1. 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.
3. 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.
8. 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.
10. Emission Spectrum
• Shows that atoms can emit
only specific energies
(discrete wavelengths,
discrete frequencies)
11. hypothesis: if atoms emit only
discrete wavelengths, maybe
atoms can have only discrete
energies
12.
13.
14. A turtle sitting on a
staircase can take on only
certain discrete energies
15. 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)
17. bottom step is called the
ground state
higher steps are called
excited states
18. 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
19. Flame Test
• Flame Test
The 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
20. 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.
22. 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
23. 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.
25. 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
30. Bohr Diagrams
To 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.
34. 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.
35.
36. Heisenberg
• We cannot know both the position and
speed of an electron
• Therefore we cannot describe how an
electron moves in an atom
47. Main levels AND THE
NUMBER OF ELECTRONS
•
•
•
•
1 = 2e
2 = 8e
3 = 18e
4 = 32e
48. 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
56. 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)
