2. 10.1 History of the Atom Dalton’s Model of the Atom
John Dalton proposed that all matter is made up of tiny particles.
These particles are molecules or atoms.
Molecules can be broken down into atoms by chemical processes.
Atoms cannot be broken down by chemical or physical processes.
Chapter 5

3. Thomson’s Model of the Atom
J.J. Thomson proposed a subatomic model of the atom in 1903.
Thomson proposed that the electrons were distributed evenly throughout a homogeneous sphere of positive charge.
This was called the “plum pudding” model of the atom.
Chapter 5

4. Rutherford’s Gold Foil Experiment
Rutherford’s student fired alpha particles at thin gold foils. If the “plum pudding” model was correct, α-particles should pass through undeflected.
At the center of an atom is the atomic nucleus, which contains the atom’s protons.
Chapter 5

5. 10.2 Radiant Energy Spectrum
The complete radiant energy spectrum is an uninterrupted band, or continuous spectrum.
The radiant energy spectrum includes most types of radiation, most of which are invisible to the human eye.
Chapter 5

6. Visible Spectrum
Light usually refers to radiant energy that is visible to the human eye.
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.
Chapter 5

7. 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.
An individual unit of light energy is a photon.
Chapter 5

8. c= ln Wave Nature of Light
Light travels through space as a wave, similar to an ocean wave.
Wavelength is the distance light travels in one cycle.
Frequency is the number of wave cycles completed each second.
Light travels at a constant speed: 3.00 × 108 m/s (given the symbol c).
c= ln
Chapter 5

9. Waves
The distance between corresponding points on adjacent waves is the wavelength ().
The number of waves passing a given point per unit of time is the frequency ().
Amplitude

10. Wavelength vs. Frequency
The longer the wavelength of light, the lower the frequency.
The shorter the wavelength of light, the higher the frequency.
Chapter 5

11. Sample Calculating Frequency from Wavelength
The yellow light given off by a sodium vapor lamp used for public lighting has a wavelength of 589 nm. What is the frequency of this radiation?
Solution
589 nm m X 10-7m
1 X 109 nm
3.00 X 108 m/s
5.89 X 10-7 m 11

12. The Nature of Energy
Einstein used this assumption to explain the photoelectric effect. (electrons are ejected from metals when light from specific wavelengths are applied)
He concluded that energy is proportional to frequency:
E = h
where h is Planck’s constant,
6.63  10−34 J-s.

13. The Nature of Energy
Therefore, if one knows the wavelength of light, one can calculate the energy in one photon, or packet, of that light:
c = 
E = h

14. Sample Energy of a Photon
Calculate the energy of one photon of yellow light that has a wavelength of 589 nm (n= 5.09 X 1014 s-)
Solution
The value of Planck’s constant, h, is given both in the text and in the table of physical constants on the inside back cover of the text, and so we can easily calculate E: 14

15. The Wave Nature of Matter
Louis de Broglie proposed that if light can have material properties, matter should exhibit wave properties.
He demonstrated that the relationship between mass and wavelength was
 = h mv

16. h mv  = What is the wavelength of an electron moving with a speed of
Sample Matter Waves
What is the wavelength of an electron moving with a speed of
5.97 × 106 m/s? The mass of the electron is 9.11 × kg.
Solution
 = h mv
6.40 X 1034 J*s
(9.11 X10-31 kg)(5.97 X106 m/s)
mass Kg
Velocity m/s
1.22 X m 16

17. 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.
Chapter 5

18. “Atomic Fingerprints”
The emission line spectrum of each element is unique.
We can use the line spectrum to identify elements using their “atomic fingerprint.”
Chapter 5

19. 10.5 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.
Chapter 5

20. 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.
Chapter 5

21. Evidence for Energy Levels
Bohr realized that this was the evidence he needed to prove his theory.
The electric charge temporarily excites an electron to a higher orbit. When the electron drops back down, a photon is given off.
The red line is the least energetic and corresponds to an electron dropping from energy level to energy level 2.
Chapter 5

22. 10.7 Quantum Mechanics
Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated.
This is known as quantum mechanics.

23. Quantum Model
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.
Chapter 5

24. 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
Chapter 5

25. 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.
Chapter 5

26. Electrons per Energy Level
Chapter 5

27. Quantum Mechanical Model
An orbital is the region of space where there is a high probability of finding an electron.
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
Chapter 5

28. Shapes of p-Orbitals
Recall that 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.
Chapter 5

29. d- orbitals

30. F-orbitals

31. 10.9 Electron Configurations
Electrons are arranged about the nucleus in a regular manner. The first electrons fill the energy sublevel closest to the nucleus.
Electrons continue filling each sublevel until it is full, and then start filling the next closest sublevel.
A partial list of sublevels in order of increasing energy is:
1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d …
Chapter 5

32. Filling Diagram for Sublevels
The order does not strictly follow 1, 2, 3, etc.
For now, use this figure to predict the order of sublevel filling.
Chapter 5

33. Orbital Diagrams Each box in the diagram represents one orbital.
Half-arrows represent the electrons.
The direction of the arrow represents the relative spin of the electron.

34. Electron Configurations
The electron configuration of an atom is a shorthand method of writing the location of electrons by sublevel.
The sublevel is written followed by a superscript with the number of electrons in the sublevel.
If the 2p sublevel contains 2 electrons, it is written 2p2.
The electron sublevels are arranged according to increasing energy.
Chapter 5

35. Electron Configurations follow 3 rules
Rules for e- Configs
Electron Configurations follow 3 rules
Aufbau- electrons fill in lowest energy first (start at the bottom)
Pauli Exclusion- 2 electrons maximum in an orbital, with opposite spins to reduce repulsion
Hund’s- everyone (box) gets one e- before anyone gets seconds- (in degenerate orbitals)

36. Writing Electron Configurations
First, determine how many electrons are in the atom. Bromine has 35 electrons.
Arrange the energy sublevels according to increasing energy:
1s 2s 2p 3s 3p 4s 3d …
Fill each sublevel with electrons until you have used all the electrons in the atom:
Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d 10 4p5
The sum of the superscripts equals the atomic number of bromine (35).
Chapter 5

37. Valence Electrons
When an atom undergoes a chemical reaction, only the outermost electrons are involved.
These electrons are of the highest energy and are furthest away from the nucleus. These are the valence electrons.
The valence electrons are the s and p electrons beyond the noble gas core.
For our purposes we will include ALL valence electrons past the noble gas core.

38. 10.10 e- Configs using the Periodic Table
We fill orbitals in increasing order of energy.
Different blocks on the periodic table (shaded in different colors in this chart) correspond to different types of orbitals.

39. Blocks and Sublevels
We can use the periodic table to predict which sublevel is being filled by a particular element.

40. Noble Gas Core Electron Configuration
Recall, the electron configuration for Na is:
Na: 1s2 2s2 2p6 3s1
We can abbreviate the e- config by indicating the innermost electrons with the symbol of the preceding noble gas.
The preceding noble gas before sodium is neon, Ne. We rewrite the electron configuration:
Na: [Ne] 3s1