Chapter 7 Periodicity

Question 1

How did Mendeleev organise his periodic table?

Answer 1

arranged using atomic mass and grouped elements that had similar properties.
–> If grouped properties did not fit, Mendeleev swapped elements around and left gaps, assuming the atomic mass measurements were incorrect

Question 2

How is the modern Periodic table organised?

Answer 2

114 elements organised into 7 horizontal periods and 18 vertical groups
–> position is linked to physical and chemical properties
–> arranged in increasing atomic number

Question 3

Groups and periods

Answer 3

vertical columns: groups which all have atoms with the same number of outer-shell electrons and similar properties

horizontal rows: periods
–> no on period gives number of highest energy electron shell in an element’s atom

Question 4

Periodicity & properties

Answer 4

repeating trends in properties of elements
–> electron configuration
–> ionisation energy
–> structure
–> melting point

Question 5

Trends across a period

Answer 5

Across period 2, 2s sub-shell fills with two electrons, then 2p sub-shell with 6 electrons e.g. Be is (He)2s^2 and Ne is (He)2s^2,2p^6

Across period 3 same pattern of filling for 3s and 3p sub-shells

Across period 4, although 3d sub-shell is involved, highest shell number=4 so from n=4, only 4s and 4p sub-shells are occupied

Question 6

Trends down a group

Answer 6

Elements have the same number of electrons in each sub-shell (can be divided into 4 distinct blocks)
–> Group 1&2 are s-block
–> Group 3-12 are d block
–>Group 13-18 are p-block

Question 7

Names and numbers of groups

Answer 7

1- Alkali metals
2- alkaline earth metals
3-12 Transition elements
15 pnictogens (N,P,As,Sb,Bi)
16 chalcogens (O,S,Se,Te,Po)
17 Halogens
18 Noble gases

Question 8

Ionisation energy

Answer 8

How easily an atom loses electrons to form positive ions

Question 9

First ionisation energy

Answer 9

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

Na(g) –> Na+(g) + e- //. X(g) –> X+(g) + e-

Question 10

Successive ionisation of Chlorine

Answer 10

Cl(g) –> Cl+(g) +e- It is not diatomic because ionisation energy is from ONE atom at a time

IE2: Cl+(g) –> Cl2+(g) + e- (only one electron is removed per ionisation energy)

Question 11

Factors that affect ionisation energy

Answer 11

Atomic radius, nuclear charge, electron shielding

Question 12

Atomic radius IE

Answer 12

The greater the distance between the nucleus and outer electrons, the less nuclear attraction
–> force of attraction falls of sharply with increasing distance so atomic radius has a large effect

Question 13

Nuclear Charge IE

Answer 13

More protons in nucleus= greater attraction between nucleus and outer electrons

Question 14

Electron Shielding IE

Answer 14

Inner-shell electrons repel outer-shell electrons and reduces the attraction between the nucleus and outer electrons

Question 15

Why is the second ionisation energy greater than the first IE?

Answer 15

Once an electron is lost, the single electron is pulled in closer as there are now more protons than electrons.
–> causes a stronger attraction on the remaining electron

Question 16

Second ionisation energy definition

Answer 16

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

Question 17

What does a large difference in ionisation energy mean?

Answer 17

This means a different shell is now losing electrons
–> much closer to the nucleus so stronger attraction
–> more energy is required to lose the electrons)

Question 18

Trend in first ionisation energies

Answer 18

As you go across each period, there is a general increase in first IE
–> higher nuclear charge= more protons in nucleus= more attraction between nucleus and outermost electron

A sharp decrease between end of one period and start of another ( He->Li/ Ne->Na/ Ar->K)

Question 19

As you go down a group, why does first IE decrease?

Answer 19

atomic radius increases
more inner shells in total so shielding increases (More repulsion)
nuclear attraction on outer electrons decreases
first ionisation energy decreases

Question 20

Trends in First IE across a period

Answer 20

Across a period, the nuclear charge increase as there are more protons
–> same shell means similar shielding levels
–> nuclear attraction increases and atomic radius decreases
–> first ionisation energy increases

Question 21

Why does Beryllium have a higher ionisation energy than Boron

Answer 21

Beryllium has no 2p subshell but Boron has one electron in the 2p sub-shell.
2p sub-shell in boron has a higher energy than 2s sub-shell in beryllium.
–> Therefore, in boron, the 2p electron is easier to remove than one of the 2s electrons in beryllium (lower ionisation IE)

Question 22

Why does oxygen have a lower ionisation energy than nitrogen

Answer 22

Due to electron pairing in p-orbitals of 2p sub-shell
–> in N & O, the highest energy electron are in a 2p sub-shell
–> in oxygen, paired electrons in one of the 2p orbitals repel one another, making it easier to remove an electron from an oxygen atom than a nitrogen atom
–> ionisation energy of oxygen is less than nitrogen

Question 23

Nitrogen and Oxygen electron configuration

Answer 23

3 2p electrons: 2px1 2py1 2pz1
–> one electron in each 2p orbital
–> spins are at right angles - equal repulsion as far apart as possible

4 2p electrons: 2px2 2py1 2pz1
–> two electrons in one 2p orbital
–> 2p electrons start to pair
–> paired electrons repel

Question 24

Metallic bonding and structure

Answer 24

in solid metal structure, each atom has donated its negative outer-shell electrons to a shared sea of electrons. (delocalised)
–> positive ions ( cations) left behind consist of nucleus and inner electron shells of metal atoms

Metabolic atoms: strong electrostatic forces of attraction between cations and delocalised electrons
–> fixed cation with mobile delocalised electrons
(giant metallic lattice)

Question 25

Properties of metals

Answer 25

strong metallic bonds
high electrical conductivity
high melting and boiling points

Question 26

Electrical conductivity

Answer 26

metals conduct electricity in solid and liquid states
–> voltage applied across metal= delocalised electrons can move through the structure, carrying charge
–> ionic compounds cannot carry charge when solid

Question 27

Melting and boiling points

Answer 27

depends on strength of metallic bonds in a giant metallic lattice
–> most require high temps to provide large amount of energy needed to overcome strong electrostatic attraction between cations and electrons
–> strong attraction results in most metals having high melting and boiling points

Question 28

Solubility

Answer 28

metals do not dissolve (can interact in a polar solvent)

Question 29

Giant covalent structures

Answer 29

strong network of covalent bonds
–> carbon and silicon both have 4 electrons
–> can form 4 covalent bonds to other carbon or silicon atoms (results in tetrahedral shape)
–> 109.5 bond angle

Question 30

Properties of giant covalent structures

Answer 30

melting and boiling points: high due to strong covalent bonds (need lots of energy to break)

solubility: insoluble as covalent bonds are way too strong to be broken by interaction with solvents

electrical conductivity: non conductors, except graphene and graphite
- in diamond all 4 electrons are used

Question 31

periodic trend in melting points

Answer 31

• Melting point increase from group 1-14 (giant structures)
• sharp decrease in melting point from group 14-15 (simple molecules)
• melting points are comparatively low from group 15-18

Question 32

Bond strength

Answer 32

Giant covalent> Giant metallic> simple molecular

simple molecules: weak London forces between molecules