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Transcript
Atomic Structure
HL and SL
2.1 The Atom
Atoms were thought to be uniform spheres
like snooker balls.
Experiments, however, have
shown that atoms consist of a
small nucleus, which contains
particles called protons and
neutrons, around which
electrons orbit in energy levels.
The table below shows the relative masses and charges of
the three sub-atomic particles:
P
a
r
tic
le
P
R
O
T
O
N
N
E
U
T
R
O
N
E
L
E
C
T
R
O
N
M
a
ss
C
h
a
r
g
e
L
o
c
a
tio
n
1
+1
In the nucleus
1
0
In the nucleus
5 x 10-4
-1
Energy levels around
nucleus
You can use Avogadro’s constant to determine the mass
of a proton and an electron. The charge on an electron is
1.60 x 10-19 C.
Atoms are NEUTRAL, this means that they have an
equal number of protons and electrons.
Important Definitions:
The ATOMIC NUMBER is the number of
protons in an atom
The MASS NUMBER is the number of
protons plus the number of neutrons in an
atom .
ISOTOPES are atoms of the same element
with the same number of protons and
electrons but differing numbers of neutrons.
Shorthand:
mass number
A
atomic number
Z
X
symbol of element
So:
mass number
23
atomic number
11
Na
sodium atom
23
11
Na
This sodium atom has
11 protons
11 electrons
12 neutrons
NUMBER OF NEUTRONS = MASS NUMBER – ATOMIC NUMBER
More on ISOTOPES:
As the chemical properties of an atom depend upon their
electron arrangements and isotopes have the same electron
arrangement (as they are atoms of the same element),
isotopes have the same chemical properties.
However, their physical properties especially those that
depend upon mass may vary. These include density and
boiling point.
Using isotopes:
Find out about the use of 14C in radiocarbon dating,
60Co in radiotherapy, 131I and 125I as medical tracers.
Do the benefits outweigh the dangers?
Ions are charged particles formed when atoms either lose or gain
electrons.
Metal atoms tend to lose electrons to form positive ions.
e.g. a sodium ion (Na+) is formed when a sodium atom loses one
electron.
A magnesium atom (Mg2+) is formed when a magnesium atom
loses 2 electrons.
Non-metal atoms tend to gain electrons to form negative ions.
e.g. a chlorine atom gains an electron to form a chloride ion (Cl-)
An oxygen atom gains two electrons to from an oxide ion (O2-)
2.2 Mass Spectrometry
Measures mass of atoms and their relative abundance enabling
relative atomic mass values to be calculated.
3. strong electromagnet
vapourised
sample
4. detector
2. electric field
1. electron gun
Mass spectrometry
Ionisation
The vapourised sample (atoms or molecules) diffuses into
the path of high energy electrons fired from the electron
gun.
These electrons knock out an electron from the sample
producing positive ions.
M(g) + e-  M+(g) + 2eSome doubly charged ions (M2+(g)) are also formed but in
small amounts as it requires more energy to knock out 2
electrons.
Molecules can be broken into ‘fragments’ by the high energy
electrons as they break covalent bonds.
Acceleration
The positive ions are accelerated by an electric field and then
focussed into a beam by passing them through a series of
slits.
Deflection
The beam of fast moving positive ions is then deflected by a
strong magnetic field. The magnitude of the deflection
depends upon the mass to charge ratio (m/z) of the ion. When
the m/z is small, the deflection is large.
The magnetic field can be increased in order to deflect heavier
ions into the detector.
Explain why 40Ca2+ and 20Ne+ are deflected by the same
amount.
Detection
The ions are detected electrically. The ions hit a plate and this
sends a current to an amplifier and then to a recorder.
The chart produced by the recorder is called a ‘mass
spectrum’.
The spectrum on the next slide is for a sample of lead.
This shows that there are 4 isotopes of
lead with mass numbers of 204, 206,
207 and 208.
The height of the peak gives a
measure of the relative abundance of
that peak.
This information can be used to
calculate the Ar of lead.
Using Mass Spectra to calculate
Ar values
The Ar is given by calculating the weighted mean of the
individual relative atomic masses of the isotopes relative to
1/12th of carbon-12.
To do this you need to find the total mass of all of the
isotopes and then divide this by the total abundance.
The total mass of all the isotopes is found by adding
together the mass present of each isotope in the sample
which is determined by multiplying the mass of the isotope
by its abundance.
Ar = total mass / total abundance
Total mass = (204 x 1.5) + (206 x 23.6) + (207 x 22.6) + (208 x 52.3)
= 20724.2
Total abundance = 1.5 + 23.6 + 22.6 + 52.3
= 100
Ar = total mass / total abundance
= 20724.2 / 100 = 207.2
Time to use brainpower:
Sketch the spectrum that would be obtained for a molecule
of chlorine. Given that chlorine exists as 2 isotopes, one
with mass number 35 and the other with mass number 37.
These two isotopes occur in the ratio 3 : 1.
2.3 Electron Arrangements
Why are street lights orange?
Why does copper give a green colour to a bunsen flame when
it is being heated?
Due to their electron arrangements!
The different colours are electromagnetic radiation.
Each colour has a particular wavelength and frequency and is
associated with a particular amount of energy.
To discover the cause of these colours we need to consider the
electromagnetic spectrum.
The velocity of travel of electromagnetic waves is related to its
wavelength and frequency by the equation:
c = f
c = velocity in ms-1,  = wavelength in m, f = frequency in s-1
The energy of electromagnetic radiation is related to its frequency
by the equation:
E = hf
E = energy in J, h is Planck’s constant (6.63 x 10-34 Js)
The smaller the wavelength, the higher the frequency and the
more energy the wave possesses.
Electromagnetic waves have a wide range of wavelengths
varying from radio waves (103 m) to gamma radiation (10-12 m).
The electromagnetic spectrum is a continuous spectrum. You
can get another version of this diagram at
www.chemsoc.org/Networks/Learnnet/data/ds_electromagnetic_spectrum.htm
If the orange light emitted by sodium is passed through a
prism, it is seen as a series of lines, each at a fixed wavelength.
Each element has its own characteristic set of lines. These
lines are known as the line emission spectrum of the element.
These lines become closer together (converge) at the high
energy end of the spectrum.
The simplest spectrum to consider is the hydrogen spectrum:
This is the series of lines found in the visible region. There is
another series in the uv region and several in the ir region.
How do these lines arise?
Niels Bohr published a model in 1913 to explain.
Electrons travel in orbits around the nucleus of
the atom, each orbit is in a fixed energy level.
If the electron is given energy it is promoted to
a higher energy level. As it drops back down it
emits a packet of light called a quantum with a
particular amount of energy. This energy
corresponds to light of a particular wavelength
and shows up as one of the lines in the
spectrum.
The spectrum is not continuous as the electrons only exist at
certain fixed energy levels.
Electrons dropping back down to the lowest energy (n = 1)
emit most energy so this produces the series in the uv region
of the hydrogen spectrum. Electrons dropping down to the
third energy level (n = 3) cause the first series in the ir region.
The visible spectrum arises due to electrons falling to the
second energy level (n = 2).
The lines get closer as the energy levels themselves get closer.
The value n is known as the principal quantum number of the
energy level.
Look at
www.chemsoc.org/Networks/Learnnet/data/int_electron_energy_hydrogen.htm
Each of the energy levels described by the principal quantum
numbers can only hold a certain number of electrons.
The lowest fills first. When one energy level is full with
electrons the next then begins to fill.
Principal quantum
number
1
Maximum number
of electrons
2
2
8
3
8 (18 for HL)
The electron arrangements of the first 20 elements can be found
on the department web pages.