1. Atomic Structure and Periodicity
AP Chemistry
Chapter 7

2. 7.1 EMR
Electromagnetic radiation – a form of energy that exhibits wavelike behavior as it travels through space.
Types include visible light, X rays, ultraviolet light, infrared light, microwaves, and radio waves.
Electromagnetic spectrum – All the forms of electromagnetic radiation together

3. Electromagnetic Radiation
Waves have a wavelength – distance between corresponding points on adjacent waves
Use the Greek letter “lambda”, , for wavelength, and units are length units (m, cm, nm)

4. Electromagnetic Radiation

5. Electromagnetic Radiation
Waves have a frequency – number of waves that pass a given point in a specific time
Use the Greek letter “nu”, , for frequency, and units are “cycles per sec” or Hertz (Hz)

7. Electromagnetic Radiation
All radiation travels at the same speed of light.
c = 3.00 x 108 m/s
 •  = c
This means that  must be in meters and  must be in Hertz (1/s) so that units cancel.

8. Electromagnetic Spectrum
Long wavelength  small frequency
Short wavelength  high frequency
increasing frequency
increasing wavelength

9. Electromagnetic Spectrum

10. Radio waves Low frequency, Long wavelength 1 m - 1km Microwaves 1 cm
                      
Low frequency,
Long wavelength 1 m - 1km
Microwaves
                 
1 cm
Infra-Red
0.01 mm
Visible light nm
Ultra-Violet
100 nm
X-Rays
1 nm
Gamma Rays
High frequency,
0.01 nm
Short wavelength

11. Problems with Wave Theory of Light
Scientific belief around the 1900’s was that there was NO relationship between matter and light
Light given off by objects that were heated to high temperatures could not be explained.

12. Black Body Radiation

13. 7.2 Nature of Matter Max Planck
Stated that objects radiated energy in small packets of energy called quanta
quantum- a specific amount of energy that can be gained or lost by an atom

14. Particle Behavior of Light
Energy and frequency are directly related
E=hn
E is energy (J)
h is Planck’s constant
h = x J s

15. Photoelectric Effect Thomson (1839)
First to observe the photoelectric effect
photoelectric effect - the emission of electrons from a metal surface when exposed to light of a specific energy.

16. Photoelectric Effect 1905 Albert Einstein
stated that EMR could be viewed as a stream of particles “photons”
photon- a quantum of light
energy of these photons could be calculated by Planck’s equation
stated that the photons strike the electrons therefore ejecting them from the metal

17. Photoelectric Effect

18. Dual Wave-Particle Behavior Of Light
The success of Einstein’s work in explaining the photoelectric effect was largely responsible for the acceptance of the particle behavior of light
Ephoton = hn
E = mc2

19. Can matter act as a wave?
Using Einstein’s and Planck’s equations, de Broglie derived:
The momentum, mv, is a particle property, whereas  is a wave property.
In one equation de Broglie summarized the concepts of waves and particles as they apply to low mass, high speed objects.

20. Sample Problem
Compare the wavelength for an electron (mass = 9.11 x kg) traveling at a speed of 1.0 x 107 m/s with that for a ball (mass = 0.10 kg) traveling at 30 m/s.

21. Dual Wave-Particle Behavior Of Matter
Energy is a form of matter.
All matter exhibits both particle and wave properties.
Large pieces of matter (i.e. baseball) exhibits mostly particle properties.
Tiny pieces of matter (i.e. photons) exhibits mostly wave properties.
Pieces of matter somewhere in the middle (i.e. electrons) clearly show both types of properties!

22. Kirchoff and Robert Bunsen (1854)
Observed that light was given off when they heated different chemicals in their designed burner
They passed the light through a prism and saw separate lines instead of a continuous spectrum.

23. Absorption and Emission Spectra
Emission spectra- the colors produced by an object when burned or heated.
Absorption spectra- the colors that are not shown, rather absorbed in the spectrum

24. 7.3 Hydrogen Spectrum Only four lines are emitted:
Red, green, blue, violet
Only certain energies are allowed.

25. Why do elements produce these lines?
To understand emission spectrum, we must understand these two terms:
Ground state: the lowest energy state for the electron
Excited state: state where electron has higher energy than ground state.

26. Why do elements produce these lines?
Atoms are heated, which adds energy. The electron become excited (thus unstable). They want to return to their normal, or ground state. To do so, they give off energy in the form of EMR.

27. Scientists associated with the H spectrum
Balmer: developed a numerical relationship between the wavelength of the lines in the spectrum and the amount of energy
Lyman: discovered lines produced in the UV range.
Paschen: discovered lines produced in the IR range

28. 7.4 The Bohr Model Neils Bohr
-worked with Rutherford to study the H spectrum.
-Bohr’s model is sometimes referred to as the “Planetary model” based upon his postulates.

29. Bohr Model of the Atom Bohr Model of the Atom
Postulates of Bohr’s model:
The single electron of hydrogen can circle the nucleus in fixed paths called orbits or stationary states.
The electron can jump to higher orbits when energy is added.
The angular momentum of the electron is quantized.
-The electron’s energy can be calculated in the different orbits.

30. Bohr Model of the Atom Bohr Model of the Atom
How does this relate to the Hydrogen spectrum?
Bohr calculated the energy that the electron would lose as it fell from higher orbits to lower orbits.
Bohr’s calculations agreed exactly with Lyman, Balmer and Paschen’s observations.

