Download Atomic Structure

Contents
 The Structure of Atoms
 Isotopes of elements
 The mass spectrometer
 Electron Arrangement
 The nature of the electron HL
Objectives
 State the position of protons, neutrons and electrons in




the atom
State the relative masses and charges of protons,
neutrons and electrons.
Define the terms mass number (Z) and isotopes of an
element.
Deduce the symbol for an isotope given it’s mass number
and atomic number.
Calculate the number of protons, neutrons and electrons
in atoms and ions from the mass number, atomic number
and charge.
 Compare the properties of the isotopes
 Discuss the uses of radio isotopes
Atomic structure
Properties of sub atomic
particles
A single Hydrogen-1 atom contains one proton
and one electron.
An electron has negligible mass compared with
that of a proton.
As 1 mol of hydrogen atoms contains 6.02 x 1023
hydrogen-1 atoms, and the mass of 1 mol of
hydrogen-1 atoms is 1.00g, we can calculate the
mass of a single proton.
It will simply be equal to the reciprocal of
Avogadro’s constant in grams.
A neutron has a mass that is almost the same as, but
very slightly larger than, that of a proton, whereas an
electrons mass is only about 1/1836 that of a proton.
These masses can be simplified by assigning relative
masses ( known as atomic mass units). Similarly,
electrons and protons have relative charges.
Can you remember the relative masses and charges
of each sub atomic particle? Copy out and complete the
table below.
Sub atomic
particle
Proton
Relative Mass
Neutron
1
1
Electron
0 or 5 x 10 -4
Relative Charge
+1
0
-1
Isotopes of elements
All atoms can be characterized by two numbers: the atomic number and the
mass number.
The atomic number (Z) is simply equal to the number of protons in the
nucleus of the atom. As atoms are electrically neutral this will also be equal
to the number of electrons in the atom. Atoms of different elements will
have different atomic numbers
The mass number (A) is equal to the number if protons and neutrons
(collectively known as nucleons) in the nucleus of the atom.
Different atoms of the same element may have the same mass number, or
a different mass number if they contain a different number of neutrons.
Atoms of the same elements must contain the same number of protons, but
if they contain a different number of neutrons they are known as
ISOTOPES
Isotopes are atoms of the same
element which have the same
number of protons but a different
number of neutrons
More on isotopes
The symbol for an isotope of an atom is written in the
form:
Examples include:
1
1
17
3
1
1
H H
hydrogen
35
2
deuterium
37
Cl Cl
17
Chlorine-35 Chlorine-37
A
Z
X
H
tritium
235
92
U
238
92
U
Uranium-235 Uranium-237
Isotopic ions
If the symbol for an ion of an isotope rather than that of
an atom is required, then the charge carried by the ion is
written on the top right hand side, some examples are
shown below:
Symbol
27
13
23
11
31
15
Al
Atomic
Mass
number Z number
A
Number
of
protons
Number
of
neutrons
Number
of
electrons
13
27
13
14
13
Na+
11
23
11
12
10
P3-
15
31
15
16
18
Properties of isotopes
The chemical properties of
atoms depend on their outer
electrons.
As all isotopes of the same
element have the same
arrangement of electrons,
their chemical properties
are identical.
Isotopic
Water
Boiling point at 1atm
H2O
100.0
D2O
101.4
T2O
101.5
H217O
100.1
H218O
100.2
D218O
101.5
HDO
100.7
HTO
100.8
However because they have different masses, their
physical properties such as density, rate of diffusion,
melting point and boiling point will differ.
Radio isotopes
•
•
•
Isotopes have many uses in chemistry and beyond.
Many but by no means all, isotopes of elements are
radioactive, because the nuclei of these atoms break
down spontaneously.
When they break down, these radioisotopes emit
radiation, which may be one of three types:
• Gamma (γ) radiation is highly penetrating,
• Beta (β) radiation which can be stopped by a thin
•
sheet of aluminium
Alpha (α) radiation which can be stopped by a few
centimeters of air.
Some uses of radioisotopes
 Radioisotopes can occur naturally or be created
artificially.
 Their uses include
 Nuclear power generation,
 The sterilization of surgical instruments in
hospitals
 Crime detection
 To find cracks and stresses in metals
 The preservation of food
 Dating artifacts
 Treating and diagnosing illness in medicine.

