1. Chapter 10- Part One Modern Atomic Theory
Objectives:
Review history… (10.1)
Describe electromagnetic radiation (10.2)
Describe the Bohr atom (10.3)
Explain energy levels of electrons and diagram atomic structures for elements (10.4 & 10.5)

2. Review… Dalton Thomson Rutherford
Model doesn’t explain how the negative electron can stay in orbit and not be attracted to the positive proton

3. Electromagnetic Radiation
Light travels in waves, similar to waves caused by a moving boat or a pebble tossed in a pond
Light is a form of Electromagnetic Radiation
Form of energy that exhibits wavelike behavior as it travels through space

4. Electromagnetic Radiation
All waves can be described in 3 ways:
Amplitude – the height of the wave, results in the brightness or intensity of the light
Wavelength (l): distance between consecutive peaks in a wave
Frequency (n): number of waves that pass a given point in a second

5. Electromagnetic Radiation
Speed of light in air: Electromagnetic radiation moves through a vacuum at speed of 3.00 x 108 m/s
Since light moves at constant speed there is a relationship between wavelength and frequency:
c = ln
Wavelength and frequency are inversely proportional
C = speed of light
L = wavelength
V = frequency

6. Electromagnetic Spectrum
Visible spectrum – white light (what we can see)
Wavelengths of visible spectrum are 350 nm to 700 nm

7. Quantum Theory Wave theory does not explain Heated iron gives off heat
1st red glow yellow glow white glow
How elements such as barium and strontium give rise to green and red colors when heated

8. Energy is released in Quanta
Quantum Theory
Max Planck ( )
Proposed that there is a fundamental restriction on the amounts of energy that an object emits or absorbs, and he called each of these pieces of energy a quantum.
Energy is released in Quanta

9. E = hn E = energy Quantum Theory
A quantum is a finite quantity of energy that can be gained or lost by an atom
E = hn E = energy
v = frequency
h = x J/s
This constant, h, is the same for all electromagnetic radiation

10. Photoelectric Effect
The emission of electrons by certain metals when light shines on them
Albert Einstein (1905) used Planck’s equation to explain this phenomenon;
proposed that light consists of quanta of energy that behave like tiny particles of light
Photon = individual quantum light (also known as a particle of radiation)

11. Photoelectric Effect
He (Einstein) explained that the photoelectric effect would not occur if the frequency and therefore the energy of each photon is too low to dislodge an electron.
Analogy:
70 cents placed in soda machine: no soda
30 cents more and you will get your soda

12. Now… Light can be described as both particles and waves
Dual Wave-Particle Nature of Light was accepted
What does this mean for the atom???

13. Line Spectrum Elements in gaseous states give off colored light
High temperature or high voltage
Always the same
Each element is unique
Also known as the “atomic emission spectrum of an element”
Visible light or electromagnetic radiation is emitted when an atom passes from a state of higher potential energy to a state of lower potential energy

14. Line Spectrum Ground state Excited state Lowest energy level available
State in which electron has a higher potential energy than in its ground state
Farther from nucleus
Higher potential energy

15. Line Spectrum
Electron falls from higher energy level to lower one…emits light at a specific frequency
Color of light emitted depends on difference between excited state and ground state
See figure 10.5 page 201

16. Line Spectrum
Each band of color is produced by light of a different wavelength
Each particular wavelength has a definite frequency and has definite energy
Each line must therefore be produced by emission of photons with certain energies

17. Line Spectrum

18. Line Spectrum
Whenever an excited electron drops from such a specific excited state to its ground state (or lower excited state) it emits a photon
The energy of this photon is equal to the difference in energy between the initial state and the final state.

19. Niels Henrik David Bohr
Physicist
Worked with Rutherford
1912
Studying line spectra
of hydrogen

20. Niels Henrik David Bohr
1913 – proposed new atomic structure
Electrons exist in specific regions away from the nucleus
Electrons revolve around nucleus like planets around the sun

21. The Bohr Atom Nucleus with protons and neutrons
Electrons move in “stationary states” which are stable (paths or orbits)
When an electron moves from one state to another the energy lost or gained is done is ONLY very specific amounts
Each line in a spectrum is produced when an electron moves from one stationary state to another

22. The Bohr Atom
Model didn’t seem to work with atoms with more than one electron
Did not explain chemical behavior of the atoms
Safeco field example

23. Wave Matters… Louis de Broglie (1924)
Proposed that electrons might have a wave-particle nature
Used observations of normal wave activity
Safeco field example

24. Wave Matters… Erwin Schrodinger (1926)
Used mathematical understanding of wave behavior – devised an equation that treated electrons moving around nuclei as waves
Quantum Theory
Safeco field example

25. Quantum Theory
Describes mathematically the wave properties of electrons and other very small particles
Applies to all elements (not just H)

26. Energy Levels of Electrons
Principal energy levels
Designated by letter n
Each level divided into sublevels
1st energy level has 1 sublevel
2nd energy level has 2 sublevels
Etc.

