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Atoms
Sub-atomic particles:
Particle
Proton
Neutron
Electron
Mass (kg) : relative
Charge (actual) : relative
-27
1.672x10 : 1
1.674x10-27 : 1
9.109x10-31 :negligible
+1.602x10-19 : +1
0 : 0
-1.602x10-19 : -1
A nuclear atom has a nucleus at its centre composed of protons and
neutrons, with electrons arranges in energy levels around it. Most of
the mass of the atom is contained within the nucleus. Any atom has
equal numbers of protons and electron, meaning it is neutral overall,
but because neutrons have no charge, the number of them in the
nucleus has no effect on the overall charge.
Atomic number/ proton number = the number of protons in the nucleus.
All the atoms of a particular element have the same proton number.
(Z)
Mass number = the number of subatomic particles in the atom’s
nucleus. (A)
A
Z
X
Isotopes are atoms that have the same number of protons but different
numbers of neutrons in their nucleus. Different isotopes are
chemically identical to each other because they contain the same
numbers of protons and electrons.
Mass Spectrometry
Ionisation:
The sample must be gaseous and must be ionised to be accelerated and
detected. A high energy electron gun fires electrons at the atoms,
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normally shooting off one electron before they are attracted to the
acceleration plates. However, it frequently happens that two
electrons are removed, or that the molecule is smashed into smaller
pieces. This is called fragmentation.
Acceleration:
The negative charge on the plates is adjustable to suit the particles
being investigated. Small molecules will not need a lot of
acceleration; big heavy ones will need a larger negative charge on
the plates to make them move fast enough. Once the particles have
passed through the slits, they form a narrow focussed beam of
positively charged particles.
Deflection:
Big, heavy molecules will need a strong magnetic field to deflect
them; small light molecules will need a smaller magnetic field. The
more charge a particle is carrying the more it will be deflected by a
magnetic field. When the field is being adjusted, the operator needs
to consider the charge and the mass of the particles in the machine
at the same time – they use the mass to charge ratio to combine both
factors. A big mass to charge ratio means a bigger field is needed
to bend the particles round the spectrometer tube.
Detection:
The ionised samples cause a tiny current in the detector; the more
ions that hit the detector, the bigger the current. The current on
the detection plate is recorded in a mass spectrum.
Detecting Isotopes
Relative isotopic mass:
RIM = mass of one atom of the isotope X12
mass of one atom of
12
C
M/Z ratio (mass to charge ratio):
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The relative mass of the ion, m, divided by its charge, z.
An ion with a large m/z ratio will be deflected more than an ion with
a large m/z ratio.
The detector reports the mass to charge ratio of the particles it
finds.
Most molecules just lose one electron then get detected.
This means that their mass to charge ratio will be mass ÷ charge =
Mr ÷ (+1)
E.g. for CH4, ionisation will give (CH4)+
Mr of (CH4)+
= 16.0
Charge = +1
Mass to charge ratio (m/z)= 16.0÷1=16.0 (no units)
Sometimes we lose two electrons:
m/z of 48Ti2+ = 48.0 ÷ 2 = 24.0
This can cause problems. If we had a sample that contained 48Ti and
24Mg, we could get a 24Mg+ ion formed with a m/z ratio of: 24.0 ÷ 1
= 24.0.
By mass spectrometry, 24Mg+ is indistinguishable from 48Ti2+.
The peak to the furthest right of the spectrum gives the Mr of the
starting molecule. Tiny peaks to the right of the molecular ion peak
are caused by isotopes, e.g. 13C.
We can use mass spectrometry to identify isotopes.
E.g. Cl has two different isotopes: 35Cl and 37Cl.
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If each atom of the isotope loses one electron, we will get peaks at
m/z 35 and 37. The ratio of the peak heights tells us the relative
abundance of the two isotopes in chlorine.
More Mass Spectrometry
To work out the relative atomic mass of an atom using a spectrometer:
Ar = Sum of (m/z ratio x relative abundance)
Sum of relative abundances
The relative molecular mass of a substance deals with compounds; it
is the mean mass of a molecule compared to one-twelfth the mass of a
12
C atom.
The mass spectrometer provides information about molecules – the
high-energy electrons in the ionisation stage can remove electrons
from molecules, creating a molecular ion. If there is a single charge
on the molecular ion, its m/z ratio is the same as its relative
molecular mass.
The peak caused by the molecular ion is the peak on the far right of
the mass spectrum, with the highest m/z ratio.
Arranging Electrons
Electrons can only occupy certain energy levels. Each energy level
contains one or more sublevels.
S sub-levels: can hold 2 electrons
P sub-levels: can hold 6 electrons
D sub-levels: can hold 10 electrons
Electron configuration:
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Electrons fill energy levels according to the Aufbau principle; means
the building up principle, so electrons fill energy levels in order
of increasing energy.
Orbitals:
Electrons are moving all the time so you cannot work out exactly
where they are, but instead we can guess the regions called orbitals.
Pauli Exclusion Principle – each orbital can hold a maximum of two
electrons, and they exist in 2 states: spin up or spin down. When two
electrons occupy one orbital, they must have opposite signs or the
repulsion is too much.
S sub-level: has 1 orbital
P sub-level: has 3 orbitals
D sub-level: has 5 orbitals
Hund’s Rule: electrons occupy orbitals in a particular sub-level
following Hund’s rule. They occupy orbitals as single unpaired
electrons all with the same spin. They only begin to pair up when
there are no more empty orbitals in the sub-level. An empty orbital
is called a vacant orbital.
The 4s sub-level is lower in energy than 3d, so it fills first
Cr and Cu pull a 4s electron into the 3d orbital to minimize
repulsion.
The periodic table is all in order of elements with increasing atomic
number. The different blocks on the periodic table show where the
highest energy outer electron is.
Electron configuration of ions:
Atoms may lose or gain electrons, and when this happens they become
ions. Hydrogen is the simplest ion. When the electron in hydrogen is
lost, the particles become a positively charged hydrogen ion.
Metal ions in the s and p blocks:
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Metal atoms can lose electrons to form positive ions. In all cases
the electrons are lost from the highest energy level
Metal ions in the d block:
Although the 3d sub-levels is higher in energy than the 4s, the
energy level 3 is overall lower in energy than 4, so when d block
atoms form ions, they lose electrons from the first 4s sub-level,
then the 3d sub-level.
Non-metal ions:
Non-metal atoms gain electrons to form negative ions. In all cases
the highest occupied energy level is the one that gains electrons.
Evidence for Energy Levels
Ionisation energy:
The first ionisation energy of needed to remove one electron from a
gaseous atom. Electrons are negatively charged, so are attracted to
the positively charged nucleus. Overcoming that force of attraction
needs energy.
The first ionisation energy decreases as you go down a group. In each
case an electron from the highest energy levels is being removed. As
you go down group two, each successive element has more occupied
energy levels. The distance between the outer electrons and the
nucleus increases, so electrons in the highest energy levels are less
strongly attracted to the nucleus, so even though the nuclear charge
increases down the group, less energy is needed to remove an outer
electron.