C7 - Periodicity Flashcards Preview

Year 12 - Chemistry > C7 - Periodicity > Flashcards

Flashcards in C7 - Periodicity Deck (30):
1

What is periodicity?

The repeating trend in properties of the elements across each period.

2

Who identified the law of octaves?

John Newlands.
Every 8th element has similar properties, known as the law of periodicity.

3

Who formulated the periodic table?

Dimitri Mendeleev

4

How are the elements of the periodic table arranged?

They have increasing atomic number from left to right.

They are in vertical groups based on their number of electrons on their outer shell and chemical properties

They are in horizontal rows. The period number gives the number of the highest energy electron shell of that element's atom.

5

Why did Mendeleev leave spaces in his periodic table and what did he predict?

He left gaps for elements which he believed existed but hadn't been discovered yet. He predicted the existence of elements e.g. gallium.

6

What is ionisation?

A process where atoms become charged ions due to the loss or gain of electrons.

7

What is the first ionisation energy?

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

M -> M+ + e-

8

What factors affect ionisation energy?

Atomic radius - the greater the distance, the weaker the nuclear attraction.

Nuclear charge - the greater the amount of protons, the greater nuclear attraction.

Electron shielding - The more inner-shell electrons present, the more shielding due to repulsion of the electrons so attraction decreases.

9

What is the second ionisation energy?

The energy required to remove one electron from each atom in one mole of gaseous 1+ ions of an element to form one mole of gaseous 2+ ions.

10

Why would a successive ionisation graph show large increases in its energy levels?

Because an electron is being removed from an inner shell.

Removing an electron from a lower energy level requires more energy.
It also shows the filling of different shells.

11

What is the trend of first ionisation energy across a period?

First ionisation energy increases as nuclear charge increases therefore nuclear attraction increases and atomic radius decreases.

Within a period, electrons are in the same shell so there is similar shielding/shielding doesn't affect the energy.

12

What is the trend of first ionisation energy down a group?

First ionisation energy decreases as atomic radius increases and there are more inner shells so shielding increases.
This means nuclear attraction decreases so 'I' energy decreases.

13

How are metals bonded?

By metallic bonding

14

What is the structure of a metal?

Rows of (positive) atoms surrounded by a sea of delocalised electrons.
This conduct electricity due to these electrons which can move carrying charge.

Each atom donates it's outer shell electrons to a shared pool of (delocalised) electrons spread throughout the whole structure.
The (positive) cations left behind consist of the nucleus and inner electrons.

The cations are fixed meanwhile the electrons are mobile.

15

Do metals conduct better when hot or cold?

Cold as the ions are moving less so are less intrusive to the electrons carrying charge.

16

What are allotropes?

Different forms of the element

17

What's graphene?

A single layer of carbon one atom thick

18

What causes the increase in metallic bond strength from Na to Al?

Charge density - Na+ ions are large with a small charge so have a low charge density.

The number of free electrons - Na has 1 free electron per ion whereas Al has 3 so more attractions must be broken within Al.

19

What is charge density?

The ratio of an ion's charge to its size

20

What is the structure of silicon?

It has a macro molecular structure similar to a diamond.
Each Si atom has 4 strong covalent bonds.

21

What's the structure of phosphorus?

It forms a P4 tetrahedral molecule.
It has the second highest boiling point in period 3.

22

What's the structure of sulfur?

Forms crown shaped molecules. (S8)
They're the largest in period 3 and have more Van der waal forces so has the highest m/b points.

23

What's the structure of chlorine?

It forms diatomic Cl2 molecules.
It's small and has weak Van de waal forces so has low m/b points.

24

What's the structure of argon?

It's a monatomic molecule with very weak m/b points.

25

Define the term 'first ionisation energy':

It's the amount of energy required to remove one electron from each atom in one mole of gaseous atoms.

26

Explain why first ionisation energies increase from Li to Ne:

Moving across the period, nuclear charge increase so nuclear attraction increases and atomic radius decreases.
Shielding (in the same period) remains the same and doesn't affect ionisation energies.

27

Write an equation to represent the second ionisation energy of oxygen:

O+ (g) -> O 2+ (g) + e-

28

Why is the second ionisation energy of oxygen greater than the first?

There are fewer electrons for the same number of protons and the atomic radius is smaller.
Shielding decreases.

29

Solid carbon exists as diamond and graphite.

Explain why it is unnecessary to refer to carbon as either diamond or graphite when discussing their (first) ionisation energies:

Because the ionisation energy is the amount of energy required to remove one electron from each atom in one mole of gaseous atoms and, when gaseous, no covalent bonds exist.

30

What is metallic bonding?

The strong electrostatic attraction between cations and delocalised electrons.