chemistry

·         RELATIVE ATOMIC MASS-is the average mass of an atom taking in to account its isotopes and their abundances compared to 1/12 of Carbon-12

·         RELATIVE ISOTOPIC MASS-is the mass of an isotope of an element compared to 1/12 of Carbon-12

·         RELATIVE MOLECULAR MASS-is the average mass of a molecule compared to 1/12 of Carbon-12

·         ISOTOPE-An isotope is an element with the same number of protons and electrons but a different number of neutrons in the nucleus.

·         FIRST IONISATION ENERGY-is the energy required to remove 1 electron from each atom in 1 mole of gaseous atoms to form 1 mole of gaseous

·         NUCLEAR CHARGE-The more protons in the nucleus the more positively charged it will be and the stronger attraction for the electrons

·         ELECTRON SHIELDING-as the number of electrons between the nucleus and the outer electrons increases the attraction of outer electrons to nuclear charge decreases.

·         A HIGH IONISATION ENERGY MEANS A HIGH ATTRACTION BETWEEN ELECTRONS AND NUCLEUS

GOING DOWN A GROUP IN THE PERIODIC TABLE

There are more filled energy levels between the nucleus and the outermost electrons.

These filled energy levels shield the outer electrons from the attractive force of the positive nucleus. As the radius of the atom increases, the distance between the nucleus and the outer electron increases and therefore the force of attraction between the nucleus and outer most electrons is reduced. These factors mean that less energy is needed to remove the first electron from an atom at the bottom of the group compared to one at the top of the group.

GOING ACROSS A PERIOD IN THE PERIODIC TABLE

As we go across the period, there are more protons in each nucleus so the nuclear charge in each element increases. This increases the attractive force acting on the outermost electrons.

So, the nuclear charge increases as each proton is added and another electron is added to the outermost energy level. This electron is poorly shielded from the nuclear charge by the other electrons in its own energy level. Overall, the electrons are drawn closer to the nucleus and are harder to remove.

This would explain a steady increase from left to right. The dips between groups II and III and between V and VI require more explanation.

THE DIP BETWEEN GROUP II AND GROUP III

The electron to be removed from the group III element has to come from a p-orbital. This is higher in energy than the s-orbital of the same main energy level. So removal of this electron is easier and so the ionisation is less.

THE DIP BETWEEN GROUP IV AND V

The electron to be removed from the group VI element comes from a p-orbital that has two electrons. These are repelled slightly more than the electrons in a p-orbital with only one electron. This is because the two electrons are occupying the same region of space. The extra repulsion makes the first ionisation energy a little less.

Subatomic Particle

Relative Mass

Relative Charge

Proton

1

+1

Neutron

1

0

Electron

1/1850

-1

This is the process of Vaporisation, Ionisation, Acceleration, Deflection and Detection.

It is important to note that this only works for gaseous samples so that is why it is vaporised first. Only the process has to happen in a vacuum so that the ions don’t react with particles in the air. In ionisation all samples are made positive even the ones that don’t really ionise like Argon. 

Vaporisation

If the sample is a gas or volatile liquid it is injected into the instrument directly. If the sample is a solid it is turned into a vapor first by heating it.

Ionisation

A beam of electrons from an “electron gun” knocks out electrons from atoms or molecules of the sample so that they form positive ions (cations). Nearly all ions lose just one electron but there are some that lose two.

Acceleration

These positive ions are attracted towards negatively charged plates and are accelerated at a high speed. The speed they reach depends on their mass, the lighter the ions the faster they go. Some ions pass through slits in the plates and from a beam.

Deflection

The beam of ions then moves into a magnetic field at right angles to its direction of travel. The magnetic field deflects the beams of ions into an arc of a circle.  The deflection of an ion depends on its mass/charge ratio. The heavier the ion is, the less it is deflected and the greater the magnetic field strength, the greater the deflection.

Detection

The magnetic field is gradually increased so that ions of increasing mass enter the detector one after another. Ions strike the detector, accept electrons, lose their charge and create a current which is proportional to the abundance of each ion. From this a mass spectrum is produce which is what we use to work out Relative Atomic Mass and Relative Molecular Mass with.

A_r--STEP ONE

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

STEP TWO

Add up these totals

STEP THREE

Divide by 100

With Relative Molecular Mass, it gives a peak in the spectrum with the highest mass (furthest to the right) and this is labeled M. It’s the M_r for the molecule.

DOUBLY CHARGED IONS

Sometimes during ionisation 2⁺ ions are formed and because of their charge are accelerated and deflected more than singly charged ions.

This means that you add up the percentage abundances and times them by the actual mass (because they’ll be peaks at half the actual mass as well because it will behave as though it is a one plus ion). And then divide by one hundred.

