Chapter 7 The Periodic Table

Question 1

How did Mendeleev initially arrange the elements?

Answer 1

In order of atomic mass.
(Also lined up the elements in groups with similar properties.)

If group properties didn’t fit, Mendeleev swapped elements around and left gaps, assuming that the atomic mass measurements were incorrect and that some elements were yet to be discovered.
He predicted properties of the missing elements from group trends.

Once protons were discovered in the early 1900s, the real reason for the order of Mendeleev’s table was revealed.

Question 2

Describe the periodic table NOW?

Answer 2

114 elements
- 7 horizontal periods
- 18 vertical groups

Question 3

What are the positions of the elements in the periodic table linked to?

Answer 3

Linked to the physical and chemical properties of the elements.
- Makes the periodic table essential for predicting the properties of elements and their compounds.

Question 4

Atomic number?

Answer 4

From LEFT to RIGHT, the elements are arranged in order of INCREASING atomic number.

Each successive element has atoms with 1 EXTRA proton.

Question 5

What are GROUPS in the periodic table?

Answer 5

= Vertical columns.

Each element in a group has atoms with the same number of OUTER-SHELL ELECTRONS (and therefore similar properties).

Question 6

What are PERIODS/PERIODICITY in the periodic table?

Answer 6

Periods = horizontal rows

The number of the period gives the number of the HIGHEST ENERGY ELECTRON shell in an element’s atoms.

Periodicity = repeating trends in properties of the elements…

(1) electron configuration
(2) ionisation energy
(3) structure
(4) melting points

Question 7

(1) Periodic trend in electron configuration?

Answer 7

Each period starts with an election in a NEW highest energy shell.

• Across Period 2, the 2s sub-shell fills with 2 electrons, followed by the 2p sub-shell with 6 electrons.
• Across Period 3, the same pattern of filling is repeated for the 3s and 3p sub-shells.
• Across Period 4, although the 3d sub-shell is involved, the highest shell number is n = 4. From n = 4 shell, only the 4s and 4p sub-shells are occupied.

Question 8

(1) Electron configuration trend DOWN a group?

Answer 8

Elements in each group also have atoms with the same number of electrons in each sub-shell.
= similar chemistry

Question 9

(1) Blocks??

Answer 9

The elements in the periodic table can be divided into blocks corresponding to their highest energy sub-shell.
= s
= p
= d
= f

Question 10

(2) What is (first) ionisation energy?

Answer 10

Measures how easily an atom loses electrons to form positive ions.

FIRST ionisation energy is the 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.

Question 11

(2) Factors affecting ionisation energy?

Answer 11

Electrons are held in their shells by attraction from the nucleus.
The first electron lost will be in the highest energy level and will experience the least attraction from the nucleus.

• ATOMIC RADIUS : the greater the distance between the nucleus and the outer electron , the less the nuclear attraction.
• NUCLEAR CHARGE : the more protons there are in the nucleus of an atom, the greater the attraction between the nucleus and the outer electrons.
• ELECTRON SHIELDING : electrons are negatively charged and so inner-shell electrons repel outer-shell electrons = shielding EFFECT - reduces the nuclear attraction.

Less the nuclear attraction = less the ionisation energy

Question 12

What are successive ionisation energies?

Answer 12

After the 1st electron is lost, the single electron is pulled closer to the nucleus. The nuclear attraction on the remaining electron increases and more ionisation energy will be needed to remove this second electron.

Definition (textbook) : 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 13

What is the link between successive ionisation energies and shells?

Answer 13

Successive ionisation energies provide evidence for the different electron energy levels in an atom.

A large increase between 2 ionisation energies suggests that the electron must be removed from a different shell, closer to the nucleus and with less shielding (higher nuclear attraction = higher ionisation energy).

Question 14

What predictions can be made from successive ionisation energies?

Answer 14

• number of electrons in outer shell
• group of an element in the periodic table
• identify of an element

Question 15

What is trend in first ionisation energy down a group?

