Learning Outcomes

• Emission and absorption spectra of the hydrogen atom .

• Balmer series in the emission spectrum as an example.

• Line spectra as evidence for energy levels.• Energy sub-levels. • Viewing of emission spectra of elements using a

spectroscope or a spectrometer.

Atomic structure

Spectra

Spectrscope

In a light spectroscope, light is focused into a thin beam of parallel rays by a lens, and then passed through a prism or diffraction grating that separates the light into a frequency spectrum.

Continuous Spectrum

Emission SpectraEmission Spectra

Continuous spectrum

A Spectrum in which all wavelengths are present between certain limits.

Emission Sprectrum

Emission spectrum

Spectrum lines

When light from an unknown source is analyzed in a spectroscope, the different patterns of bright lines in the spectrum reveal which elements emitted the light. Such a pattern is called an emission spectrum.

Absorption spectrum

Emission Spectrum

• Shows that atoms can emit only specific energies (discrete wavelengths, discrete frequencies)

hypothesis: if atoms emit only discrete wavelengths, maybe atoms can have only discrete

energies

A turtle sitting on a staircase can take on only certain discrete energies

energy is required to move the turtle (electron) up the steps (energy levels) (absorption) energy is released when the turtle (electron) moves down

the steps (energy levels) (emission)

energy staircase diagram for atomic hydrogen

bottom step is called the ground state

higher steps are called excited states

Balmer Series

• Balmer analysed the hydrogen spectrum and found that hydrogen emitted four bands of light within the visible spectrum:

• Wavelength (nm) Color• 656.2 red• 486.1 blue• 434.0 blue-violet• 410.1 violet

Flame Test

• Flame TestThe following metals emit certain colours of light when their atoms are excited.

• Metal      Colour• Sodium (Na) Yellow• Lithium (Li) Pink/Red• Potassium (K) Purple• Copper (Cu) Green• Calcium (Ca) Pink• Barium (Ba) Yellow/Orange• Strontium (Sr) Red/Orange

Learning Outcomes

• Energy levels in atoms.• Organisation of particles in atoms of

elements nos. 1–20 (numbers of electrons in each main energy level).

• Classification of the first twenty elements in the periodic table on the basis of the number of outer electrons.

Bohr

Bohr’s theory

• Electrons revolve around nucleus in orbits

• Electron in orbit has a fixed amount of energy

• Orbits called energy levels• If electron stays in level it neither

gains nor loses energy

Bohr

• Atom absorbs energy • Electron jumps to higher level• Atom unstable at higher levels. Electron falls back

to a lower level• Atom loses or emits energy of a particular

frequency.

quantisation

• Electrons can have only certain particular values of energy

EVIDENCE FOR ENERGY LEVELS

• In Hydrogen electron in lowest (n=1) level; ground state

• Energy given; electron jumps to higher level excited state

• Falls back and emits a definite amount of energy

• Energy appears as a line of a particular colour

colours

• Energy emitted depends on the jumps

• Different jumps emit different amounts of energy and hence different colours

Main energy levels (shells)

• Spectroscopic notation for shells . • N shell name

1 = K

2 = L

3 = M

4 = N

Bohr Diagram

Bohr DiagramsTo draw Bohr Diagrams:1.Draw the nucleus as a solid circle.2.Put the number of protons (atomic number) in the nucleus with the number of neutrons (atomic mass – atomic number) under it.3.Place the number of electrons (same as protons) in orbits around the nucleus by drawing circles around the nucleus.

Remember, 1st shell – 2 electrons, 2nd shell – 8 electrons, 3rd shell – 8 electrons, 4th shell – 18 electrons.

Valency & Groups

Valencies

Atomic structure 2

Learning Outcomes

• Energy sub-levels.• Heisenberg uncertainty principle.• Wave nature of the electron. (Non-mathematical treatment in• both cases.)• Atomic orbitals. Shapes of s and p orbitals.• Building up of electronic structure of the first 36 elements.• Electronic configurations of ions of s- and p-block elements only.• Arrangement of electrons in individual orbitals of p-block atoms.

Heisenberg

• We cannot know both the position and speed of an electron

• Therefore we cannot describe how an electron moves in an atom

Einstein

• . •

.

                                                       

       

De Broglie

• Matter has wave characteristics

2-slit expt..

                                                      

      

Expected Result if light and electrons are particles:

                                                           

        

Actual result for light and electrons – demonstrates their wavelike nature

:

                                                           

        

Electrons were both particles and waves

Same for all sub-atomic particles

Matter exists as particles and waves at the same time.

The electron as a wave

Orbital

• A region in space where the probability of finding an electron of a particular is high

Electrons moving

Electron paths

Main levels AND THE NUMBER OF ELECTRONS

• 1 = 2e• 2 = 8e• 3 = 18e• 4 = 32e

Sub-levels

• Each main level has sub-levels• 1has s sub-level only• 2 has s and p sub-levels• 3 has s,p and d sub-levels• 4 has s,p,d and f sub-levels• Energy of sub-levels spd

1s

2s

2p

3d

Electrons in sub-levels

• s = 2e• p = 6e• d = 10e• f = 14e

Sub-levels

• 1 = s(2e)• 2 = s(2e) + p(6e) = 8e• 3 = s(2e) + p(6e) + d(10e) = 18e

The "p" orbital is dumb belled shaped and each P sub level is made of three "p" orbitals (because the P sub level can hold 6 electrons and every orbital holds 2 electrons)

P-orbitals

P-orbitals

Electrons in orbitals

• S holds 2e• 3 p orbitals each holds only 2e• 5 d orbitals each holds only 2e

Pauli’s exclusion principle

• Orbital can only hold 2electrons and these electrons must have opposite spins

Pauli's exclusion principle

Aufbau principle

• Electrons fill levels in a specific order.

• 1s 2s 2p 3s 3p 4s 3d 4p

AUFBAU

Hunds rule

• When filling up the orbitals in a sublevel electrons fill then singly at first.

5 electrons

6 electrons Hund’s rule

Electron Configurations

• He, 2, helium : 1s2 • Ne, 10, neon: 1s2 2s2 2p6 • Ar, 18, argon : 1s2 2s2 2p6 3s2 3p6 • Kr, 36, krypton : 1s2 2s2 2p6 3s2 3p6 4s2 3d10

4p6

Exceptions to Electron configuration rules

• Cr • Half-filled orbitals give greater stability • 1s2 2s2 2p6 3s2 3p6 3d4 4s2 1s2 2s2 2p6 3s2 3p6 3d5

4s1

• Cu • Full 3d sub-level gives greater stability• 1s2 2s2 2p6 3s2 3p6 3d9 4s2 1s2 2s2 2p6 3s2 3p6 3d10

4s1

Electron Configurations (ions)

• F-, 10, Flouride: [1s2 2s2 2p6 ]-

• Cl-, 18, Chloride : [1s2 2s2 2p6 3s2 3p6]- • Na+, 10, Sodium ion: [1s2 2s2 2p6 ]+