1. Electrons in Atoms

2. The Evolution of Models of the Atom
We have previously described atoms using three different models.
As a group,
1) Build a replica of the model of the atom as described by the scientist your group is assigned.
2) Discuss in your group the experimental evidence that led to your model.

3. Dalton Dalton gave us the basis of the modern atomic theory.
- All matter are composed of tiny indivisible particles called atoms.
- Atoms of the same element are identical. Atoms of different elements are different.
- Atoms combine in simple, whole number ratios to form compounds.
- Atoms are separated, rearranged, and recombined during chemical changes.

4. Dalton’s Model of the Atom

5. J.J. THomson
J.J. Thomson discovered the electron, a negatively charged particle, through his cathode ray tube experiment.
This discovery resulted in a revised model of the atom called the plum pudding model.
- Atom was made of a positively charged material.
-Negatively charged electrons were distributed throughout the atom.

6. Thomson’s model of the atom

7. Rutherford
After Rutherford’s discovery of the positively charged nucleus through his gold foil experiment, another revision of the model of the atom occurred.
He proposed that the protons were concentrated in a small dense nucleus and the electrons were located in the large space surrounding the nucleus.

8. Rutherford’s model of the atom

9. Bohr
Bohr, a student of Rutherford, questioned how the electrons, being oppositely charged from the protons, did not fall into the nucleus due to the attraction that forms between opposite charges.
Bohr proposed that the negatively charged electrons must be moving in circular orbits around the nucleus, much like planets orbit the sun.
This model was referred to as the planetary model.

10. Bohr’s model of the atom

11. Evolution of the model of the atom

12. The atom and unanswered questions
Bohr’s model did not address the differences and similarities in chemical behavior among the various elements.
In the early 1900’s, chemists observed that certain elements emitted visible light when heated in a flame or exposed to other energy sources.
Analysis of the emitted light revealed that this behavior was related to the electron arrangement within its atoms.
In order to understand this relationship, we must first review the nature of light.

13. The wave nature of light
Visible light is a type of electromagnetic radiation-a form of energy that exhibits wavelike behavior as it travels through space.
Other forms of electromagnetic radiation include TV/radio waves, microwaves, infrared waves, ultraviolet waves, x-rays, and gamma rays.
All of these types of waves together make up the electromagnetic spectrum.

14. Electromagnetic spectrum

15. Characteristics of waves
The only differences among the different types of waves making up the electromagnetic spectrum are the frequencies and wavelengths.
1) wavelength (represented by the Greek letter lambda, λ) is the shortest distance between equivalent points on a continuous wave. Wavelength is usually expressed in meters or nanometers (1 nm = 1 x 10-9 m)
The wavelength of visible light is between 400 and 700 nm.

16. Wavelength

17. Characteristics of waves
2) frequency (represented by the Greek letter nu, ν) is the number of waves that pass a given point per second. Frequency is measured in hertz, Hz. One hertz equals one wave per second.

18. Frequency

19. Relationship between frequency and wavelength
What do you notice about the relationship between the frequency and the wavelength of a wave?
It is an inverse relationship: as the frequency increases the wavelength decreases.

20. Characteristics of waves
3) Amplitude of a wave is the wave’s height from the origin to a crest, or from the origin to a trough. Wavelength and frequency do not affect the amplitude of a wave.

21. Characteristics of waves
4) All electromagnetic waves, including visible light, travel at a speed of 3.00 x 108 m/s in a vacuum. This value is often referred to as the speed of light. Because this is such an important and universal value, it is given its own symbol, c.
The speed of light = wavelength x frequency.

22. Practice problem
Calculate the wavelength of the yellow light emitted by the sodium in the pickle if the frequency of the radiation is 5.10 x 1014 Hz.
c = νλ
3.00 x 108 = 5.10 x 1014 (λ)
λ = 5.88 x 10-7 m or 588 nm
The red-colored light given off by the compound held in the flame has a wavelength of x 10-7 m. What is the frequency of such red light?
ν = 4.62 x 1014 Hz

23. The particle nature of light
Sunlight, our most common form of light is continuous and covers all of the wavelengths in the visible spectrum.
It can be separated by passing the light through a prism.
This model cannot explain, however, why heated elements emit only certain frequencies of light at a given temperature.
Scientists realized that a new model or revision of the wave model of light was needed.

24. The quantum concept
When objects and elements are heated or exposed to energy, they emit glowing light of different colors.
These colors correspond to different frequencies and wavelengths.
A German scientist, Max Planck, studied this light phenomenon and came to a conclusion: matter can gain or lose energy in only small, specific amounts called quanta.
A quantum of energy is the minimum amount of energy that can be gained or lost by an atom.

25. The quantum concept
Prior to this discovery, physicists thought that energy could be absorbed and emitted in varying quantities, with no minimum limit to the amount.
Example: think about heating a cup of water on a hot plate. The temperature on the thermometer seems to rise continuously, not in small jumps.
Actually, the water’s temperature is increasing in very tiny steps as the water molecules absorb quanta of energy.
Because these steps are so small, the temperature just appears to be continuous.

26. Planck’s constant
Planck demonstrated mathematically that a relationship exists between the energy of a quantum and the frequency of the radiation that is emitted.
E = hν
h = x J s (the joule, J, is a unit of energy)
In 1905, Einstein described these quanta of energy as particles that he called photons.

27. Practice problem
What is the energy of a photon from the violet portion of the Sun’s light if it has a frequency of 7.23 x 1014 Hz.
E = hν
E = 6.626x10-34 (7.23 x 1014)
E = 4.79 x 10-19J

28. Bohr’s model of the atom and the particle nature of light
Recall that Bohr stated that electrons exist on energy levels.
When enough energy is added to an atom, electrons jump to a higher energy level.
(This is sometimes referred to as a “quantum leap” or the electrons are said to be in an excited state)
This is an unstable situation and the electron will drop back to its former energy level, the ground state, releasing the absorbed energy in the form of light.
The color of light emitted depends on what energy level an electron is dropping to.
A revision in Bohr’s model of the atom was necessary to explain the relationship among atomic structure, electrons, and the unique frequencies of light emitted.