1. SCH4C UNIT 1: MATTTER AND QUALITATIIVE ANALYSIS Atomic Theory 2
Development of the Bohr Model of the Atom

2. Learning Objectives
By the end of today’s lesson, you should be able to:
Describe the Bohr model of the atom
Explain the relationship between atomic structure and line Emission and Absorption spectra
Explain how line spectra can be used to identify the element.

3. Summary of the Subatomic Particles
Location
Charge
Mass (amu)
Proton
Nucleus +1
1.0073
Neutron
1.0087
Electron
Outside Nucleus -1

4. Problems with the Planetary Model
Electrons orbiting the nucleus should lose energy and emit light. This loss of energy would cause the electrons to spiral into the nucleus, resulting in the collapse of the atom.
Excited electrons should emit a continuous spectrum of white light when they are excited. Instead, the emission spectrum of elements are all unique.

5. The Wave Theory of Light
Light is a form of electromagnetic radiation. All electromagnetic radiation is made up of electric and magnetic fields. These fields oscillate in a wave pattern as electromagnetic radiation moves through space. Electromagnetic radiation travels at a constant speed in a vacuum (“the speed of light”), but the wavelength and frequency can change.

6. Basic wave terms: Wavelength (): the distance between two crests of a wave. Frequency (f): the number of waves per second. Speed of light (c) = 300,000,000 m/s
short wavelength
higher frequency
long wavelength
lower frequency
direction of movement

7. The Electromagnetic Spectrum
All colours and types of electromagnetic radiation have different wavelengths and frequencies. Note: 1 nanometer (nm) = 1 x 10-9 m The “visible” part of the spectrum is between 390 nm and 750 nm.

8. Note that the wavelength and frequency or inversely proportional:
Name of Radiation
Wavelength (m)
Frequency (c/s or Hz)
Radio Waves
Microwaves
Radar Waves
Infrared Light
Visible Light *
UV Light
X-Rays
Gamma Rays
102
Increasing
Wavelength
10-14
106
Frequency
1022
Colour
Wavelength (m)
Frequency (c/s or Hz)
Red
Orange
Yellow
Green
Blue
Indigo
Violet
6.5 x 10-7
4.1 x 10-7
4.6 x 1014
7.3 x 1014
Note that the wavelength and frequency or inversely proportional:
If  , then f 

9. The Line Spectrum of the Elements
When a pure element is excited, it will produce a unique colour. This colour is produced by a line spectrum.
Line spectrum: a discontinuous spectrum produced when light is given off by an element.

10. Line Spectra and Elements
The line spectrum of an element is unique and can be used to identify the element.

11. The Bohr Model of the Atom (1913)
Bohr developed a model of the atom which explained the line spectrum and why an atom doesn't just collapse: Postulate 1: Electrons can only move in certain fixed orbits. Each orbit corresponds to a specific energy level and an electron can move within an orbit without losing any energy. Postulate 2: An electron can only move from one orbit (or energy level) to another when it gains or loses energy.

12. The Line Spectrum Explained
In a discharge tube, the electrons become temporarily excited and thus move to orbits or shells that are farther from the nucleus. These transitions are only temporary, though, so the electrons return to their lowest (or ground) states. When they do so, they give off (emit) light.

13. Why Only Four Visible Lines?
Since electrons can only undergo transitions between certain specific energy levels, only certain quantities (quantums) of energy are given off. Since the colour of light corresponds to the energy of possessed by a quanta (or photon), only certain coloured lines are observed. Animation of the Bohr Model of the Atom (Emission) Emission and Absorption Spectra (Youtube Video)

14. Significance of the Bohr Model
Bohr’s model explains the chemical and behaviour of elements is related to the filling of Bohr’s energy levels (2, 8, 8,18 etc.).