Inorganic - Periodicity

Question 1

What is periodicity?

Answer 1

The pattern in properties across a row which is repeated in each row.

Question 2

How does the atomic radius change across Period 3?

Answer 2

• Decreases across the period.
• Nuclear charge increases as the number of protons increases.
• Similar shielding of electrons along the group due to all being in the same energy level.
• Atomic size decreases as electrons feel a stronger electrostatic attraction towards the nucleus.

Question 3

How does the 1st ionisation energy change across Period 3?

Answer 3

• General increase along the period.
• Nuclear charge increases so electrons feel a stronger electrostatic attraction towards the nucleus and are harder to remove.
• Aluminium is lower than magnesium due to electrons being removed from a sub-shell (3p) further from the nucleus (increased shielding).
• Sulphur is lower than phosphorus due to pairing of electrons in the 3p orbital causing a repulsive force making an electron easier to remove.

Question 4

Why do successive ionisation energies increase?

Answer 4

Once you have removed the first electron, you are left with a positive ion. Trying to remove a negative electron from a positive ion is going to be more difficult than removing it from an atom, because of the force of attraction between the negatively charged electron and the positively charged ion.

Question 5

How does the electronegativity change across Period 3?

Answer 5

• Nuclear charge increases so electrons feel a stronger electrostatic attraction.
• Shielding is similar across period 3 and not significant enough to overcome increasing nuclear charge.
• Argon does not have an electronegativity value as it does not form bonds.

Question 6

How does the electrical conductivity change across Period 3?

Answer 6

• Sodium, magnesium and aluminium are the only substances able to conduct electricity as they can form structures with free electrons.
• Aluminium has the greatest number of delocalised electrons.
• So more are free to carry charge.

Question 7

How does the melting and boiling point change across Period 3?

Answer 7

• Sodium, magnesium, aluminium have strong metallic bonding with increasing electrostatic attraction between positive ions and delocalised electrons.
• Silicon is a giant covalent structure with many strong covalent bonds.
• Phosphorus, sulphur and chlorine are simple covalent structures with weak van der Waals forces.
• Argon is monatomic with weak van der Waals forces.

Question 8

Describe the structure of phosphorus (P4).

Answer 8

In red phosphorus, each molecule exists in a tetrahedral structure. The atoms are joined by covalent bonds within the molecule. Melting point drops dramatically as intermolecular attractions are now due to weak van der Waals forces.

Question 9

Describe the structure of sulphur (S8).

Answer 9

Molecule can exist in an eight membered ring structure. The atoms are joined by covalent bonds within the molecule. Melting point rises slightly as the molecule is bigger/heavier so has slightly stronger van der Waal’s forces.

Question 10

Describe the structure of chlorine.

Answer 10

Exists as a linear diatomic molecule. Melting point falls slightly as the molecule is smaller so has slightly lower van der Waal’s forces.

Question 11

Describe the structure of argon.

Answer 11

Argon is a noble gas that exists as individual atoms with a full outer shell of electrons. This makes the atom very stable and as it is also lighter than Cl2, the van der Waal’s forces between them very weak. As a result, the melting point of Argon is very low and it exists as a gas at room temperature.

Question 12

Where are the metalloids found on the periodic table? What are they?

Answer 12

The staircase line divides metals (on its left) from non-metals (on its right). Elements that touch this line have a combination of metallic and non-metallic properties. They are called metalloids or semi-metals.

Question 13

Why is silicon an example of a metalloid?

Answer 13

Silicon is a non-metal but looks quite shiny and conducts electricity, although not as well as a metal.

Question 14

Why are elements found in the s/p/d block?

Answer 14

If their highest energy electrons are in that orbital.

Question 15

What is a group?

Answer 15

A group is a vertical column of elements. The elements in the same group form a chemical ‘family’ - they have similar properties. Elements in the same group have the same number of electrons in the outer main level.

Question 16

Describe the general trend in reactivity in the periodic table.

Answer 16

In the s-block, elements (metals) get more reactive going down a group. To the right (non-metals), elements tend to get more reactive going up a group.

Transition elements are a block of rather unreactive metals. This is where most of the useful metals are found.

Lanthanides are metals which are not often encountered. They all tend to form +3 ions in their compounds and have broadly similar reactivity.

Actinides are radioactive metals. Only thorium and uranium occur naturally in the Earth’s crust in anything more than trace quantities.

Question 17

What is a period?

Answer 17

Horizontal rows of elements in the Periodic Table are called periods. There are trends in physical properties and chemical behaviour as you go across a period.

Question 18

Explain the electron arrangement of sodium, magnesium and aluminium.

Answer 18

These elements are metals. They have giant structures. They lose their outer electrons to form ionic compounds.

Question 19

Explain the electron arrangement of silicon.

Answer 19

Silicon has four electrons in its outer shell with which it forms four covalent bonds. The element has some metallic properties and is classed as a semi-metal.

Question 20

Explain the electron arrangement of phosphorus, sulphur and chlorine.

Answer 20

Phosphorus, sulphur and chorine are non-metals. They either accept electrons to form ionic compounds, or share their outer electrons to form covalent compounds.

Question 21

Explain the electron arrangement of argon.

Answer 21

Argon is a noble gas - it has a full outer shell and is unreactive.

Question 22

Which element in period 3 has the highest and lowest melting point?

Answer 22

highest - silicon

lowest - argon

Question 23

Which element in period 3 has the highest and lowest boiling point?

Answer 23

highest - aluminium

lowest - argon

Question 24

Why does the first ionisation energy increase across a period?

Answer 24

As you go across a period from left to right, the number of protons in the nucleus increases but the electrons enter the same main level. The increased charge on the nucleus means that it gets increasingly difficult to remove an electron.

Question 25

Why does the first ionisation energy decrease going down a group?

Answer 25

The number of filled inner levels increases down the group. This results in an increase in shielding. Also, the electron to be removed is at an increasing distances from the nucleus and is therefore held less strongly. Thus the outer electrons get easier to remove going down a group because they are further away from the nucleus.

Question 26

Why is there a drop in ionisation energy from one period to the next?

Answer 26

Moving from neon in Period 0 (far right) with electron arrangement 2.8 to sodium 2.8.1 in Period 1 (far left), there is a sharp drop in the first ionisation energy. This is because at sodium, a new main level starts and so there is an increase in atomic radius, the outer electron is further from the nucleus, less strongly attracted and easier to remove.