     ^{Excluding the d block when atoms react they either loose, gain or share electrons to achieve the electron configuration of the nearest noble gas due to their stable electron configuration. Called becoming IOSELECTRONIC with the nearest noble gas.}

^{Group 1-Alkalin Metals                                                                                  CATIONS}

 ^{·                 Lithium atom                                     lithium ion}

^{Lithium ---------------------------> Li+ + e-}

^{1s2 2s1                                 1s2                          (He) iso electronic}

^{·                 Potassium Atom                                               Potassium ion}

^{K--------------------------------------->K+ + e-}

^{1s2 2s2 2p6 3s2 3p6 4s1     1s2 2s2 2p6 3s2 3p6           (Ar) iso electronic}

^{Group 2- Alkaline-Earth Metals}

^{·                 Magnesium atom                                              Magnesium ion}

^{Mg------------------------------------>Mg2+ + 2e-}

^{1s2 2s2 2p6 3s2                                   1s2 2s2 2p6}

                                                                             ^{Ne iso electronic}

^{·                 Calcium Atom                                   Calcium ion}

^{Ca ------------------------------------>Ca2+ +2e-}

^{1s2 2s2 2p6 3s2 3p6 4s2----------->1s2 2s2 2p6 3s2 3p6}

                                                                             ^{Ar iso electronic}

^{·                 Aluminium atom                                                Aluminium ion}

^{Al-------------------------------------->Al3+ +3e-}

^{1s2 2s2 2p6 3s2 3p1                            1s2 2s2 2p6}            

                                                                             ^{(Ne) iso electronic}

^{·                 Nitrogen atom                                   nitride ion}

^{N+3e- --------------------------------->N3-}

^{1s2 2s2 2p3                                          1s2 2s2 2p6}

                                                                             ^{(Ne) iso electronic}

^{·                 Phosphorus Atom                                             Phosphide ion}

^{P+3e- ---------------------------------->P3-}

^{1s2 2s2 2p6 3s2 3p3                            1s2 2s2 2p6 3s2 3p6}

                                                                             ^{(Ar) Iso electronic}

^{·                 Oxygen atom                                    oxide ion}

^{O +2e- --------------------------------->O2-}

^{1s2 2s2 2p4                                          1s2 2s2 2p6}

                                                                             ^{(Ne) iso electronic}

^{·                 Sulphur atom                                    sulphide ion}

^{S+2e- ---------------------------------->S2-}

^{1s2 2s2 2p6 3s2 3p4                            1s2 2s2 2p6 3s2 3p6}

                                                                             ^{(Ar) iso electronic}

^{Fluorine Atom                                        Fluoride ion}

^{F+e- ----------------------------------->F-}

^{1s2 2s2 2p5                                          1s2 2s2 2p6  (Ne) iso electronic}

       ^{Chlorine atom                                   chloride ion}

^{Cl + e- -------------------------------->Cl-}

^{1s2 2s2 2p6 3s2 3p5                            1s2 2s2 2p6 3s2 3p6 (Ar) iso electronic}

     ^{Bromine atom                                                   Bromide ion}

^{Br + e- ---------------------------------------->Br-}

^{1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p5     1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6    (Kr) iso electronic}

Ionisation Energies

·         Ionisation energies provide the evidence for the existence of energy levels and sublevels in atoms ionisation energy is defined as “the energy required to remove one mole of electrons from one mole of atoms in the gaseous state to produce a mole of gaseous unipositive ions

·         X(g)--->X+(g) + e-

·         Electrons can be removed by hitting the atom with a beam of electrons from an electron gun, allowing us to measure the energy it takes to remove the electron

·         All ionisations require energy, this is  an endothermic process and have a positive enthalpy change

·         1^st ionisation energies need the least amount of energy as they are being removed from a neutral atom

·         The energy needed increases in succession as you get closer to the nucleus and are being removed from a 1+ ion, 2+ ion, 3+ ion ect = Successive ionisation energies

Value of first ionisation energy depends on:

·         Number of protons in nucleus(nuclear charge) this causes strong forces of attraction between the Ps and Ns

·         Energy level(period) of the electron being removed (as energy level increases the forces of attraction decrease). Meaning more shielding

      1^st Ionisation energy decreases going down a group because: Electrons come from a higher energy level(further from the nucleus) Nuclear charge increases but the outer electron is shielded from the nuclear charge by the inner electrons

  Trends across a period

    1^st ionisation energies show a general increase due to:The nuclear charge is increasing(no of protons) making it more difficult for the electron to be removed. The electron comes from the same energy level(same shielding)

 The break in pattern- Mg and Al -  Al outer electron in the 3p orbital   Mg outer electron in the 3s orbital.  Al 3p is of a slightly higher energy than the 3s of Mg therefore needs less energy to removed

Break in pattern-P and S   P, each electron is in an orbital, S, one of the 3p orbitals has two electrons

·         The pair of electrons in the 3p orbital is easier to remove due to spin pair repulsion than the unpaired outer electron of  P in the 3p orbital

These breaks that occur across the period show evidence that there is s and p sublevels. Predicted by the quantum theory and Schrödinger’s equation

^{·          when atoms react together they share or transfer electrons to achieve a more stable electron arrangement usually a noble gas.}

^{·           the noble gas has a full outer main level of electrons which means that they are very un-reactive}

^{·          three types of strong chemical bonds Ionic, Covalent and Metallic}

^{Ionic bonding}

^{·          ionic bonding occurs between metals and non metals}

^{·          the electrons are transferred from metals to non metals, this is due to metals having one, two or three electrons in their outer main levels and the easiest way for them to attain the electron structure of the nearest noble gas is to loose these electrons.}

^{·          positive and negative ions are formed}