Chapter 7 - Periodicity

Question 1

Who created the modern periodic table?

Answer 1

Dmitri Mendeleev

Question 2

How were the elements ordered by Mendeleev?

Answer 2

Atomic mass

Question 3

What did Mendeleev do that no one before him did?

Answer 3

• Put elements into groups with similar properties. Swapping some elements around so they fit trends.
• Left gaps for undiscovered elements so the patterns would fit. He therefore predicted properties of these elements.

Question 4

What are the elements now ordered in?

Answer 4

Atomic number

Question 5

What are the vertical columns that elements are arranged in?

What are the horizontal columns that elements are arranged in?

Answer 5

Groups

Periods

Question 6

What is periodicity?

Answer 6

Repeating trend across a period

Question 7

How does electron configuration change across a period?

Answer 7

The S and P sub-shells are filled the same way.
S subshell fills with 2 electrons, then p subshell to to 6. - up n=4 shell
Periodic pattern

Question 8

How does electron configuration change down the group?

Answer 8

• Same number of electrons in outer shell.

- Same number of electrons in each sub-shell.

Question 9

Give the definition of the first ionisation energy?

Answer 9

Energy required to remove one electron from each atom in one mole of gaseous atoms of an element toform one mole of 1+ ions.

Question 10

Give the first ionisation half equation of Sodium?

Answer 10

Na(g) -> Na+(g) +e-

Question 11

Give 3 factors that affect ionisation energy

Answer 11

• Atomic radius - Greater the distance between nucleus and the outer electrons the less the nuclear attraction. Inversely proportional to nuclear attraction.
• Nuclear charge - More protons = greater attraction between nucleus and outer electrons.
• Electron shielding - More electrons shells = more shielding effect, reducing attraction between nucleus and outer electrons.

Question 12

How many ionisation energies does Magnesium have?

Answer 12

12 - as magnesium has 12 electrons to lose

Question 13

Define second ionisation energy

Answer 13

Energy required to remove one electron from each ion in one mole of 1+ ions of an element to form one mole of gaseous 2+ ions.

Question 14

What does large jump in ionisation energy between ionisation energy number 7 and 8 signify in analysis of successive ionisation energy graph?

Answer 14

8th electron must be removed from a different from a different shell, closer to the nucleus with less shielding - and therefore takes more energy to remove.

Question 15

How does ionisation energy change down the group? And why?

Answer 15

Decreases down the group. Atomic radius increases, due to more electron shells meaning more shielding - so decreases nuclear attraction on outer electrons, even though nuclear charge increases - this is outweighed by shielding. And so less energy is required to remove an electron from its outer shell.

Question 16

How does the first ionisation energy change across a period? And why?

Answer 16

Increases across a group. This is because the atomic radius decreases due to same or similar shielding in all elements but an increase in nuclear charge across the group and so increased nuclear attraction. And so increased first ionisation energy.

Question 17

When is the rule of increasing ionisation energies across a period wrong?

Answer 17

• When a new sub-shell starts e.g. when first adding electrons to P sub-shell. e.g. Be to B
• When starting to pair electrons as paired electrons repel and less energy required to remove e.g. N to O

Question 18

What is a metallic bond?

Answer 18

Strong electrostatic attraction between cations (positive ions and delocalised electrons.

Question 19

Why do metals conduct electricity?

Answer 19

Delocalised electrons (given out by cations) in sea of delocalised electrons carry charge through the structure.

Question 20

Why do metals have high boiling and melting points?

Answer 20

High temperatures/ large amounts of energy required to overcome strong metallic bonds holding the atoms together in the giant metallic lattice.

Question 21

List the properties of a giant covalent molecule

Answer 21

• High melting + boiling point
• Insoluble
• Non-conductors of electricity (except hexagonal layers of carbon e.g. graphene and graphite)

Question 22

How does the melting point change across a period?

Answer 22

Melting point increases from group 1 to group 4 (giant structures). Then a sharp decrease between group 4 and 5 and then the melting points are low from group 5 to 8 as these are simple molecules.

Question 23

Why does melting point increase from group 1 to group 4?

Answer 23

Ions have greater charges (greater charge density) and also more electrons - so metallic bond is stronger due to increased attraction between ions and electrons.