slowly and the hydroxide barely dissolves

38
Q

what is the trend in alkalinity down group 2?

A

more strongly alkaline down the group as the hydroxides become more soluble

39
Q

explain group 2 uses

A

1- Acidic soils:
group 2 metals can neutralise acids
Therefore Ca(OH)2 can be used in agriculture to neutralise acid soils

2- Antacids
Magnesium hydroxide Mg(OH)2 is used to neutralise excess stomach acid

40
Q

what is the trend in boiling point in group 7?

A

increases down the group as the London forces increase due to the increasing size and relative mass of the atoms
- the physical state goes from gas at the top of group 7 to solid at the bottom

41
Q
A

Chemistry topic 3 (2 decks)

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