The Periodic table

Question 1

What did Johann Dobereiener do?

Answer 1

1817 - attempted to group similar elements - Dobereiners triads.
Saw Chlorine, Bromine and iodine had similiar characteristics.
Lithium, sodium and potassium.
Saw that bromine atomic weight fell between chlorine and iodine.

Question 2

What did John Newlands do?

Answer 2

1863- noticed if he arranged elements in order of mass, elements with similar chemical and physical properties appeared at regular intervals.
The law of octaves, and listed known elements in rows of 7 so similar were in columns.

Question 3

What is the Periodic Law?

Answer 3

The law of octaves developed into this.
States that if you arrange elements in order of increasing atomic number then their chemical and physical properties will repeat in a systematic way that can be predicted.
But the pattern broke down on the third row.

Question 4

What did Mendeleev do?

Answer 4

1869 - built on Newlands, by leaving gaps in the table where the next element didn’t fit.
So elements with similar chemical properties were in the same groups.
He predicted the properties of missing elements, and was correct.

Question 5

What are periods?

Answer 5

All the elements in a period (row) have the same number of electron shells.
This results in a repeating pattern of physical and chemical properties across a period (periodicity).

Question 6

What are groups?

Answer 6

All elements in a group have the same number of electrons in their outer shell.
So they have similar physical and chemical properties.
Group number tells the number of electrons in the outer shell.

Question 7

What are the blocks of elements?

Answer 7

s-block elements have an outer shell electron configuration of s1 or s2.
p-block have configuration of s2p1 to s2p6.
d-block have configurations in which d-sub shells are being filled.

Question 8

What is first ionisation energy?

Answer 8

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

Question 9

How should ionisation energies be written?

Answer 9

Use the gas state symbol (g).
Always refer to one mole of atoms.
The lower the ionisation energy, the easier to form an ion.

Question 10

What is high ionisation energy?

Answer 10

There is strong electrostatic attraction between the nucleus and the electron, so more energy is needed to overcome the attraction and remove the electron.

Question 11

How does nuclear charge affect ionisation energy?

Answer 11

The more protons in the nucleus, the more positively charged the nucleus is, and the stronger the attraction for the electrons.

Question 12

How does atomic radius affect ionisation energy?

Answer 12

Attraction falls off very rapidly with distance.
An electron close to the nucleus will be much more strongly attracted than ones further away.

Question 13

How does shielding affect ionisation energy?

Answer 13

As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge.

Question 14

What is the trend in ionisation down a group?

Answer 14

Ionisation energies generally fall.
Elements further down have extra electron shells, so atomic radius is larger, and inner shells shield the outer electrons from the attraction.
The increasing nuclear charge is overridden by this.

Question 15

What is periodicity of ionisation energy?

Answer 15

Ionisation energies generally increase across a period.
Nuclear charge increases, and all the extra electrons are at the same energy level, so similar shielding and atomic radius.

Question 16

What are the differences in ionisation energies in groups?

Answer 16

The drop in energy between Groups 2 and 4 is due to the extra electrons in group 3 elements being in a p-orbital. Increases shielding and energy level.
The drop between 5 and 6 is due to electron repulsion.

Question 17

What is electron repulsion?

Answer 17

Electron pairs around a central atom tend to orient themselves as far apart as possible.
Electrons removed from a shared orbital means electrons are easier to remove than from singly occupied orbitals.

Question 18

What are successive ionisation energies?

Answer 18

Each time an electron is removed.
Second ionisation of oxygen: O+ (g) –> O^2+(g) +e-

Question 19

What is the trend of ionisation energies in shells?

Answer 19

Within each shell, successive ionisation energies increase.
More electrons are being removed from an increasingly positive ion.
There’s less repulsion, so more energy needed.
Big jumps happen when a new shell is broken into - an electron is being removed from a shell closer to the nucleus.

Question 20

How do you use graphs to find groups?

