Download Overview Atomic Structure

Atoms: very small
Comparing the size of
a helium atom to a
tennis ball...
... is like comparing a
tennis ball to the Earth.
The ancient Greeks theorised that
everything was made of fire, earth,
water and air. Only after over 1600
years did scientists realise matter was
composed of different elements.
In 1805 John Dalton refined a vague
idea that matter was indivisible.
In 1805, John Dalton proposed explanations
for some observations he had made during
experiments. He developed an atomic theory.
1. Elements are made of extremely small particles called atoms.
2. Atoms of a given element are identical in size, mass, and other
properties; atoms of different elements differ in size, mass, and
other properties.
3. Atoms cannot be subdivided, created, or destroyed.
4. Atoms of different elements combine in simple whole-number
ratios to form chemical compounds.
5. In chemical reactions, atoms are combined, separated, or
rearranged.
Dalton predicted
that each element
was made of a type
of atom and he
created symbols for
them.
He also put them
together to make
compounds.
After he discovered the
electron in 1897, J J
Thomson suggested in
1904 that electrons are
embedded in a blob of
positive charge.
It was nicknamed the
“plum pudding model”.
Rutherford’s gold foil experiment
In 1909, Ernest Rutherford bombarded a thin gold foil
with alpha particles.
He thought they would go
through, but be “bent” by
the positive charge in the
atoms
Instead, most went straight
through and a few bounced
straight back.
Rutherford had discovered the positivelycharged nucleus and disproved Thomson’s
theory. But where were the electrons?
Niels Bohr suggested they orbited the
nucleus in “shells” around it, like planets
orbiting the sun.
Quantum Theory, 1920s
The modern
atomic model
Electrons exist in
a “cloud” around
the nucleus.
They can be thought of
as waves, particles or
mathematical functions
How big is a nucleus?
The helium atom is not in the right
proportions.
The three subatomic particles are
wrongly enormous in comparison to
the atom’s radius.
Most of the atom is empty space.
The nucleus is extremely small.
But we draw it like this anyway, for obvious reasons.
Ease of drawing.
The number of protons is what defines an element
A is the mass number:
protons + neutrons
Z is the atomic number,
or proton number.
4
2
He Z = 2
4
Helium
He
2
No. of protons, Z
Number of
protons
=
Electrons
2
Protons
2
Neutrons
2
Number of
electrons
Subatomic particles in more detail
Subatomic
particle
Relative
charge
Relative
mass
Common
depiction
Proton
+1
1
+
Neutron
0
1
-1
1
____
1836
Electron
Actual mass of electron, me = 9.11 × 10−31 kg
-
Lithium
Mass number, A
7
Li
3
Atomic number, Z
Number of
protons (+)
=
Number of
electrons (-)
The charges must be
balanced in a neutral atom.
Lithium
Electrons
3
Protons
3
Neutrons
4
Number of protons does not always
equal the number of neutrons.
Isotopes
The number of protons is the “soul” of an atom.
Most elements have different isotopes: atoms of the
same element with different numbers of neutrons.
Chlorine (Z = 17) has 2 naturally-occurring isotopes:
37Cl
Neutrons
35Cl
37Cl
18
20
% abundance 75.53 24.47
35Cl
Thus the Ar of chlorine is:
(  )+(  )
=
and
Isotopes
Write the symbol for each of the following isotopes
1. Xenon-133
2. Oxygen-15
3. Iodine-125
4. Phosphorus with 18 neutrons
5. Cobalt with number of neutrons = Z + 3
6. The three isotopes of hydrogen
7. The most abundant isotope of carbon
Relative atomic mass
Symbol: Ar
Since elements have different isotopes, the mass
stated in the periodic table is a weighted average.
It is stated relative to a particular “standard”. It
has a mass of 1 and is equal to 1/12 of the mass of
an atom of 12C.
Relative atomic mass = average mass of the
naturally-occurring isotopes of an element relative
to C-12.
Physical properties of isotopes
10B
and 11B have different atomic masses.
