Chemistry AS Unit 1 Topic 2: Atomic Structure and the Periodic Table

Edexcel AS, mass spectrometry, ionization energys and electron shells

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  • Created by: Alisha
  • Created on: 03-01-11 17:23

Mass Spectrometry

The mass spectrometer measures the masses of positive ions formed from atoms as atoms themselves cant be directly weighed. The stages in a mass spectrometer are:

  • Vapourisation: the sample must be a gas so it can move throught the mass spectrometer-> so it is heated to very high temperatures.
  • Ionization: Sample is bombarded with high-energy high speed electrons, knocking off an electron off the atoms/molecules forming positive +1 ions.
  • Acceleration: The ions are accelerated by an electric field.
  • Deflection: The sample is passed through a agnetic field where ions are deflected depending on their mass:charge ratio. This can be steadily increased.
  • Detection: Only ions with correct mass:charge ratio will be deflected round to the detector where it is recorded.

Most of the ions formed have a charge of +1 so the mass:charge ratio corresponds to the mass of the ion.

A Mass Spectrum shows the masses of the ions detected and their relative abundance - isotopic composition of an element.

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Mass Spectrometry and compounds

At the second stage, not only are electrons knocked off but the molecule is fragmented.

The mass spectrum of an organic molecule comprises of a number of peaks corresponding to the various fragments of differing masses produced during electron bombardment.

The peak with the largest mass:charge ratio is due to the molecular ion (M+) (also called the parent ion).

It is formed when the molecule loses just 1 electron and is not fragmented.

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Calculating Relative Atomic Mass

A simple atomic mass spectrum shows the proportions of all the isotopes present in a sample of an element. The RAM of an element can be found from the atomic mass spectrum. The RAM is the weighted average of the masses of the atoms in a naturally occurring sample of the element.

To calculate the weighted average:

1) measure each peak height

2) total them

3) calculate percentage relative abundance:

(Peak height / total) x 100%

4) the weighted average is the sum of the products of each percentage abundance and its corresponding mass number

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Uses Of Mass Spectrometer

There are a wide range of applications:

  • in the pharmaceutical industry - identifying molecules with potential for use as drugs. the likely structure of new compounds can be identified from its mass spectrum.
  • Drug testing - detecting illegal drugs in athletes' urine samples (e.g. anabolic steroids). each molecule can be identified by its unique mass spectrum.
  • in Space - identification of molecules detected by planetary probes.
  • Radioactive dating - percentage abundance of radioactive isotopes are measured.
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Ionization Energy

This is the amount of energy required to remove an electron from an atom. It is Endothermic and is measured as the energy required to remove 1 mole of electron from 1 mole of atoms in the gaseous state.

  • The First ionization energy gives an idea of how easily an atom loses an electron to form a 1+ ion. It is the energy required to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of plus 1 ions under standard conditions of 25 degrees celcius and 1 atm pressure. e.g.
     Na (g) --> Na+ (g) + e-
  • The second ionization energy is the energy required to remove an electron from one mole of gaseous plus 1 ions to form one mole of plus 2 ions under standard conditions of 25 degrees celcius and 1 atm pressure. e.g.
    Na + (g) --> Na 2+ (g) + e-
  • The total energy required to form a 2+ ion is the sum of the first and second ionization energies.
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Evidence for electron shells

  • The first ionization energy is lowest as the electron is removed from the outer shell which is furthest from the attractive pull of the nucleus with complete energy levels shielding the outer electron from the nuclear charge. The ratio of p+ : e- is 1:1.
  • There is a "jump" between certain ionization energys as the 2nd electron must be removed from an energy shell closer to the nucleus so there is less shielding and also a greater attractive force as the ratio of p+: e- has increased so the nucleus has a greater hold on each e-.
  • For other ionization energys, there is a steady increase as the only change now is the ratio of p+:e- so the effective nuclear charge is increasing but they are all in the same energy level so there is no jump as shielding and distance from the nucleus are the same.
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First Ionization energy against atomic numbers

This provides evidence for electron sub-shells.