Describe and explain the operation of a mass
spectrometer

Describe how the mass spectrometer may be used to
determine relative atomic mass using the 12 – Carbon
scale

Calculate non-integer relative atomic masses and
abundance of isotopes from given data.
The Mass Spectrometer
If you throw three balls, which
all have the same diameter, at
exactly the same speed on a very
windy day. They present the
same profile to the wind, but
their masses are distinctly
different – e.g. a foam ball, a
tennis ball and a cricket ball.
Where would each ball land?
Throw direction
Cricket ball
wind
Tennis ball
Foam ball
Atoms have masses in the range of about 1x10-24 to 1x10-22 grams, and
you can’t weigh them in any conventional sense. You can, however, get
around the problem.
Atoms can be deflected in a similar way to the balls in the wind by
magnetic fields – provided the atom is first turned into an ion.
The sample is first vaporized – turned to a gas
The vapour is ionized by bombarding it with a stream of high-energy
electrons from an electron gun to generate positive ions.
M(g)
Vaporised
atom
+
e-
→
High-energy
electron
Mg+(g)
+
2e-
Unipositive ion
The positive ions pass through slits in negatively charged parallel
plates, where they are accelerated.
The ions are deflected by a magnetic field. The amount of the deflection
depends both on the mass of the ion and on its charge. Heavier and less
highly charged ions will be deflected less than lighter and more highly
charged ions.
Ions with a particular mass to charge ratio (m/z) are then recorded on a
detector, which measures both the mass-to-charge ratio and the
relative amounts of all the ions present.
In practice, the machines electron beam energy can be adjusted so
that only positive ions with a single charge are detected, so that the
mass-to-charge ratio is the same as the mass.
The mass spectrometer produces a mass spectrum.
From the mass
spectrum of an
element it’s relative
atomic mass Ar can
be calculated, as it is
equal to the
weighted mean mass
of all the naturally
occurring isotopes
of that element
relative to 1/12
carbon-12.
Detector current
(arbitrary units)
12.17
9.13
6.83
2.60
2.60
m/z (mass charge ratio)
The mass spectrum above is of naturally occurring germanium.
From the mass spectrum it can be seen that;
total detector current is = (6.83 + 9.13 + 2.60 + 12.17 + 2.60) = 33.33
The relative abundance of germanium-70 =
6.83
33.33
x 100 = 20.5%
The relative abundance of all the isotopes can be calculated in a
similar way.
Copy out this diagram in your notes and calculate the relative
abundance of the other germanium isotopes.
Isotope
70
72
Relative abundance / %
20.5
27.4
73
7.8
74
36.5
76
7.8
The relative atomic mass of germanium is given by
Ar = (70 x 20.5) + (72 x 27.4) + (73 x 7.8) + (74 x 36.5) + (76 + 7.8)
= 72.7
100
This is the number we see on the periodic table
The mass spectrum for
chlorine
 Chlorine has two isotopes,





35Cl
and 37Cl, in an approximate
ratio of 3 atoms of 35Cl to 1 atom of 37Cl.
So obviously the mass spectrum will consist of two lines
at m/z 35 and 37, with the 35 line three times higher
than the 37 line. But there is more……
Chlorine consists of molecules, not individual atoms
So when chlorine passes into the ionisation chamber, an
electron is knocked off the molecule to give a molecular
ion, Cl2+
Some of these ions fragment (fall apart) to give a
chlorine atom and a Cl+ ion.
Cl2 +
→
Cl
+
Cl+ (fragmentation)
 If the Cl atom formed is not then ionized by
collision with an electron , it simply gets lost in the
machine – neither accelerated nor deflected.
 The Cl+ ions will pass through the machine and will
give lines at 35 and 37, depending on the isotope,
with the m/z = 35 line 3 times taller than the 37
line. This is what you would expect.
 It will also record lines for unfragmented Cl2+ ions.
 There are three different possible masses for a
Cl2+ ion depending on what combination of 35 Cl and
37 Cl atoms it contains.
 What could the masses be?
 The masses could be:
35 + 35 = 70
35 + 37 = 72
37 + 37 = 74
 So in addition to the lines at 35 and 37, there will
also be lines at 70, 72 and 74
 You should be able to work out the relative height
ratio to be 9:6:1 but you cannot make predictions
about the relative heights of the 35/37 compared
to those at 70/72/74. That depends on what
proportion of the molecular ions break up into
fragments.
Problems on isotopes:
 1) Calculate the relative atomic mass of silicon, given:
Relative isotopic mass
Relative abundance
28
100
29
5.10
30
3.36
 2) Calculate the relative atomic mass of gallium given
the percentage abundances:
 3) Bromine has two isotopes,
69Ga
79Br
60.2%, 71Ga 39.8%
and 81Br. At what
values of m/e would you expect to find lines in the mass
spectrum of bromine, Br2? (assume that only 1+ ions are
formed),

