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Mass Spectroscopy

Mass Spectroscopy

Mass Spectrometry

A mass spectrometer machine can be used to analyse elements or compounds.

Time Of Flight (TOF) mass spectrometry

  • There are 4 things that occur when a sample is squirted into a TOF mass spectrometer:
  • Electronspray ionisation:
  • The sample is dissolved in a polar solvent and is pushed through a small nozzle at a high pressure.
  • A high voltage is applied which results in the particles in the sample to lose an electron.
  • These ionised particles are then separated from the solvent, leaving a gas of positive ions
  • Acceleration:
  • The positive ions are accelerated by an electric field
  • The particles need to be positively charged to accelerated by the electric field
  • The electric field gives the same kinetic energy to all of the ions regardless of size
  • The ions with a lower m/z (mass/charge) ratio experience greater accelerations
  • Ion Drift:
  • The ions leave the electric field at a constant speed with equal amounts of kinetic energy
  • They enter a region with no electric field and drift through it at the speed they left the electric field
  • Ions with lower mass/charge rations will be drifting at a higher speed
  • Detection:
  • As ions with a lower mass/charge ration travel through the drift region at a higher speed, they will reach the detector in a quicker time relative to the heavier ions
  • The detector detects the change in current created when the ions (cations as they have had an electron removed) collides with the detector plate. It then records how long they took tp pass through the spectrometer. The data produced can then be used to determine the mass/charge value need to produce a mass spectrum
  • Time of Flight uses alcohol groups (ethanol) which has a space hydrogen bond. This hydrogen bonds with X (x being the substance in question) causing it to break apart

X + H+ -> XH+

  • The various atoms are accelerated towards a negative plate which then measures which atoms get there at what time to calculate the masses.
  • In older methods of spectrometry the electron ionisation. An electron gun is fired at an atom, causing it to lose 1 electron as it is knocked out.

X(g) -> X+(g) + e

  • The test however is not greatly accurate as often X+ would be further broken down.

Calculating Relative Atomic Mass

  • Averages are used to find the RAM of an element
  • Take the value of the x axis and multiply it to the Y axis.
  • Add the value of step 1 for all major peaks on the graph
  • Divide by the total number of the Y axis

e.g.

polarity

(15 x 18) + (29 x 62) + (31 x 100) + (32 x 63)

243

7184

243

= 29.56

Identifying Elements

  • Mass spectrometry can be used to identify elements
  • Used in determining different isotopes of an element as it produced more than one line in a mass spectrum as the isotopes all have different masses
  • The patterns produced can be used as a ‘fingerprint’ and can identify specific elements that are contained
Hund’s Rule

Hund’s Rule

Every orbital in a subshell is singly occupied with one electron before any one orbital is doubly occupied, and all electrons in singly occupied orbitals have the same spin.

Single electrons occupy all empty orbitals before they start to form pairs in orbitals. Two electrons in the same orbital have a repulsion between them due to their negative charge. The more stable configurations is with single electrons in different orbitals.

Evidence for Hund’s Rule

The first ionisation energy for the elements in period 3 has a general increase.

Sulphurs value is below that of phosphorus. The highest energy electrons are both in 3p sub levels this is evidence for Hund’s Rule.

Phosphorus has 3 electrons in its 3p sub level whereas sulphur has 4.

The lowest first ionisation energy for sulphur is because it has a pair of electrons in one of the 3p orbitals. Mutual repulsions between these two electrons make it easier to remove one from the other.

Phosphorus ionisation energy = 1012 KJ mol-1

Sulphur ionisation energy = 1000 KJ mol-1

Spin Diagrams

Spin Diagrams

Spin diagrams give a visual representation of electron configurations along a graph. Each arrow represents an electron in the atom. The direction of the arrow represents the spin the electron has.

The Pauli Exclusion Principle

The Pauli Exclusion Principle states that each orbital may contain no more than two electrons.

This introduced a property of electrons called spin, which has two states ‘up’ or ‘down’. The spin of an electron in the same orbital must spin in different directions.

Hund’s Rule

Hund’s rule states that single electrons occupy all empty orbitals within a sub-level before they start to form pairs in orbitals.

If two electrons enter the same orbitals there is a repulsion between them due to their negative charges. The most stable configuration is with single electrons in different orbitals.

Hydrogen

1s1

Helium

1s2

Lithium

1s2 2s1

Beryllium

1s2 2s2

Boron

1s2 2s2 2p1

Carbon

1s2 2s2 2p2

Nitrogen

1s2 2s2 2p3

Oxygen

1s2 2s2 2p4

Fluorine

1s2 2s2 2p5

Neon

1s2 2s2 2p6

Sodium

1s2 2s2 2p6 3s1

Magnesium

1s2 2s2 2p6 3s2

Aluminium

1s2 2s22p6 3s3p1

Silicon

1s2 2s22p6 3s3p2

Phosphorous

1s2 2s22p6 3s3p3

Sulphur

1s2 2s22p6 3s3p4

Chlorine

1s2 2s22p6 3s3p5

Argon

1s2 2s22p6 3s3p6

Potassium

1s2 2s22p6 3s3p4s1

Calcium 

1s2 2s22p6 3s3p4s2

After the 1s2 2s22p6 3s3pformat, there is a break in the logical pattern in that 4s sub-level fills before the 3d sub level. 
This is because the 4s sub level is of lower energy than the 3d sub level. Once the 3dd sub level is filled, then the 4p sub level is filled

Electron Configuration of Transition Metals

When transition metals form ions, it is the 4s electron that are removed before the 3d electron.

