3.1.1 Atomic structure

Question 1

What is the first ionisation energy?

Answer 1

The energy required to remove one electron from each atom of one mole of gaseous atoms.

Question 2

What is the equation for the first ionisation energy?

Answer 2

X(g) → = X+(g) + e-

Question 3

Which factors affect ionisation energy?

Answer 3

-Nuclear charge
-Shielding
-Atomic radius

Question 4

Why does the first ionisation energy generally increase across a period?

Answer 4

-Increasing nuclear charge
-Shielding remains constant
-Therefore there is a stronger attraction of electrons towards the nucleus.

Question 5

Why does aluminium have a lower first ionisation energy than magnesium?

Answer 5

Even though it has more protons,
-The outer electron is shielded by the full 3s orbital
-The 3p orbital is also more further away from the nucleus
-Therefore attraction between the outer electron and the nucleus decreases

Question 6

Why does sulfur have a lower first ionisation energy than phosphorus?

Answer 6

-Sulfur has 4 electrons in the P orbital, so, there’s an electron pair creating electron-electron repulsion.
-Phosphorus has no electron pairs and they have the same shielding and atomic radius.

Question 7

Why do first ionisation energies decrease down a group?

Answer 7

-The outer electron will occupy orbitals further away from the nucleus
-There’s increased shielding from inner electrons
-The effect of the nuclear charge decreases

Question 8

Why does helium have the highest first ionisation energy of all the elements?

Answer 8

-It has the configuration 1s2 and has no shielding.
-So, the electron experiences a very strong electrostatic force of attraction towards the 2 protons in the nucleus.

Question 9

Why is the second ionisation energy of an atom always greater than the first?

Answer 9

The remaining electrons will experience a greater effect of nuclear charge pulling on each electron.

Question 10

Why does atomic size decrease across a period?

Answer 10

-The atomic radius and shielding remains stable
-The nuclear charge increases so, there’s more effective nuclear charge and the electrons are pulled closer.

Question 11

What is ion drift?

Answer 11

Where ions enter a region with no electric field so they just drift through this region.
Lighter ions drift faster as their velocity will be higher whereas heavier ions drift slower as their velocity will be lower. This is because every particles kinetic energy within the mass spectrometer is constant.

Question 12

What does a mass spectrometer tell you?

Answer 12

Relative atomic mass
Relative molecular mass
Relative isotopic abundance

Question 13

What are the 4 different phases of a mass spectrometer?

Answer 13

Electro spray/ electron spray ionisation
Acceleration
Ion drift
Detection

Question 14

Describe elctrospray ionisation.

Answer 14

A sample is dissolved in a volatile liquid
Its forced through a needle connected to a positively charged terminal with a high voltage
Each particle gains a H+ ion
The solvent evaporates

Question 15

Describe electron impact ionisation.

Answer 15

The sample is vaporised
Fired at by high energy electrons
1 electron is knocked off

Question 16

Describe the process of ion acceleration.

Answer 16

They accelerate towards a negatively charged plate as they’re attracted to it.
Lighter ions have a higher acceleration
All ions have the same kinetic energy

Question 17

Describe the process of ion drift.

Answer 17

The ions pass through a hole in the plate and form a beam
They stop accelerating as there’s no electrical field
Lighter ions drift at a faster velocity

Question 18

Describe the process of ion detection.

Answer 18

Lighter ions arrive at the detector and they gain an electron, so there’s a flow of current.
The time taken to reach the detector = mass of the isotope
The size of current = abundance of isotopes

Question 19

Why is the mass spectrometer kept under a vacuum?

Answer 19

To prevent ions colliding with molecules of air.

Question 20

Why are positive ions formed in the mass spectrometer?

Answer 20

To accelerate them to the detector plate
Ions pass through hole forming a beam
So that they can be detected

Question 21

Why is Phosphorus’s sixth ionisation energy much larger than its fifth?

Answer 21

The electron is being removed from the second energy level , which is closer to the nucleus.

Question 22

For bigger molecules, which ionisation technique is used?

Answer 22

Electro spray as bombardment would fragment the molecule.

Question 23

What are the uses of mass spectrometry?

Answer 23

To identify elements
Detecting illegal drugs
Forensic science
Space exploration

Question 24

Explain, in detail, how the relative atomic mass of this element can be calculated from data obtained from the mass spectrum of an element.

Answer 24

Spectrum gives (relative) abundance (1)
And m/z (1)
Multiply m/z by relative abundance for each isotope (1)
Sum these values (1)
Divide by the sum of the relative abundances (1)

Question 25

State how you would collect hydrogen. State the measurements that you would make in order to calculate the number of moles of hydrogen produced.

Answer 25

Hydrogen collection
Using a gas syringe or measuring cylinder/ graduated vessel over water
Measurements
(i) P 1
(ii) T 1
(iii) V 1
Use ideal gas equation to calculate mol hydrogen or mass/Mr
Mol H2 = mol Mg (Mark consequentially to equation)

Question 26

In terms of structure and bonding explain why the boiling point of bromine is different from that of magnesium. Suggest why magnesium is a liquid over a much greater temperature range compared to bromine.

Answer 26

-Structures
Bromine is (simple) molecular / simple molecules
Magnesium is metallic / consists of (positive) ions in a (sea) of delocalised electrons

-Strength
Br2 has weak (van der Waals) forces between the molecules / weak IMFs
So, more energy is needed to overcome the stronger (metallic) bonds or converse.

-Liquid range
Mg has a much greater liquid range because forces of attraction in liquid / molten metal are strong(er)

Question 27

Name the strongest type of intermolecular force between hydrogen fluoride molecules and draw a diagram to illustrate how two molecules of HF are attracted to each other.
In your diagram show all lone pairs of electrons and any partial charges. Explain the origin of these charges.
Suggest why this strong intermolecular force is not present between HI molecules.

Answer 27

Hydrogen bonding
Draw diagram
Dipole results from electronegativity difference
Fluorine more/very electronegative
HI dipole weaker or bonding e- more equally shared

Question 28

Crystals of sodium chloride and of diamond both have giant structures. Their melting points are 1074 K and 3827 K, respectively. State the type of structure present in each case and explain why the melting point of diamond is so high.

Answer 28

NaCl is ionic (lattice)
Diamond is macromolecular/giant covalent/giant atomic/giant molecular
(Many) covalent/C-C bonds need to be broken / overcome

Question 29

Describe the bonding in, and the structure of, sodium chloride and ice. In each case draw a diagram showing how each structure can be represented. Explain, by reference to the types of bonding present, why the melting point of these two compounds is very different.

Answer 29

NaCI is ionic
cubic lattice
ions placed correctly
electrostatic attraction between ions
Covalent bonds between atoms in water
Hydrogen bonding between water molecules
Tetrahedral representation showing two covalent and two hydrogen bonds
2 hydrogen bonds per molecule
Attraction between ions in sodium chloride is very strong
Covalent bonds in ice are very strong
Hydrogen bonds between water molecules in ice are much weaker

Question 30

Explain how the concept of bonding and non-bonding electron pairs can be used to predict the shape of, and bond angles in, a molecule of sulfur tetrafluoride, SF4
Illustrate your answer with a diagram of the structure.

Answer 30

4 bonding electron pairs and one lone pair
repel as far apart as possible
lone pair - bond pair repulsion > bp-bp
pushes S-F bonds closer together
shape is trigonal bipyramidal with lone pair
angles <90
and < 120