  • Peaks correspond to Noble Gases - Noble gases require the most energy to remove an e-
  • Peaks correspond to the filling of an energy level - peaks are found at intervals of 2,8,8,18 which suggests the number of e-s in each energy level.
  • There is a repeating pattern for the 1st 20 elements so it is a periodic property (repeated every 8 atomic numbers).
  • Trend down a group- becomes less endothermic as each time you go down 1 place in a group you add another energy level meaning the e- is being removed further from the nucleus so there is increased shielding of the nuclear charge.
  • Although there is a general increase in ionization energy across a period it is not steady as there are breaks in values (e.g Be/B or N/O and 8 units later in period 3 Mg/Al and P/S)

This is evidence the e-s in an energy level are not all in the same place.

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Orbitals and sub-shells

In an energy level there are sub-energy levels called Orbitals.

An orbital is the region in space where there is the greatest probability of finding an e-.

There are four types of orbitals:

  • s -spherical
  • p - dumb-bell shaped (3D in the x y and z directions)
  • d
  • f

If they are available in an energy level there are:

  • 1 x S orbital
  • 3 x P orbital
  • 5 x D orbital
  • 7 x F orbital.
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Electronic Configurations

Each orbital can hold a maximum of 2 electrons (Of opposite spin). Shown in boxes as half arrows.

The energies of subshells differ from each other because of the effect of penetrating towards the nucleus.

Electron Density Maps of electron orbitals show how the electron cloud is distributed within the orbital.

The electronic configuration of an atom specifies the number of e-s in each electron shell or sub-shell. Electronic configurations for atoms in their ground state ( lowest energy state) generally reflect the order in which shells and sub-shells are filled.

Lowest Energy shells are filled first - 1s 2s 2p 3s 3p 4s 3d 4p.

Electrons populate orbitals singly at first before pairing up explaining why removing the outer e- from boron requires less energy than for beryllium- the outermost e- goes to a new p shell rather than filling the s. The p-sub shell is is at a higher energy level than the s so less energy is required to remove an e- from it.

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Electronic Configurations

The electronic structure of an atom determines the chemical properties of that element. The outer shell electrons are those which take part in chemical reactions (bonding) and have the most influence over the chemistry of the atom.

The periodic table is divided into blocks - s block, p block and d block- in which the last electrons to be added occupy specific orbitals - s p or d respectively. this helps predict electronic configurations from the order of electron filling of sub shells.

  • The s-block elements have 1 or 2 outer electrons and are reactive metals.
  • The d-block elements are transition metals- have similar chemical properties as the electrons are not being added to outer electron shells; the d orbitals are inside the s orbitals.
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Periodic Properties

A Periodic property has a trend which repeats across periods. e.g.

  • the first ionization energy tends to increase across a period and decrease down a group.
  • atomic radius decreases across a period. going right across a period, the outer electron is drawn closer to the increasingly positive nucleus. There is constant shielding by the full inner shells. Overall there is a greater force of attraction between the nucleus and outer electron so it becomes more difficult to remove it from the atom.
  • going down a group, the outer electron is further and further from the nucleus and electron shielding by the additional inner shells reduce the attraction to the nucleus - it becomes easier to remove it from the atom.
  • Melting temperature is also a periodic property and depends on its structure and bonding. Metallic structures are strong meaning it takes a lot of energy to break the bonds that hold the atoms/ions tgoether. Giant molecular structures are held together by covalent bonds which are strong- it is very difficult to melt. Simple molecular structures are held by weak forces of attraction between molecules- requiring little energy to overcome these intermolecular forces of attraction, the melting temperatures are low.
  • Melting temperatures rise from group 1-2-3 as the number of shared delocalized electrons increases, thus increasing the strength of metallic bonding.
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