Vanadium:

1s2 2s22p6 3s3p4s3d3

Nickel:

1s2 2s22p6 3s3p4s3d5

If there is a positive charged ion then to remove electrons, they are taken from the highest energy level.

e.g. Ni = 1s2 2s22p6 3s3p4s3d5

Ni2+= 1s2 2s22p6 3s3p3d 

As there is a 2+ charge on the second nickel, two electrons are taken from the highest energy level, which is the 4s shell.

e.g. Fe = 1s2 2s22p6 3s3p4s3d6 (or) [Ar] 4s2 3d6

Fe2+ = 1s2 2s22p6 3s3p3d6 (or) [Ar] 3d6

Energy Levels

Energy Levels

  • The energy levels are called Principle Energy Levels
  • The principal energy levels contain sub-levels. Each principal energy contains a different number of sub-levels.

Principal Energy Level

Number of Sub-levels

1

1 (S)

2

2 (S,P)

3

3 (S,P,D)

4

4 (S,P,D,F)

Sub-level

Max number of Electrons

S

2

P

6

D

10

F

14

Principal Energy Level, n

Sub-level

Max Number of Electrons

1

1s

2

2

2s,2p

8

3

3s,3p,3f

18

4

4s,4p,4d,4f

31

e.g.

  1. a) Sodium = S block
  2. b) Iron = D block
  3. c) Sulphur = P block
  4. d) Chlorine = P block
  5. c) Radium = S block
  6. e) Neon = P block

Electron Configurations

Element

Electron Configuration

Hydrogen

1s1

Helium

1s2

Lithium

1s1, 2s1

Beryllium

1s1, 2ss

Boron

1s1, 2ss,2p

Carbon

1s1, 2ss,2p2

Neon

1s1, 2ss,2p3

Oxygen

1s1, 2ss,2p4

Fluorine

1s1, 2ss,2p5

Neon

1s1, 2ss,2p6

Sodium

1s1, 2ss,2p6, 3s

Relative Atomic and Molecular Masses

Relative Atomic and Molecular Masses

Avogadro’s number: 6.02×1023

(Ar) Relative Atomic Mass

mass of the protons and Neutrons
Mean mass of 1 atom of the element
1/12 the mass of atom of Carbon-12

(Mr) Molecular mass

The sum of the Ar for all of the atoms in the formula
Mean mass of the molecule
1/12 the mass of atom of carbon-12

The Mole

The weight of the relative atomic mass of any element e.g. carbon-12 = 12.00g (1 mole)

To weigh one atom of an element divide the moles in grams by Avogadro’s number

e.g. 1 mole of carbon-12 = 12g

6.02×1023

Examples

a) 1 mole of hydrogen = 1.0g
b) Mass of 1 atom of H = 12g
6.02×1023

a) 1 mole of Copper = 24.3g
b) Mass of 1 atom of Cu = 3g
6.02×1023

a) 1 mole of Hydrogen Molecule = 2.0g
b) Mass of 1 atom of H2 = 0g
6.02×1023

a) 1 mole of Calcium Nitrate= 64.1g
b) Mass of 1 atom of Ca(NO3)2 = 1g
6.02×1023

Relative Masses

Relative Masses

The relative mass of an electron is sometimes quoted as zero as its weight is very nearly zero especially compared to the weight of the proton and neutron.

The protons are found in the centre of the atom in the nucleus. Whereas the electrons orbit in shells around the nucleus

The atomic number of an element is the number of protons I the nucleus. It has the symbol Z.

Element

Atomic Number

Mass

Number

Number of

Protons

Number of

Electrons

Number of

Neutrons

Hydrogen

1

1

1

1

0

Helium

2

4

2

2

2

Lithium

3

7

3

3

4

Beryllium

4

9

4

4

5

Boron

5

11

5

5

6

Carbon

6

12

6

6

6

Nitrogen

7

14

7

7

7

Oxygen

8

16

8

8

8

Fluorine

9

19

9

9

10

Neon

10

20

10

10

10

Sodium

11

23

11

11

12

Magnesium

12

24

12

12

12

Aluminium

13

27

13

13

14

Silicon

14

28

14

14

14

Phosphorus

15

31

15

15

16

Sulphur

16

32

16

16

16

Chlorine

17

35.5

17

17

17

Argon

18

40

18

18

22

Potassium

19

39

19

19

29

Calcium

20

40

20

20

20

Isotopes

Isotopes are atoms of the same elements with the same element with the same atomic number but different mass numbers.

Number of neutrons = Mass Number – Atomic Number

Isotopes

Protons

Neutrons

Electrons

Chlorine-35

17

18

17

Chlorine-37

17

20

17

Silicon-28

14

14

14

Silicon-29

14

15

14

Silicon-30

14

16

14

Examples

  1. About 75% of naturally occurring chlorine is Chlorine-35 and 25% of chlorine-37

Average= (75×35) x (25×37) =35.5

100

  1. There are 2 isotopes of Bromine Br-79 has an abundance of 50.69%. Br-81 has an abundance of 49.31%. What is its mass?

(50.69 x79) x (49.31 x 81) = 79.9862

100