This affects their properties:
• Density
• Boiling point
• Melting point
• Deflection by a
magnetic field
• Rate of diffusion
Heavier things
require more
energy to
move!
This means it is possible to separate isotopes.
Useful isotopes
14C
“Carbon dating” of organic matter up
to about 60,000 years old. In use since
1949.
131I
Radiotherapy, particularly to test
functioning of thyroid gland, or to
treat cancer of the thyroid gland,
medical tracer
125I
Used to treat brain and prostate
cancer.
60Co
Radiotherapy, sterilisation of medical
equipment, industrial radiography,
thickness and levelling tools, medical
tracer
t½ : 5730 years
β- decay to 14N
t½ : 8.07 days
β- & γ decay to 131Xe
t½ : 60 days
γ decay to 125Te
t½ : 5.27 years
β- decay to 60Ni
Electronic Structure
The first 20 elements have their electrons
arranged in a simple manner:
The first energy level can take up to 2 electrons.
The second energy level can take up to 8 electrons.
The third energy level can take up to 8 electrons.
After calcium (Ca) things start to get more
complicated. Studying the transition metals
requires a new method of arranging electrons on
top of the first 20 elements (orbital theory).
Electronic Structure
All the Group 1 elements have 1 electron in the
outermost energy level.
2,1
Li
Lithium
Na
Sodium
K
Potassium
Rb
Rubidium
Cs
Caesium
2,8,1
2,8,8,1
The group number tells you
how many electrons the atom
has in its outer energy level.
Magic!
How many protons, electrons and neutrons do these have?
Atom
Na
Rh
phosphorus
The last of the halogens
Xe
The only liquid non-metal
Li+
FCarbon-14 (14C)
4
He
2
A helium atom
P
N
2 protons
2 electrons
2 neutrons
E
How many protons, electrons and neutrons do these have?
Atom
Na
Rh
phosphorus
The last of the halogens
Xe
The only liquid non-metal
Li+
FCarbon-14 (14C)
4
He
2
A helium atom
P
11
45
15
85
54
35
3
9
6
N
12
58
16
125
77
45
4
10
8
2 protons
2 electrons
2 neutrons
E
11
45
15
85
54
35
2
10
6
Fill in the table
Species
238
92
Protons
Electron
Neutrons Electrons configuration
U
X
40Ca2+
35
Cl
Br-81
Fluoride ion
X
10
Carbon-14
Mg ion in MgO
133
Xe
12
X
Fill in the table
Species
238
Protons
Electron
Neutrons Electrons configuration
U
92
146
92
X
40Ca2+
20
20
18
2,8,8
Cl
17
18
17
2,8,7
Br-81
35
46
35
2,8,8,17
Fluoride ion
9
10
10
2,8
Carbon-14
6
8
6
2,4
Mg ion in MgO
12
12
10
2,8
54
79
54
X
92
35
133
Xe
Exciting electrons
Electrons can jump from one energy level to another.
Absorption of energy
Emission of energy
Electron jumps from
n=3 to n=4
Electron jumps from
n=4 to n=3
Exciting electrons – so what?
Well, without these transitions the
world would be boring and colourless!
Yes, it’s true! Electron transitions in
atoms are responsible for colour!!
We can disturb these poor
electrons by heating atoms up.
When they go back to their
original arrangement, energy is
emitted in the form of light.
The light pattern is
called a
spectrum
In the 1850s, Robert
Bunsen and Gustav
Kirchhoff invent
spectroscopes.
They use them to study
the spectra of blood and
other substances for
criminal investigations.
Just like a fingerprint,
each substance has its
own spectrum
Lockyer
1868
Discovered helium while
looking through a spectroscope
at the sun during an eclipse
The spectroscope
A simple device:
1. A slit allows
light in.
2. The light is
split by a
prism into a
spectrum.
3. The resulting
spectrum is
projected onto
a screen.
Not the whole thing
The continuous spectrum (“rainbow”) is the
result of white light. If you start off with only a
fraction of white light (i.e. a colour) then you
won’t get a “rainbow”. Instead, you’ll get only
certain bits of it. This is a line spectrum.
