1. The Modern Atom Figure: 05-00CO Caption:
A silicon wafer is “doped” with GaAs and yields dozens of computer chips.

2. Emission Line Spectra
When an electrical voltage is passed across a gas in a sealed tube, a series of narrow lines is seen.
These lines are the emission line spectrum. The emission line spectrum for hydrogen gas shows three lines: 434 nm, 486 nm, and 656 nm.

3. Wave Nature of Light
Light is made of particles (photons) with no mass carrying energy.
Photons travels through space as a wave, similar to an ocean wave.
A wave has characteristics:
a) Wavelength is the distance light travels in one cycle.
b) Frequency is the number of wave cycles completed each second.
c) Speed: Light has a constant speed: 3.00 × 108 m/s.

4. Wavelength vs. Frequency
The longer the wavelength of light, the lower the frequency.
The shorter the wavelength of light, the higher the frequency.

5. Energy and Frequency
There is a relationship between the energy and the frequency of photons: E = h x , where E is the energy, h is called Planck’s constant (h= 6.62x J.s), and , is the frequency.
The frequency of an electromagnetic radiation determines the color of the light.
Photons with high frequency carry more energy than photons with low frequency.

6. Visible Spectrum
Radiation composed of only one wavelength is called monochromatic.
White light is made of photons of different wavelengths.
These photons can be separated into a continuous spectrum of colors.
The visible spectrum is the range of wavelengths between 400 and 700 nm.
Radiant energy that has a wavelength lower than 400 nm and greater than 700 nm cannot be seen by the human eye.

8. Radiant Energy Spectrum
The complete radiant energy spectrum is an uninterrupted band, or continuous spectrum.

9. Bohr Model of the Atom
Niels Bohr speculated that electrons orbit about the nucleus in fixed energy levels.
Electrons are found only in specific energy levels, and nowhere else.
The electron energy levels are quantized.

10. Evidence for Energy Levels
The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off.

11. “Atomic Fingerprints”
The emission line spectrum of each element is unique.
We can use the line spectrum for the identify of elements, using their “atomic fingerprint”.

12. Bohr Model
Colors from excited gases arise because electrons move between energy states in the atom.

13. The Wave/Particle Nature of Light
In 1900, Max Planck proposed that radiant energy is not continuous, but is emitted in small bundles. This is the quantum concept.
Radiant energy has both a wave nature and a particle nature.

14. The Photoelectric Effect
The photoelectric effect provides evidence for the particle nature of light -- “quantization”.
If light shines on the surface of a metal, there is a point at which electrons are ejected from the metal.
The electrons will only be ejected once the threshold frequency is reached.
Below the threshold frequency, no electrons are ejected.
Above the threshold frequency, the number of electrons ejected depend on the intensity of the light.

15. The Photoelectric Effect

16. The Quantum Concept
The quantum concept states that energy is present in small, discrete bundles.
For example:
A tennis ball that rolls down a ramp loses potential energy continuously.
A tennis ball that rolls down a staircase loses potential energy in small bundles. The loss is quantized.

17. Energy Levels and Sublevels
It was later shown that electrons occupy energy sublevels within each level.
These sublevels are given the designations s, p, d, and f.
These designations are in reference to the sharp, principal, diffuse, and fine lines in emission spectra.
The number of sublevels in each level is the same as the number of the main level.

18. Quantum Mechanical Model
An orbital is the region of space where there is a high probability of finding an atom.
In the quantum mechanical atom, orbitals are arranged according to their size and shape.
The higher the energy of an orbital, the larger its size.
s-orbitals have a spherical shape.

19. Shapes of p-Orbitals There are three different p sublevels.
p-orbitals have a dumbbell shape.
Each of the p-orbitals has the same shape, but each is oriented along a different axis in space.

20. s-orbitals
p-orbitals
d-orbitals

21. Orbitals and their Energy
H-Atom Other atoms

22. Energy Levels and Sublevels
The first energy level has 1 sublevel: 1s
The second energy level has 2 sublevels:
2s and 2p
The third energy level has 3 sublevels:
3s, 3p, and 3d

23. Electron Occupancy in Sublevels
The maximum number of electrons in each of the energy sublevels depends on the sublevel:
The s sublevel holds a maximum of 2 electrons.
The p sublevel holds a maximum of 6 electrons.
The d sublevel holds a maximum of 10 electrons.
The f sublevel holds a maximum of 14 electrons
The maximum electrons per level is obtained by adding the maximum number of electrons in each sublevel.

24. Figure: 05-T03
Title:
Table 5.3
Caption:
Distribution of Electrons by Energy Level
Notes:
The higher the energy level, the more possible sublevels there are.

25. Conclusions Light has both the properties of waves and particles.
The particles of light are referred to as photons.
The energy of photons is quantized.
Electrons exist around the nucleus of atoms in discrete, quantized energy levels.
Electrons fill energy sublevels starting with the lowest energy sublevel and filling each successive level of higher energy.