Atomic Structure

Still study the sheet with the details!!!

Scientific Model

A simplified representation of reality used to understand, explain or predict phenomena

Continuous Spectrum

An unbroken band of colours produced when white light is separated into its component wavelengths.

Line Emission Spectrum

A series of separate coloured lines producced when light emitted by an excited element is separated into its component wavelengths.

Energy level

The fixed amount of energy of an electron in an atom.

Quantised energy

Energy that can have only particiular fixed values; values between these allowed amounts are not possible.

Ground state

The state in which the electrons occupy the lowest available energy levels.

Excited State

The state in which one or more electrons occupy higher energy levels than those available in the ground state.

Photon

A quantum, or discrete packet, of electromagnetic energy.

Energy Sublevel

A subdivision of a main energy level.

The different sublevels are…

… s, p, d and f

Heisenburg’s Uncertainty Principle

It is impossible to know both the exact position and the exact momentum of an electron at the same time.

Orbital

A region of space around the nucleus within which there is a high probability of finding an electron.

Why might a scientific model need to be changed?

When new experimental evidence cannot be explained by the existing model.

Particulate Nature of Matter

Proposed by Dalton.

The idea that all substances are made of tiny, indivisible particles.

Who proposed the Plum Pudding model?

J.J Thomson.

What was the Plum Pudding model?

The proposition that the atom was a sphere of positive charge with negatively charged electrons embedded throughout it.

The Plum Pudding model explained why the atom is electronically neutral, but…

it did not include a nucleus and was later contradicted by Rutherford’s evidence.

1st assumption of the Nuclear Model.

The atom is very dense, very small and positively charged.

2nd assumption of the Nuclear Model.

The atom is mostly empty space.

3rd assumption of the Nuclear Model.

The nucleus takes up most of the atoms mass.

4th assumption of the Nuclear Model.

The electrons are located in the space surrounding the nucleus.

1st limitation of the Nuclear Model.

Didn’t show the…

Did not show the arrangement of electrons.

2nd limitation of the Nuclear Model.

…why electrons didn’t…

Didn’t explain why electrons didn’t lose energy and spiral into the nucleus.

3rd limitation of the Nuclear Model.

It gave no clear account of…

It gave no clear account of electron behaviour.

Rutherford’s gold-foil experiment.

An experiment conducted by Rutherford in which positively charged alpha particles were directed at a thin sheet of gold foil.

Rutherford’s Gold foil experiment

1. Observation → 2. What it showed → 3. Conclusion 1

1. Most particles passed straight through

2. They met very little matter

3. Most of the atom is empty space

Rutherford’s Gold foil experiment

1. Observation → 2. What it showed → 3. Conclusion 2

1. Some particles were deflected.

2. They were repelled by positive charge.

3. Positive charge is concentrated in a small nucleus

Rutherford’s Gold foil experiment

1. Observation → 2. What it showed → 3. Conclusion 3

1. Very few particles bounced back

1. Very few particles bounced back

2. They met a very dense region

3. The nucleus is very dense and contains most of the mass.

Spectrometer

A tool used to examine line spectra. When an electric current is passed through a gas in a discharge tube, the gas emits light.

Continous Spectrum vs Line Spectrum

Continous spectrum contains every visible wavelength

Line spectrum contains only particular wavelengths

Why is each elements spectrum unique?

Each element has its own arrangement of electrons + set of allowed energy levels.

The energy gaps between these levels are therefore different.

Electron transitions in each element release a unique set of photon energies and frequencies, producing a characteristic line emission spectrum.

1st assumption of the Bohr Model.

Electrons move…

Electrons move in fixed orbits around the nucleus.

2nd assumption of the Bohr Model.

Each orbit corresponds to a particular energy level.

3rd assumption of the Bohr Model.

Electrons can only have certain fixed, quanties amounts of energy.

4th assumption of the Bohr Model.

Energy levels are labelled…

Energy levels are labelled n=1, n=2, n=3 and so on.

1st limitation of the Bohr Model.

Didn’t account for wave-particle duality

2nd limitation of the Bohr Model.

Worked well for hydrogen, but did not work so well for multi-electron atoms.

3rd limitation of the Bohr Model.

Didn’t account…

Didn’t account for sublevels

Electron transition step 1

An atom absorbs energy when it is heated or when electricity is passed through it. It becomes excited.

Electron transition step 2

An electron can move to a higher energy level only if it absorbs an amount of energy equal to the gap between the two levels.

Electron transition step 3

The excited state is unstable. The electron later loses energy and returns back to the lower energy level.

Electron transition step 4

The electron releases the energy difference as a photon of electromagnetic radiation.

Photon frequency

Because the allowed energy levels are fixed, electrons can release only particular amounts of energy.

Each photon therefore has a definite frequency and wavelength.

A photon in the visible region appears as a particular coloured line in the emission spectrum.

Eₘ - Eₙ = hf

Eₘ = energy of the higher energy level

Eₙ = energy of the lower energy level

h = Planck’s constant

f = frequency of the emitted photon

Higher photon frequency means…

Higher photon energy and a shorter wavelength.

Lower photon frequency means…

Lower photon energy and a longer wavelength.

Balmer series

The visible line spectrum

Excited electrons fall to n=2

Hydrogen Emission Spectrum

n=6 to n=2: Violet line

n=5 to n=2: Blue-violet line

n=4 to n=2: Blue-green line

n=3 to n=2: Red line

UV Lyman Series

Transitions ending at n=1

Infrared Paschen Series

Transitions ending at n=3

Potassium

Lilac

Sodium

Yellow-orange

Copper

Blue-green

Lithium

Crimson

Barium

Apple green

Strontium

Red

Method for the flame test

Soak separate wooden splints in water.

Crush a small sample of each salt using clean equipment.

Dip a damp splint into the salt and hold it in the blue Bunsen flame.

Record the flame colour and use a different splint for each salt.

Evidence for energy sublevels

Improved spectrometers showed that some lines that had appeared to be single lines were actually several very closely spaced lines.

Sodium’s yellow line is a doublet.

The small energy differences could not be explained by transitions between completely different main energy levels.

Scientists concluded that most main energy levels are divided into smaller, closely spaced energy sublevels.

1st assumption of the Orbital Model

Electrons do not follow fixed paths

2nd assumption of the Orbital Model

Electron behaviour…

Electron behaviour is described using probability.

3rd assumption of the Orbital Model

Electrons show…

Electrons show wave-particle duality

1st limitation of the Orbital Model

2D models cannot fully display the 3D structure

2nd limitation of the Orbital Model

The exact path and poisition of an individual electron is not provided.

3rd limitation of the Orbital Model

Orbital boundaries and their shapes are simplified probability representations.

Evidence leading to the orbital model

• Electrons exhibit wave-particle duality; they show both particle-like and wave-like behaviour.

• The splitting of spectral lines showed that main energy levels contain smaller sublevels.