Answer 15

First ionisation energies DECREASE down a group.

Although the nuclear charge increases, its effect is outweighed by the increased radius (and the increased shielding).

EXAM ANSWER…
1) atom is radius increases
2) more inner shells so shielding increases
3) nuclear attraction on outer electron decreases
4) first ionisation energy decreases

Question 16

What is the trend in ionisation energy ACROSS a PERIOD?

Answer 16

First ionisation energies INCREASE going across (L —> R) periods.

EXAM ANSWER…
1) nuclear charge increases
2) same shell : similar shielding (essentially not much effect)
3) nuclear attraction increases
4) atomic radius decreases
5) first ionisation energy increases

Question 17

Describe and explain metallic bonding?

Answer 17

= the strong electrostatic attraction between CATions and delocalised electrons

Bonding in METALS…

Solid metal, each atom has donated negative outer-shell electrons = delocalised electrons spread thru the structure.

Positive ions (CATions) left behind consist of the nucleus and the inner electron shells of the metal atoms.

The CATions are fixed in position, maintaining structure and shape of the metal.

The delocalised electrons ate mobile (free to move) - allowing conduction.

Question 18

What is a giant metallic lattice?

Answer 18

In a metal structure, billions of metal atoms are held together by metallic bonding in a giant metallic lattice.

Question 19

What are the typical properties of metals?

Answer 19

STRONG metallic bonds (CATions and delocalised electrons)

HIGH electrical conductivity [delocalised electrons can move and carry a charge (when a voltage is applied across a metal)]

HIGH melting and boiling points (large amount of energy required to overcome the strong electrostatic attraction between CATions and delocalised electrons)

Question 20

Solubility of metals??

Answer 20

Metals do NOT dissolve.

• Any interaction would lead to a reaction rather than dissolving.

Question 21

SIMPLE covalent structures??

Answer 21

Non-metallic elements exist as simple covalent bonded molecules.

• Solid state = simple molecular lattice structure held by weak intermolecular forces

LOW melting and boiling points

Question 22

GIANT covalent structures??

Answer 22

• BORON
• CARBON - group 4 = tetrahedral structure (109.5)
• SILICON - group 4 = tetrahedral structure (109.5)

All have different lattice structures. Held together by a network of strong covalent bonds = GIANT COVALENT LATTICE

Question 23

What are the typical properties of substances with a giant covalent lattice structure? (Melting, boiling, solubility, conductivity)

Answer 23

HIGH melting and boiling points as covalent bonds are strong and require a large quantity of energy to overcome.

INSOLUBLE in almost all solvents.
- Covalent bonds holding together the atoms in the lattice are far too strong to be broken by interaction with solvents.

NON-CONDUCTORS of electricity except for GRAPHENE and GRAPHITE (carbon).
- Carbon - all 4 outer-shell electrons are involved in bonding (NOT AVAILABLE).

Question 24

Describe the structure of GRAPHENE?

Answer 24

• Single layer of graphite
• Hexagonally arranged carbon atoms linked by strong covalent bonds
• Only 3 electrons used in covalent bonding.
• Remaining electron is released into a pool of delocalised electrons shared by all atoms in the structure = GOOD electrical conductor.

Question 25

Describe the structure of GRAPHITE?

Answer 25

• Parallel layers of hexagonally arranged carbon atoms.
• Layers bonded together by weak LONDON FORCES.
• The bonding in the hexagonal layers only uses 3 of carbon’s 4 outer-shell electrons. The spare electron is delocalised between the layers = GOOD electrical conductor.

Question 26

What is the periodic trend in MELTING POINTS?

Answer 26

ACROSS PERIOD 2, 3…

• Melting point increases (group 1 - 4)
• Melting point decreases (group 4 - 5)
  Sharp decrease = marks change from GIANT to SIMPLE covalent structures.
• Melting points relatively low (group 15 - 18)