Atomic structure

Atoms are made up of protons, neutrons and electrons

• Electrons have a -1 charge , whizzing around the nucleus in orbitals- orbitals take up most of the volume of the atom
• Most of the mass of an atom is concentrated in the nucleus

Proton- relative mass (1)- relative charge (+1)

Neutron- relative mass (1)- relative charge (0)

Electron- relative mass(1/2000)- relative charge (-1)

Ions have different numbers of protons and electrons

• Atoms form by gaining or losing electrons
• Negative ions have more electrons than protons....
• Positive ions have fewer electrons than protons

E.g. Br-

the negative charge means there is 1 more electron than protons

Br has 35 protons, Br- must have 36 electrons

Overall charge= +35-36= -1

Isotopes are atoms of the same element with different numbers of neutrons

• learn definition of isotope= isotopes of an element are atoms with the same number of protons but different numbers of neutrons
• E.g.

• Number and arrangement of electrons that determines chemical properties- isotopes have the same configuration of electrons so they have the same chemical properties
• Isotopes of an element do have slightly different physical properties, because physical properties tend to depend on the mass of the atom

Relative masses are masses of atoms compared to carbon-12

• The relative atomic mass, Ar, is the average mass of an atom of an element on a scale where an atom of carbon-12 is 12.

• The relative isotopic mass is the mass of an atom of an isotope of an element on scale where an atom of carbon-12 is 12.

• The relative molecular mass, Mr, is the average mass of a molecule on a scale where an atom of carbon-12 is 12

Mass spectrometer can tell you the relative atomic mass, relative molecular mass, relative isotopic abundance

4 things happen when a sample is squirted into a (TOF) mass spectrometer

• 1. Ionisation- (electrospray) - the sample is dissolved and pushed through a small nozzle at high pressure. High voltage applied, causing each particle to gain an H+ ion. Sample is turned into gas amde up of + ions. (Electron impact)- sample vaporisd and an 'electron gun' used to fire high energy electrons at it. Knocks 1 electron off each particle, so they become +1 ions.
• 2. Acceleration- + charged ions are accelerated by an electric field, so they all have the same kinetic energy (lighter ions move faster than heavier ones)
• 3. Ion drift- ions enter a region with no electric field, just drift through it. Lighter ions will drift through faster than heavier ones.
• 4. Detection- because lighter ions travel faster in the drift region, they reach the detector in less time than heavier ones . The detector detects charged particles and a mass spectrum is produced.

A mass spectrum is mass/charge plotted against abundance

• y axis- abundance of ions. For an element the height of the peak gives the relative isotopic abundance. If the sample is an element, each element will represent a different isotope of the element
• x axis units are given as a 'mass/charge' ratio. Since the charge on the ions is mostly +1 you can usually assume the x-axis is the relative isotopic mass

E.g.

You can work out relative atomic m*** from a m*** spectrum

1. for each peak, read the % relative isotopic abundance from the y- axis and the relative isotopic mass from the x- axis. Multiply them together to get the total mass for each isotope

2. Add up these totals

3. Divide by 100

if the relative abundance is not given as a percentage, the total abundance may not add up to 100. In this case, do steps 1 and 2 as normal, but then divide by the sum of the relative abundance instead of 100

Mass Spectrometry can be used to identify elements

• Elements with different isotopes produce more than one line in a mass spectrum because the isotopes have different masses. This produces characteristic patterns which can be used as 'fingerprints' to identify certain elements
• Many elements only have one stable isotope, they can still be identified in a mass spectrum by looking for a line at their relative atomic mass

Mass spectrometry can be used to identify molecules

• a molecular ion, M+, is formed in the mass spectrometer when one electron is removed from the molecule
• this gives a peak in the spectrum with a mass/ charge ratio equal to the relative molecular mass of the molecule
• this can be used to help identify an unknown compound, there's more about using mass spectrometry to idenitfy compounds

Electron shells are made up of sub- shells and orbitals

• electrons have fixed energies- they move around the nucleus in certain regions called shells or enegry levels
• each shell is given a number called the principal quantum number - the further away from the nuclues the higher its enegry and the larger the principal qunatum number

This table shows the number of electrons that fit in each type of sub-shell;

This table shows the sub-shells and the electrons in the first 4 energy levels;