27. Energy Levels of Electrons

28. Orbitals Electrons don’t actually orbit like planets
Orbital: region in space where there is a high probability of finding a given electron
Each orbital sublevel can hold 2 electrons
Safeco field example

29. Orbitals Each sublevel (orbital) has a specific shape
Safeco field example

30. Orbitals
Pauli exclusion principle: an atomic orbital can hold a maximum of two electrons which must have opposite spins
Electrons can only spin in two directions
Shown with arrows
Safeco field example

31. Rules for Orbital Filling
Pauli’s Exclusion Rule
No two electrons have the same set of quantum numbers
Hund’s Rule
Electrons will remain unpaired in a given orbital until all orbitals of the same sublevel have at least one electron
1s s 2p s 3p

32. Rules for Orbital Filling
Diagonal Rule
The order of filling once the d & f sublevels are being filled
Due to energy levels

33. Rules for Orbital Filling
Safeco field example

34. Quantum Numbers
Numbers that specify the properties of atomic orbitals and their electrons
Principle Quantum Numbers:
Symbolized by n, indicates the main energy levels surrounding a nucleus, which indicates the distance from the nucleus (shells or levels)

35. Quantum Numbers Orbital Quantum Number:
Indicates the shape of an orbital
(subshell or sublevels)
s, p, d, f
Principal Quantum # Orbital Quantum #
1 1s
2 2s, 2p
3 3s, 3p, 3d
4 4s, 4p, 4d, 4f

36. Quantum Numbers Magnetic Quantum Number:
Indicates the orientation of an orbital about the nucleus
Orbital position with respect to the 3-dimensional x, y, and z axes

37. Quantum Numbers Spin Quantum Number:
Indicates two possible states of an electron in an orbital
Type of Orbital Number of Orbitals
s 1 ( )
p 3 (x, y, z) ( , , ,)
d 5 ( , , , , )
f 7
Each orbital holds a maximum of 2 electrons

38. Application of Quantum Numbers
Several ways of writing the address or location of an electron
Lowest energy levels are filled first
Electron Configuration: using the diagonal rule, the principal quantum number (n), and the sublevel write out the location of all electrons
12C:
32S:
1s22s22p2
1s22s22p63s23p4

39. Application of Quantum Numbers
Orbital filling electron diagram: using Hund’s rule and the diagonal rule write out the location of all electrons
See examples on whiteboard

40. Chapter 10 – Part Two The Periodic Table
Objectives:
Understand the arrangement of the Periodic Table (10.6)
Identify connections between electron configuration and placement on the periodic table

41. The Periodic Table 1869 – arrangement proposed by Dmitri Mendeleev
And Lothar Meyer (different layout)
Still similar today
Based on increasing atomic masses and other characteristics
Was able to predict properties of elements not yet discovered….and was correct!

42. The Periodic Table Horizontal rows Vertical Columns Periods
Corresponds to outermost energy level
Vertical Columns
Groups or families
Similar properties; reactions

43. The Periodic Table Several systems for naming groups
Left to right, 1-18
Roman numerals and A and B
Used in this book
Group A: Representative Elements
Noble Gases
IA – Alkali Metals
IIA – Alkaline Earth Metals
VIIA - Halogens
Group B: Transition Elements

44. The Periodic Table
Chemical behavior and properties of elements in a particular family similar
Have the same outer shell electron configuration
Figure page 211
Noble gas configuration (shortcut)
Use previous noble gas in square brackets
Finish with valence electrons

45. The Periodic Table Examples:
K is 1s22s22p63s23p64s1 or [Ar]4s1
Ca is 1s22s22p63s23p64s2 or [Ar]4s2
Write abbreviated configuration for the following elements: Fr Y

46. The Periodic Table
Arrangement of Periodic Table also means that elements filling similar orbitals are grouped
s block
p block
d block
f block
Know these blocks…

47. The Periodic Table - Highlights
The number of the period corresponds to the highest energy level occupied by electrons in that period
The group numbers for the representative elements are equal to the total number of valence electrons in that group

48. The Periodic Table - Highlights
The elements of a family have the same outermost electron configuration
(just different energy levels)
The elements within each of the s, p, d, and f blocks are filling the corresponding orbitals
There are some discrepancies with order of filling
(not covered in this book)