Describe the electromagnetic spectrum

Distinguish between a continuous spectrum and a line
spectrum

Explain how the lines in the emission spectrum of
hydrogen are related to electron energy levels.

Deduce the electron arrangement for atoms and ions
up to Z = 20
The Electromagnetic
spectrum
 Electromagnetic waves can travel through space
and, depending on the wavelength, also through
matter.
 The velocity of travel, c, is related to its
wavelength, λ and its frequency, f.
 Velocity is measured in m s-1, wavelength in m and
frequency in s-1 so it is easy to remember the
relationship between them.
c (m s-1) = λ (m) x f (s-1)
The electromagnetic spectrum
Although all e-m waves travel at the same speed, their
wavelength [] and frequency [ƒ] can be different.
Waves that cook food.
Waves that cause
sun-tans.
The properties, dangers and uses of e-m waves depends
on the wavelength [].
 Electromagnetic radiation is a form of energy.
 The smaller the wavelength and thus the higher
the frequency, the more energy the wave
possesses.
 Electromagnetic waves have a wide range of
wavelengths ranging from low energy radio waves
to high energy gamma (γ), radiation.
 As you have seen visible light occupies a very
narrow part of the spectrum.
The Atom
 It was the Greek philosopher Democritus who first
considered the idea that matter is made up of
particles in about 400BC
 The idea was not accepted because there was no
experimental evidence for it.
 John Dalton revived the discussion around 1801 and
compiled experimental evidence which convinced
people.
 Around the year 1900, physicists began to find
evidence that atoms are made up of smaller particles
Atomic spectra
 If sunlight or light from an electric light bulb is formed
into a beam by a slit and passed through a prism on to a
screen, a rainbow of separated colours is seen.
 The spectrum of colours is composed of visible light of all
wavelengths and is called a continuous spectrum
 The colours Red, Orange, Yellow, Green, Blue, Indigo and
Violet make up the colours in the visible part of what is
known as the electromagnetic spectrum.
 Each colour is a wavelength which represents a
particular amount of energy.
 If atoms and molecules are heated to sufficiently high
temperatures, they emit light of certain wavelengths.
 The observed spectrum is called an atomic emission
spectrum or line spectrum.
 All substances give emission spectra when they are
excited in some way, by the passage of an electric
discharge or by a flame.
 The atomic emission spectra of elements are in the
visible and ultraviolet regions of the spectrum.
 When sodium or a sodium compound is put into a
flame, it colours the flame yellow.
 A tube of hydrogen gas which has been excited by an
electric discharge glows a reddish-pink colour.
The Hydrogen Spectrum
 Viewed through a spectrometer, the emission
spectrum of hydrogen is seen to be number of
separate sets of lines or series of lines.
 These series of lines are named after their
discoverers
The Balmer series of
hydrogen
 The Balmer series is in the visible part of the spectrum.
 In each series, the intervals between the frequencies of
the lines become smaller and smaller towards the high
frequency end of the spectrum until the lines run
together or converge to form a continuum of light.
 Why do atomic spectra consist of discrete (separate)
lines?
 Why do atoms absorb or emit light of certain
frequencies?
 Why do the spectral lines converge to form a
continuum?
 The Rutherford picture of the atom offers no
explanation.
 The theories continued………
Niels Bohr
 In 1913, Niels Bohr (1885-1962)
put forward his picture of the
atom to answer these questions.
 Bohr referred to Max Plank’s
recently developed quantum
theory, according to which energy
can be absorbed or emitted in
certain amounts, like separate
packets of energy , called quanta.
The Bohr model:
 Bohr Suggested
 An electron moving in an orbit can have only certain
amounts of energy, not an infinite number of values: its
energy is quantised
 The energy that an electron needs in order to move in
a particular orbit depends on the radius of the orbit.
An electron in an orbit distant from the nucleus
requires higher energy than an electron in an orbit
near the nucleus.
 If the energy of the electron is quantised, the
radius of the orbit also must be quantised. There is
a restricted number of orbits with certain radii,
not an infinite number of orbits.
 An electron moving in one of these orbits does not
emit energy. In order to move to an orbit farther
away from the nucleus, the electron must absorb
energy to do work against the attraction of the
nucleus. If an atom absorbs a photon (a quantum of
light energy), it can promote an electron from an
inner orbit to an outer orbit.