Working out electron config by filling the lowest energy levels first

• You can fill up the lowest energy sub-shells first- BUT the 4s sub-shell has a lower energy than the 3d sub-shell (even though the principal quantum number is bigger)
• Electrons fill orbitals singly before they start sharing
• For the config of ions from S and P blocks of the periodic table just remove or add electrons to or from the highest occupied sub-shell

Transition metals behave unusually

• Chromium and Copper are badly behaved, they donate one of their 4s electrons to the 3d sub-shell, because they're happier with a more stable full or half- full shell
• When transititon metals become ions they lose their 4s electrons before their 3d electrons

Electronic strucutre decides the chemical properties of an element

• number of outer electrons decides the chemical properties of an element
• The s block elements have 1 or 2 outer electrons- easily lost to form posititve ions with an inert gas config. Groups 4 to 7 can also share electrons when they form covalent bonds
• Group 0 (inert gases) have completely filled s and p sub-shells and don't bother gaining, losing or sharing electrons making them inert
• the d block elements (transition metals) tend to lose s and d electrons to form positive ions

Ionisation is the removal of one or more electrons

When an electron has been removed from an atom or molecule it has been ionised

The first ionisation energy is the energy needed to remove 1 electron from each atom in 1 mole of gaesous atoms to form 1 mole of gaseous 1+ ions

Here is the equation for the first ionisation energy

- The lower the ionisation energy, the easier it is to form an ion

Factors affecting ionisation energy are...

• Nuclear charge- more protons in the nuclus the more + charged the nucleus is and the stronger the attraction for the electrons
• Distance from nucleus- attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away
• Sheilding- as the number of electrons between outer electrons and the nuclues increases, outer electorns feel less attraction towards the nuclear charge- lessening pull of nucleus

A high ionisation energy means there's a high attraction between the electron and the nucleus and so more energy is needed to remove the electron

Successive ionisation energies involve removing additional electrons

• Each time you remove an electron, there's a successive ionisation energy
• Definiton for second ionisation energy is.....

The second ionisation energy is the energy needed to remove 1 electron from each ion in 1 mole of gaseous 1+ ions to form 1 mole of gaseous 2+ ions

• Here is the equation for the second ionisation energy of oxygen..

The equation for the nth ionisation energy is....

Successive ionisation energies show the shell structure

• Here is a graph of successive ionisation energies

• Within each shell, successive ionisation energies increase- due to electrons being removed from an increasingly positive ion there's less repulsion amongst the remaining electrons so they're held more strongly by the nucleus

Ionisation trends across periods

• Graph below shows the first ionisation of the elements in period 3 

• When you move across a period, the general trend for the ionisation energies is to increase 
• it gets harder to remove the outer electron 
• this can be explained by the fact the number of protons increasing = stronger nuclear attraction

The drop between groups 2 and 3 

Mg 1s2 2s2 2p6 3s2      1st ionisation energy= +738KJ mol-1

Al 1s2 2s2 2p6 3s2 3p1  1st ionisation energy= +578KJ mol-1

• Aluminium's outer shell is in a 3p orbital rather than a 3s. The 3p orbtial hasa slightly higher energy than the 3s orbital, so the electron on average is found further away from the nucleus
• The 3p orbital has additional sheilding provided by the 3s electrons 
• These 2 factors are strong enough to overide the nuclear charge, resulting in the ionisation energy dropping 
• This pattern provides evidence that electron sub-shells exist 

Drops between groups 5 and 6

P  1s2 2s2 2p6 3s2 3p6   1st ionisation energy= +1012KJ mol-1 

S   1s2 2s2 2p6 3s2 3p4  1st ionisation energy= +1000KJ mol-1

• The shielding is identical in the phosphorous and the sulphur atoms and the electron is being  removed from an identical orbital 
• In phosphorous's case the electron is being removed from a singly- occupied orbtial. But in sulphur, the electron is being removed from an orbtial containing two electrons. 
• The replusion between 2 electrons in an orbtial means that electrons are easier to remove from shared orbtials. It's yet more evidence for the electronic structure model. 

Ionisation energies and shell structure 

• If you know the successive ionisation energies of an element you can work out the number of electrons in each shell of the atom 
• providing eveidence for the shell structure of atoms 

• Within each shell, successive ionisation energies increase. This is because electrons are being removed from an increasingly positive ion- less repulsion so they're held more strongly by the nucleus 
• Big jumps in ionisation energies is when a shell is being broken into