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OCR A level Chemistry > The Periodic table trends > Flashcards

The Periodic table trends Flashcards

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1
Q

What do elements in a period show?

A

A trend in physical and chemical properties

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2
Q

Why do elements in a period show a trend in physical and chemical properties?

A

They have different electron configurations

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3
Q

What is a group?

A

A vertical column in the periodic table

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4
Q

What is a period?

A

A horizontal row in the periodic table

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5
Q

Do elements in the same group have the same or different number of electrons in their outer shell?

A

Same

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6
Q

Do elements in the same group have the same or different types of orbital in the outer shell

A

Same

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7
Q

Do elements in the same group have outer electrons in the same or different quantum shell?

A

Different

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8
Q

How is the periodic table arranged?

A

In order of increasing atomic number

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9
Q

Do elements in the same period have the same or different numbers of electrons in their outershells?

A

Different

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10
Q

Do elements in the same period have the same or different types of orbitals in their outer shells?

A

Different

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11
Q

Define periodicity

A

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

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12
Q

What do elements in the same group show?

A

Similar physical and chemical properties

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13
Q

Why do elements in the same group have similar chemical and physical properties?

A

They have a similar electronic configuration

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14
Q

Define 1st ionisation energy

A

The amount of energy required to remove one electron from EACH ATOM in one mole of gaseous atoms to form one mole of gaseous 1+ ions

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15
Q

What is the general trend in 1st ionisation energy across a period?

A

It increases

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16
Q

What is the general trend in 1st ionisation energy down a group?

A

It decreases

17
Q

What is the trend observed between group 0 1st of one period and the group 1 element of the next period?

A

There is a sharp drop

18
Q

Why is there a sharp drop in 1st IE observed between group 0 of one period and group 1 element of the next?

A
  • electrons are being removed from a new electron shell
  • this shell is further from the nucleus and experiences more shielding despite the increase in nuclear charge
  • the overall nuclear attraction is weaker
  • therefore less energy is required to remove the outer electron
19
Q

What 4 factors affect the value of ionisation energy?

A
  • number of electron shells
  • atomic radius
  • number of protons
  • nuclear attraction
20
Q

Explain the general trend in 1st ionisation energy across a period

A
  • 1st ionisation energy increases
  • nuclear charge increases across a period
  • shielding does not change across a period
  • atomic radius gets smaller across a period
  • therefore attraction between the nucleus and outer electron becomes greater
21
Q

Explain the general trend in 1st ionisation energy down a group

A
  • 1st IE decreases down a group
  • shielding increases down a group
  • atomic radius increases down the group
  • although nuclear charge increases, the other factors outweigh this
  • there is a decrease in nuclear attraction
22
Q

Some points do not obey the general trend in 1st ionisation energy- where are these?

A
  • between groups 2 and 3

- between groups 5 and 6

23
Q

Explain the difference in 1st ionisation energy between Be and B

A
  • the fall in IE marks the start of filling the p subshell
  • in Be the outer electron is in a 2s orbital
  • in B the outer electrons is in a 2p orbital
  • the 2p orbital in B is of higher energy and experiences more shielding from inner electrons as it is also shielded by the electrons in the 2s subshell
  • despite the fact the nuclear charge increasing by +1 from Be to B the difference in shielding and energy means the nuclear attraction is weaker in B
  • less energy required to remove the electron
24
Q

Explain difference in 1st ionisation energy between N and O

A
  • the fall in IE between groups 5 and 6 marks the start of electron pairing in the p-orbitals in p-subshell
  • In N the outer electron is unpaired in a 2p orbital
  • In O the outer electron is paired in a 2p orbital
  • The paired electrons in O repel slightly
  • This makes it easier to remove an electron from O than N
25
Q

What is the trend in melting points across periods 2 and 3?

A
  • Groups 5-8 have much lower melting points
  • Group 4 elements occupy the peaks of melting points
  • melting points increase from group 1 to group 2 to group 3
26
Q

Why do group 4 elements occupy the peaks of melting points?

A
  • Carbon and Silicon have giant covalent structures
  • each atom makes 4 covalent bonds
  • all of these bonds must be broken which requires a lot of energy
  • whereas Boron only makes 3 covalent bonds
27
Q

Why do elements with diatomic molecules or single atoms have the lowest melting points?

A
  • simple covalent structure
  • only intermolecular forces must be overcome, no bonds
  • the strength of London forces is related to the no. electrons that can take part in instantaneous induced dipole-induced dipole interactions
  • smaller the molecule the smaller the London forces the lower melting point
28
Q

Why do melting points increase from group 1 to group 2 to group 3?

A
  • metallic bonding present
  • metallic bonding increases as the number of electrons present that can be delocalised increases and as nuclear charge increases
  • The charge on the metal ion increases from group 1 (1+) to group 2 (2+) to group 3 (3+)
  • therefore there are more delocalised electrons in the structure
  • therefore the strength of electrostatic attraction between positive ions and delocalised electrons increases along the group
  • More energy is needed to overcome these attractions
29
Q

What is the trend in electrical conductivity across period 3?

A

Electrical conductivity increases from sodium to magnesium to aluminium

30
Q

Why does electrical conductivity increase from sodium to magnesium to aluminium?

A
  • the charge on ions increases
  • so the number of delocalised electrons within the metallic bond increases
  • delocalised electrons act as mobile charge carriers so conduct electricty
  • more delocalised electrons means greater electrical conductivity
31
Q

Why do elements with molecular structures have poor electrical conductivity?

A

-because they do not contain any mobile charge carriers

32
Q

What is the trend in atomic radius across periods?

A

-atomic radius decreases

33
Q

Why does atomic radius decrease across periods?

A
  • across a period there is little change in shielding
  • there is a regular increase in nuclear charge
  • the nuclear attraction increases
  • electrons are pulled more strongly towards the nucleus
34
Q

What the trend in atomic radius down the group?

A

It increases

35
Q

Why does atomic radius increase down the group?

A
  • there are increasing numbers of electron shells
  • therefore shielding increases
  • the attraction of electrons by the nucleus is less

OCR A level Chemistry (30 decks)

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