Hydrogen
Iron
Light source
Spectroscope
Projection
Continuous
spectrum
White light
Slit
Heated gas
Prism
Emission
spectrum
Cold gas
Absorption
spectrum
White light
Light source
Spectroscope
Projection
Continuous
spectrum
White light
Slit
Heated gas
Prism
Emission
spectrum
Cold gas
Absorption
spectrum
White light
Absorption + emission spectra
Hydrogen
The Balmer series
In the visible light spectrum (about 380 to 750nm)
From n > 2 to n = 2
Ultraviolet
visible region
Infrared
656 nm
n = 3 to n = 2
486 nm 434 nm 410 nm
n = 4 to
n=2
n = 5 to
n=2
n = 6 to
n=2
Balmer series
CONVERGENCE
Infrared region
Visible region
Ultraviolet region
Atomic Structure (HL)
WRONG
Electrons don’t orbit the nucleus.
They don’t spin around it along a
set path.
Bohr model (1913)
Also rather wrong. It’s
not that simple.
The electrons don’t
spend their time a
specific distance from
the nucleus.
But this diagram is partly useful if one
considers the “shells” not to be orbits
but energy levels.
A better model
The electron can be considered
to be a particle, but also a wave
or wave packet.
This means there is a probability associated with
the location of the electron – it’s not in a
definite place.
Heisenberg’s uncertainty principle
“You cannot determine both the position and
momentum of an electron at the same time.”
Electron density
How do we describe where the electron is, then?
Electron density is a measure of the probability of
finding the electron in a particular place.
Covalent bonds
are areas of
high electron
density
Pentacene, C22H14
How do you take a
photograph of a molecule?
Atomic Force
Microscopy (AFM)
Very fine tip, one atom
thick. Carbon
monoxide molecule
attached to tip to
make it sharper and
produce better images
The tip oscillates above the sample, measuring the
forces between them
February 2012:
Image of charge distribution
around a naphthalocyanine
molecule
September 2012 – New images of a similar molecule showing
different bond lengths
s orbitals
d orbitals
p orbitals
z
The long form of the periodic table
Normally rather inconvenient, so the f-block is
detached out of the way.
Note that helium is actually in the s-block.
s
f
d
p
The f block and its orbitals
Starts to be filled after 6s. 7 different f orbitals,
meaning up to 14 electrons go in each f sub-shell.
Relative Energies of Sub-shells
This explains the order or the s, p, d
and f blocks in the periodic table
Rules for arranging electrons
Aufbau (“building up”) principle
“Electrons enter the lowest available
energy level.”
Hund’s rule of “maximum multiplicity”
“When in orbitals of equal energy,
electrons will try to remain unpaired.”
This is because electrons repel each other
Pauli exclusion principle
“No two electrons can have the same
four quantum numbers.”
i.e. A full orbital must contain oppositespin electrons.
n
>0
principal quantum number (main shell)
l
0 to n-1
azimuthal q.n. (sub-shell: s, p, d or f)
m
-l to +l
magnetic q.n. (orbital within sub-shell)
s
½
spin q.n. (up or down)
3p
E
•n=3
3s
2p
2s
Example: Silicon (Z=14)
The last electron has:
• l = 1 (i.e. p orbital)
• m = 0 (py - middle p)
• s = + ½ (spin up)
1s
No other electron in this atom has the
same set of four quantum numbers
Matching orbitals to quantum numbers
Evidence: ionisation energies
Ionisation Energy is the energy required to remove
a mole of electrons from a mole of gaseous atoms.
It is determined largely by the Effective Nuclear
Charge
This is the attractive force from the nucleus, minus
the effect of inner electrons counteracting it.
I.E. is also affected by electron-electron repulsion
within the same main energy level.
Successive ionisation energies increase due to decreased
electron-electron repulsion, and increasing ENC.
First n=2
electron
First n=1
electron
Variation in 1st I.E. with atomic number
Why the drop between Mg-Al?
Why the drop between P-S?