Periodicity

Question 0

Alkaline Earth Metals? Block?

Answer 0

Group 2:
Beryllium (Be)
Magnesium (Mg)
Calcium (Ca)
Strontium (Sr)
Barium (Ba)
Radium (Ra)

s-block.

Question 1

Alkali Metals? Block?

Answer 1

Group 1:
Lithium (Li)
Sodium (Na)
Potassium (K)
Rubidium (Rb)
Caesium (Cs)
Francium (Fr)

s-block.

Question 2

Transition Metals? Block?

Answer 2

d-block.

Question 3

Noble gases? Block?

Answer 3

Group 0:
Helium (He)
Neon (Ne)
Argon (Ar)
Krypton (Kr)
Xenon (Xe)
Radon (Rn)

p-block.

Question 4

Halogens? Block?

Answer 4

Group 7:
Fluorine (F)
Chlorine (Cl)
Bromine (Br)
Iodine (I)
Astatine (At)

p-block.

Question 5

Do d block elements (transition metals) tend to lose or gain electrons? Which electrons? s/p/d/f? Do they form positive or negative ions?

Answer 5

The d block elements tend to lose s and d electrons to form positive ions.

Question 6

What makes Group 0 elements inert?

Answer 6

Full sub-shells.

Question 7

Which groups of the Periodic Table tend to gain electrons to form negative ions?

Answer 7

Groups 5, 6 and 7.

Question 8

In the s-block (metals) do elements become more or less reactive as you go down a group?

Answer 8

More reactive.

Question 9

Are transition metals reactive?

Answer 9

No, they are unreactive.

Question 10

What is a Group in the Periodic Table?

Answer 10

A group is a vertical column of elements.

Question 11

What is a period in the Periodic Table?

Answer 11

Horizontal rows of elements.

Question 12

Elements in Groups 1, 2 and 3 lose outer electrons to form what?

Answer 12

Ionic compounds.

Question 13

Explain the meaning of the term First Ionisation Energy.

Answer 13

The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous positive ions.

Question 14

Li and Mg are ______ block elements?

Answer 14

s.

Question 15

Cl is a ________ block element?

Answer 15

p.

Question 16

Co is a ________ block element?

Answer 16

d.

Question 17

State and explain the general trend in first ionisation energies for the elements Na to P.

Answer 17

Trend: Increases
Explanation: The nuclear charge increases, shielding stays the same and there is a greater attraction for the electrons.

Question 18

State which one of these elements from Na to P deviates from this general trend and explain why it occurs.

Answer 18

Element: Aluminium
Explanation: Electron is in 3p orbital (higher orbital) and therefore increased shielding makes the electron easier to remove. Less energy is required.

Question 19

State which one of the elements from Na to P has the highest melting point and explain your answer.

Answer 19

Element: Silicon
Explanation: Silicon has a macromolecular structure therefore it has strong covalent bonds that require a lot of energy to break bonds.

Question 20

State the type of structure shown by crystals of Sulphur and Phosphorus. Explain why the melting pointing of Sulphur is HIGHER than the MP of Phosphorus.

Answer 20

Type of Structure: Simple Molecular.

Sulphur is a bigger molecule than Phosphorus. Sulphur also contains more Van der Waals to be broken.

Question 21

Explain why the melting point of Aluminium is higher than the melting point of Sodium.

Answer 21

Aluminium has a bigger charge than Sodium (Al = 3+, Na = 1+).
Al has more free delocalised electrons.
Aluminium has stronger metallic bonding.

Question 22

Identify, from Period 2 elements Li to N, the element that has the largest atomic radius.

Answer 22

Lithium.

Question 23

State the general trend in first ionisation energies for the Period 2 elements Lithium to Nitrogen.

Answer 23

Increases.

Question 24

Identify the element that deviates from this general trend, from Li to N and explain your answer.

Answer 24

Element: Boron Explanation: Electron is in 2p orbital (higher orbital) therefore the shielding increases and the electron becomes easier to remove.

Question 25

State and explain the trend in atomic radius of the Period 3 elements from Sodium to Chlorine.

Answer 25

Trend: Decreases Explanation: The nuclear charge increases, shielding stays the same and the attraction for the electrons becomes greater so atoms shrink.

Question 26

State which of the first, second or third ionisations of Al would produce an ion with the electronic configuration 1s(2), 2s(2), 2p(6), 3s(1).

Answer 26

Second.

Question 27

Explain why the value of the first ionisation energy of Sulphur is less than the value of the first ionisation of Phosphorus.

Answer 27

The pair of electrons in the 3p orbital repel each other and therefore less energy is required.

Question 28

Write an equation to show the process that occurs when the first ionisation energy of Al is measured.

Answer 28

Al(g) > Al+(g) + e-

Question 29

Identify the element in Period 2 that has the highest first ionisation energy and give its electron configuration.

Answer 29

Element: Neon | Electron Configuration: 1s(2), 2s(2), 2p(6)

Question 30

State the trend in the first ionisation energies in Group 2 from Be to Ba. Explain your answer in terms of a suitable model of atomic structure.

Answer 30

Trend: Decreases Explanation: As you go down a group there is increased shielding, a greater distance from the nucleus and so the outer electrons become easier to remove.

Question 31

Explain why the second ionisation energy of Carbon is higher than the first ionisation energy of Carbon.

Answer 31

An electron is being removed from a positive ion (closer to the nucleus).

Question 32

Deduce the element in Period 2, from Li to Ne, that has the highest second ionisation energy.

Answer 32

Lithium.