For an electron to move from an orbit of energy E1
to one of energy E2, the light absorbed must have a
frequency given by Planck’s equation:
hv = E2 – E1 where v = frequency, h = Planck’s constant
 The emission spectrum
arises when electrons
which have been excited
(raised to orbits of high
energy) drop back to
orbits of lower energy.
They emit energy as light
with a frequency given by
Planck’s equation
 Bohr assigned quantum numbers to the orbits. He gave
the orbit of lowest energy (nearest to the nucleus)
the quantum number 1 an electron in this orbit is in its
ground state. The next energy level has quantum
number 2 and so on.
 If the electron receives enough energy to remove it
from the attraction of the nucleus completely, the
atom is ionised.
To summarise…
 When energy is supplied to an atom electrons are
excited (gain energy) from their lowest state to n
excited state.
 Electrons can only exist in certain fixed energy levels.
 When electrons drop from a higher level to a lower
level they emit energy.
 This energy corresponds to a particular wavelength
and shows up as a line in the spectrum.
When electrons return
to the first level (n = 1)
the series of lines
occurs in the UV region
as this involves the
largest energy change
When an electron is
‘excited ‘ (gains energy)
it can move from a lower
energy to a higher
energy level
n=4
n=3
n=2
When it returns to
it’s ‘ground state’ it
releases the energy
it absorbed as a
specific wavelength
n=1
Each shell of an atom
represents an energy
level, these are shown
by n = 1, n = 2, n = 3 to
n = infinity
The visible region
spectrum is formed by
electrons dropping back
to the n = 2 level
HEAT
The first series in the
infra red is due to
electrons falling to the
n = 3 level.
Development of ideas
 The idea of the atom is still developing but what you
must appreciate is that these are theories backed up
by scientific evidence, but evidence is not proof!!!!!
 Bohrs model is an accepted model of the atom and is
used as a building block for other ideas and theories,
it is by no means the end of the story.
Using Bohr’s model
 We can use Bohr’s model to help illustrate bonding in
metallic, ionic and covalent compounds.
 For this we need to be able to show electron
arrangements.
 Can you remember how many electrons can fit into the
first shell?
 How many can fit into the second?
 How do you know this??????????
 First we need to look at ionisation energies….
Ionisation energies
 When an electron receives enough energy to remove it
from the attraction of the nucleus completely, the
atom is ionised.
 The ionisation energy of an element is defined as:
‘The energy required to remove one mole of electrons
from one mole of atoms in the gas phase to form one mole
of cations in the gas phase’
A(g)
A+(g) + e-
Patterns of ionisation
energies in the periodic table
 Ionisation energy values can be found from emission
spectra, and the values can be found in any chemical
book of data.
 Look up the first ionisation
energies for each of the first
20 elements in the periodic
table and draw (as accurately
as is possible) a graph of
atomic number (x axis) versus
first ionisation energy (y axis)
2
8
8
Explaining Ionisation Energies
 A graph of first ionization energies plotted against
atomic number shows a repeating pattern.
 It can be seen that the highest value is for Helium, an
atom that contains two protons and two electrons.
 The two electrons are in the lowest level (n = 1) and
are held tightly by the two protons.
 For lithium it is relatively easy to remove an electron.
Can you suggest a reason for this?
 It suggests that the third electron in lithium is in a
higher energy level than the first two.
 The value then generally increases until element 10,
neon, is reached before it drops sharply for sodium.
 This graph provides evidence that the levels can
contain different numbers of electrons before they
become full.
Relating Ionisation Energies to
Atomic Structure
n=1
n=2
n=3