ARRANGEMENT OF
ELECTRONS IN THE ATOM
• Identify the first 20 elements in
terms of the number of outer
electrons.
• Evaluate Bohr’s study of the
spectra
WHERE ARE THE ELECTRONS?
Ernest Rutherford-electrons revolving in orbit.
defied the law of physics-electrons spiralling around the nucleus would lose
energy and spiral into the nucleus (hitting the atom).
Niel’s Bohr-hypothesised a line spectra to show the electron arrangement.
(white light broken up into a continuous spectrum)
Light from a hydrogen discharge tube created narrow lines instead of a
continuous spectra when passed through a prism.
Hydrogen discharge tube is a glass tube filled with H at low pressure and an
electric current is passed through.
EMISSION LINE SPECTRUM
Light is only emitted when an electric current is passed through hydrogen.
When hydrogen was replaced by other elements, different spectra resulted.
(fingerprint spectrum)
An emission spectrometer is used to analyse steel for the presence of different
elements.
Examples of element colours in everyday life.
Bohr’s atom
1.
Electrons
travel in
orbits around
the nucleus.
2)
Electrons have a fixed amount of
energy in any one orbit so the orbits
are called energy levels. Each energy
level has a definite value. The fixed
amount of energy is called a quantum
of energy.
3)
Electrons don’t gain or lose energy if
they are in one particular energy level.
4)
When an atom absorbs energy,
electrons jump from a lower energy
level to a higher energy level. (higher
the level the less stable)- further from
nucleus
5)
Energy is lost when electrons go from
higher to lower levels but electrons can
only fall to certain energy levels. (only a
fixed amount of light energy is given
off)
N=1
N=2
What do you remember? KWL

Today we will….

Define energy level, ground state and
excited state.
Homework

P24 3.2

Bohr’s study of spectra

To be continued.

Ground state: where electrons
in an atom occupy the lowest
available energy levels where
they have energies of fixed
value.

Emission spectra:

Absorption spectra:

Energy levels are represented by the
letter N. The lowest energy level is n=1

Atoms normally exist in ground state i.e
the electrons will have the lowest
amount of energy possible and occupy
the lowest energy orbitals. N=1
N=1
N=2

When a fixed amount of energy
is absorbed, an electron will
jump from n=1 to n=2 i.e a
higher energy level.

An energy level is defined as a
fixed energy value that an
electron in an atom may have.
Ground state

Electrons:

Occupy lowest energy level.

Have a fixed energy value.

Need a specific amount of
energy to jump from lower to
higher energy levels to an
excited state.

The energy level an electron
moves to depends on the atom.
Excited state

Energy absorbed=difference in
energy between the lower level
(ground state) and the higher level
(excited state).

Electrons in the excited state are
unstable, the fall back down after a
certain time period.

The extra energy is released in the
form of a photon of light with a
definite amount of energy due to
falling to a definite energy level. Light
of a definite frequency is given off.

Frequency of light: energy
difference

E2-E1=hf

E2 higher energy level

E1 lower energy level

h= Planck’s constant

f= Frequency of light emitted

The energy difference s
proportional to the frequency of
light emitted.

This frequency appears on an
emission spectrum as a line of a
particular colour.

Each definite amount of energy
gives rise to a line in the emission
spectrum. Proves energy levels
exist!

Each element has a unique
emission line spectrum because
each element has a definite
number of electrons and has its
own arrangement of these
electrons in energy levels. This
is called electron configuration.
True/False
Was Bohr Correct?
1.
After he realised the coloured line on
the emission spectrum of hydrogen
had its own frequency…
2.
He Calculated mathematically the
energy corresponding to n=1 n=2 etc.
3.
Calculated the energy an electron
would lose when falling from higher
energy level to lower energy levels.
1.
Used E2-E1=hf to calculate the
frequency and wavelength of the
coloured lines corresponding to the
transitions. He calculated frequency
from the above equation and
velocity of light = frequency X
wavelength allowed him to calculate
wavelength
Safety in the lab!

Barium chloride in is harmful by ingestion and
inhalation.

Copper(II) chloride is toxic if swallowed, and
is an eye and skin irritant.

Lithium chloride n is harmful by ingestion, and
is a severe eye and skin irritant.

Potassium chloride is an eye irritant.
Predict observe explain!!

Copper sulfate: Blue-green

Lithium chloride

Lithium carbonate: Deep red

Sodium chloride

Sodium sulfate: Yellow

Potassium chloride

Strontium nitrate: Red

Barium chloride

Barium nitrate: Yellow-green

Strontium chloride

Potassium sulfate: Lilac

Copper(II) chloride

https://www.youtube.com/watch?v=Rf3
fKLLPpRY
Atomic Absorption spectroscopy

Scientists have discovered that atoms
can absorb light.

When white light is passed through a
sample of a gaseous element, the light
that comes out has certain wavelengths
missing. i.e dark lines in the spectrum

This is called an atomic absorption
spectrum.

The dark lines represent missing
wavelengths and tell us what
elements are present as the dark
lines represent the electrons that
have absorbed the light.

The wavelengths that are
absorbed correspond exactly to
the wavelengths detected in the
emission spectrum.

What does that tell us?

You need:

White light

Gaseous sample

Prism

AAS is an important tool
used by scientists to
determine certain
elements and their
concentrations.