Answer 20

Count how many electrons are removed before the first big jump to find the group number.

Question 21

What are giant covalent lattices?

Answer 21

Large networks of covalently bonded atoms.
Carbon and silicon both form 4 strong covalent bonds.

Question 22

What is the structure of graphite?

Answer 22

The carbon atoms are sheets of flat hexagons with 3 covalent bonds.
The fourth electron is delocalised.
The sheets are bonded by weak induced dipole dipole forces.

Question 23

What are the properties of graphite?

Answer 23

The weak forces are easily broken so it’s slippery - dry lubricant and pencils.
Delocalised electrons move along sheets - carry electron flow.
Layers are far apart so has low density, used to make strong, light sports equipment.
The strong covalent bonds give it a high melting point.
It is insoluble.

Question 24

What is the structure of diamond?

Answer 24

Each carbon atom is covalently bonded to 4 other carbons.
Arranged in a tetrahedral shape.

Question 25

What are the properties of diamond?

Answer 25

Very high melting point.
Extremely hard - diamond-tipped drills and saws.
Vibrations travel easily so is a good thermal conductor.
Can’t conduct electricity.
Insoluble.

Question 26

What is silicon?

Answer 26

Same group as carbon.
Each one forms 4 strong covalent bonds.
It makes a crystal lattice structure and has similar properties to diamond.

Question 27

What is the structure of graphene?

Answer 27

1 layer of graphite.
1 atom thick, a 2-dimensional compound.
Giant covalent lattice.

Question 28

What are the properties of graphene?

Answer 28

Delocalised electrons move quickly above and below the sheet - the best electrical conductor.
Delocalised electrons strengthen covalent bonds, so is extremely strong.
A single layer is transparent and incredibly light.

Question 29

What is graphene used for?

Answer 29

Its’ high strength, low mass, and good electrical conductivity, means it can be used in high-speed electronics and aircraft technology.
Its’ flexibility and transparency can be used as a touch screen.

Question 30

What is metallic bonding?

Answer 30

The metal cations (+) are electrostatically attracted to the delocalised electrons.
They form a lattice of closely packed cations in a sea of delocalised electrons.

Question 31

What is melting point of metallic bonding?

Answer 31

The more delocalised electrons per atom, the stronger bonding, so the higher melting point.
A smaller ionic radius holds delocalised electrons closer to the nuclei.

Question 32

What is the shaping of metallic bonding?

Answer 32

No bonds holding specific ions together, so metal ions can slide over each other when pulled.
Metals are malleable and ductile (drawn into a wire).

Question 33

What is the conductivity of metallic bonding?

Answer 33

The delocalised electrons pass kinetic energy to each other, so are good thermal conductors.
Can carry a charge, so good electrical conductors.

Question 34

What is the solubility of metallic bonding?

Answer 34

Metals are insoluble, except in liquid metals, because of the strength of metallic bonds.

Question 35

How do melting and boiling points change across periods?

Answer 35

Melting points of period 2 and 3 elements increase from first to fourth elements, then decrease from fourth to eighth.
Across the period, the type of bonding and structure changes from giant metallic to giant covalent to simple molecular lattice

Question 36

How does melting point of metals change across the period?

Answer 36

Boiling points increase because metal-metal bonds get stronger.
Metal ions have a greater charge, an increasing number of delocalised electrons, and decreased ionic radius.
Causes higher charge density, attracts ions together stronger.

Question 37

How does melting points of simple molecular structures change?

Answer 37

(N2, O2, F2, S8, Cl2) They form simple molecular /covalent lattices.
Covalent bonds between atoms are very strong, but melting point depends on strength of induced dipole dipole forces.
More atoms in a molecule means stronger induced forces, so higher mp.
Noble gases have low mp as they’re individual atoms, so weak induced forces.

Question 38

What is electrical conductivity of simple molecular lattices?

Answer 38

Do not conduct electricty because charge carriers/electrons are not mobile.