1. Atomic structure ( ORBITALS + IONISATION)

Question 1

Define an orbital

Answer 1

a region of space in an atom containing up to two electrons with opposite spins

Question 2

Give the 4 orbitals in an atom

Answer 2

s p d f

Question 3

What is the shape of the s orbital compared to the p orbital

Answer 3

s orbital is spherical and p orbital is dumbell shaped

Question 4

What are the names of the 3 P orbitals

Answer 4

Px, Py and Pz

Question 5

How many electrons can the p orbital hold

Answer 5

6

Question 6

Give the maximum number of electrons in each shell

Answer 6

shell 1 - 2
shell 2 - 8
shell 3 - 18
shell 4 - 32

Question 7

Give the order in which electrons fill subshells

Answer 7

1s–> 2s –> 2p –> 3s –>3p –> 4s –> 3d –> 4p –> 5s –> 4d –> 5p

Question 8

Give the type of orbital found at each energy level

Answer 8

1 - s
2 - sppp
3- spppddddd
4- spppdddddfffffff

Question 9

Why is the 4s orbital filled before the 3d orbital

Answer 9

Because 4s is slightly lower in energy

Question 10

Describe 2 ways of writing electron structure

Answer 10

Spin diagrams - electrons in boxes with arrows
Short hand - uses the noble gas that comes before it.
eg: oxygen is 1s2 2s2 2p4 but can be written as [He] 2s2 sp4. (must use square brackets)
–> This is because helium has 2 electrons in its outer shell so it can substitute the 1s2

Question 11

When forming ions, is the 4s orbital removed before the 3d orbital

Answer 11

Yes ( apart from copper and chromium )

Question 12

What is the exception to the rule of filling 4s orbital before 3d orbital

Answer 12

Chromium and copper
Its more stable this way

Question 13

How many electrons can the d orbital hold

Answer 13

10
(so it could be written as eg: 3d^10)

Question 14

How many electrons can the f orbital hold

Answer 14

14

Question 15

In what order do electrons fill subshells

Answer 15

First singularly then pairing up

Question 16

What is meant by s block elements

Answer 16

The elements in group 1 and 2 have highest energy electron in an s orbital

Question 17

What is meant by d block elements

Answer 17

Transition metals

Question 18

What is meant by p block elements

Answer 18

Elements from group 3 to group 7

Question 19

Define isoelectronic

Answer 19

When either atoms, isotopes or ions ( referred to as species) have the same number of electrons

Question 20

Define first ionisaton energy

Answer 20

First ionisation energy is the energy required to remove one mole of electrons from one mole of atoms in the gaseous state
[ M(g) → M+(g) + e- ]

Question 21

Define second ionisation energy

Answer 21

is the energy required to remove one mole of electrons from one mole of 1+ ions in the gaseous state.
[ M+(g) → M2+(g) + e- ]

Question 22

Give 3 factors that affect ionisation energy

Answer 22

size (distance of the outer shell electrons from the nucleus) , shielding, nuclear charge

Question 23

What happens to first ionisation energies going down a group

Answer 23

It decreases

Question 24

Explain why first ionisation energies decreases down a group

Answer 24

Atomic radius (size) increases - increased number of shells so distance from outer shell electrons to nucleus increases
Shell number increases - more shielding between outer shell electron and nucleus due to the inner electrons, so force of attraction decreases
Nuclear charge increases - more protons in nucleus ( but this increase is outweighed by the other factors)

Question 25

What happens to ionisation energy going across a period

Answer 25

They generally increase

Question 26

Explain why ionisation energy increases across a period

Answer 26

Size of the atom decreases - due to the greater force of electrostatic attraction between the positive nucleus and outer shell electron
Nuclear charge increases - due to higher number of protons in the nucleus
Shielding stays the same - number of shells stays the same as electrons are removed from the same shell so experiences simillar sheilding

Question 27

Explain why sucessive ionisation energies increase in size

Answer 27

The electrons are now closer to the nucleus and there are more protons than electrons. The electron must also overcome the increasing attraction of the positive ion.

Question 28

What do the big jumps on a ionisation energy graph show + what do they provide evidence for

Answer 28

shows when an electron is being removed from a shell
Provides evidence for shells

Question 29

Describe how to work out the group of an element and its structure from an ionisation energy graph

Answer 29

To work out the group, look for how many electrons are removed before the first big jump. Look from left to right

To work out structure, look from right to left and count how many points there are before a big jump

Question 30

Define periodicity

Answer 30

The regularly repeating patern of atomic, physical and chemichal properties with increasing atomic number

Question 31

How do chemical and physical properties of isotopes differ

Answer 31

the chemical properties are the same (because the number of protons doesn’t change and electron configuration remains the same) but the physical properties (such as mass/density) are different because of different amounts of neutrons

Question 32

Why is there a drop in ionisation energy between group 2 and 3

Answer 32

An element in group 3 has its outer shell electron in a p orbital, not an s orbital.
P orbital has a slightly higher energy than s orbital
P orbital is therefore found further away from the nucleus.
These 2 factors override the increased nuclear charge so ionisation energy decreases

Question 33

Why is there a drop in ionisation energy between group 5 and 6

Answer 33

Group 5 elements have their outer shell electrons in singularly half filled p orbitals which are more stable than those with paired filled subshells.
Elements in group 6 have their outer shell electron paired with another electron in the p orbital. The repulsion between the 2 electrons makes it easier to remove so ionisation energy decreases

Note: Ionisation energy for group 6 elements is less than group 5 but greater than group 3

Question 34

When going across the period does melting and boiling point increase or decrease for metals

Answer 34

Increases as metallic bonds get stronger due to decreased atomic radius - there is a stronger electrostatic attraction between positively charged ions and a sea of
delocalised electrons

Question 35

Describe the melting and boiling point for giant covalent lattice structures ( Carbon, Silicon)

Answer 35

Very high due to many strong covalent bonds between atoms so high energy needed to break covalent bonds

Question 36

What is the trend of melting/boiling points across a period?

Answer 36

1) for metals, metallic bonding , melting/boiling points increase across period as there are more electrons that are delocalised, smaller sized ions and ions with a greater charge
2) then a peak (e.g. silicon) as silicon forms a giant covalent structure (macromolecular) many strong covalent bonds
3) the next non-metals are simple molecular with weak London forces however S forms S8 so is slightly higher than P (P4)
4) then decreases to halogens (diatomic) and then lowest at noble gases (monatomic)

Question 37

Give a reason why log is used to plot ionisation energies on a graph

Answer 37

Numbers produced are smaller and easier to plot on a graph

Question 38

What does successive ionisation energies provide evidence for

Answer 38

Existence of quantum shells and the group to which the element belongs to

Question 39

What does first ionisation energy of successive elements provide evidence for

Answer 39

Electron sub - shells