31. Bohr Model of the Atom

32. Sample Problem
Calculate the energy required to excite the hydrogen electron from n=1 to n=2.

33. Sample Problem
Calculate the energy required to completely remove the electron from a hydrogen atom in its ground state.
ninitial = 1 to nfinal = ∞

34. Bohr Model of the Atom Bohr Model of the Atom
Bohr’s model worked very well for the Hydrogen atom.
Through Bohr’s work, as well as the other scientists mentioned, a very good understanding of the electron within the atom was now in place.

35. Downfalls to Bohr’s Model
Downfalls to Bohr’s Model of the Atom
Downfalls to Bohr’s Model
1. Bohr’s model of the atom worked very well for the hydrogen atom and the He+, but failed when applied to multielectron atoms.
2. Bohr’s model could not explain why the electron could not exist between orbits.

36. Now What? We need a new approach to the atom!
Big Three: de Broglie, Heisenberg & Schrodinger
Developed wave mechanics AKA quantum mechanics (7.5)

37. Heisenberg’s Uncertainty Principle
With respect to atomic particles, we cannot determine exactly
1. the position
2. direction of motion
AND
3. speed
simultaneously.

38. Schrodinger’s Wave Equation
Schrödinger proposed an equation that contains both wave and particle terms.
Solving the equation leads to wave functions .
The wave function gives the probability distribution of an electron.
We call wave functions orbitals.

39. 7.6 Quantum Numbers
When we solve the Schrödinger equation for the hydrogen atom, we find many wave functions (orbitals) that satisfy it.
Each orbital is characterized by a series of numbers called quantum numbers that describe various properties of the orbital.

40. Principal Quantum Number, n
Related to the size and energy of the orbital – think energy level
n has integer values: 1,2,3…
As n becomes larger, the atom becomes larger and the electron is further from the nucleus.
A larger n value also corresponds to higher energy because the electron is less tightly bound to the nucleus.

41. Angular Momentum Quantum Number, l
Related to the shape of the atomic orbitals
This quantum number depends on the value of n. The values of l begin at 0 and increase to (n - 1).
Because we use numbers to describe the first quantum number, we usually use letters for l (s for l =0, p for l = 1, d for l =2 and f for l = 3).
Usually we refer to the s, p, d and f-orbitals.

42. Magnetic Quantum number, ml
Provides the 3D orientation of the orbital in space
Value depends on l. The magnetic quantum number has integer values between -l and +l.

43. Quantum Numbers

44. 7.7 Orbital Shapes & Energies
Each orbital has a unique probability distribution.
Nodes = areas of zero probability
To simplify, we think of orbitals in terms of their overall shapes, which becomes larger as n increases.

45. p orbitals

46. d orbitals

47. f orbitals

48. Energies of orbitals in Hydrogen
For hydrogen, energy is determined by value of n
All orbitals with the same value of n have the same energy – they are degenerate.

49. 7.8 Electron Spin & Pauli Exclusion Principle
Developed by Samuel Goudsmit & George Uhlenbeck (University of Leyden in the Netherlands)
4th quantum number necessary to account for the details of emission spectra of atoms
Electron has a magnetic moment with two possible orientations when placed in an external magnetic field.
Magnetic spin quantum number ms can only have two possible values +½ and -½

50. 7.8 Electron Spin & Pauli Exclusion Principle
Wolfgang Pauli developed Pauli exclusion principle
In a given atom, no two electrons can have the same set of four quantum numbers
An orbital can hold only 2 electrons, and they must have opposite spins

51. 7.9 Polyelectronic Atoms Effective Nuclear Charge
Electrons are attracted to the nucleus, but repelled by the electrons that screen it from the nuclear charge.
The nuclear charge experienced by an electron depends on its distance from the nucleus and the number of core electrons.
As the average number of screening electrons (S) increases, the effective nuclear charge (Zeff) decreases.
As the distance from the nucleus increases, S increases and Zeff decreases.

53. 7.10 History of the Periodic Table
Dobereiner – triads (groups of 3 elements share similar properties)
Newlands – octaves (certain properties repeat for every eighth element)
Meyer & Mendeleev – present form of periodic table
Mendeleev – considered father of periodic table because he predicted the existence and properties of still unknown elements and left space for them in his periodic table
Fundamental difference – modern periodic table organized by atomic number not mass

54. 7.11 Aufbau Principle & the Periodic Table
As protons are added one by one to the nucleus to build up the elements, electrons are similarly added
Electron configurations tells us in which orbitals the electrons for an element are located.

55. Periods 1 - 3
Three rules:
electrons fill orbitals starting with lowest n and moving upwards;
no two electrons can fill one orbital with the same spin (Pauli);
for degenerate orbitals, electrons fill each orbital singly before any orbital gets a second electron (Hund’s rule).

57. Period 4 and Beyond After Ca the d orbitals begin to fill.
After the 3d orbitals are full the 4p orbitals being to fill.
From Ce onwards the 4f orbitals begin to fill.
Note: La: [Xe]6s25d14f0
Elements Ce - Lu have the 4f orbitals filled and are called lanthanides.
Elements Th - Lr have the 5f orbitals filled and are called actinides.
Most actinides are not found in nature.

58. Electron Configurations and the Periodic Table
The periodic table can be used as a guide for electron configurations.
The period number is the value of n.
Groups 1 and 2 have the s-orbital filled.
Groups have the p-orbital filled.
Groups have the d-orbital filled.
The lanthanides and actinides have the f-orbital filled.
Note that the 3d orbital fills after the 4s orbital. Similarly, the 4f orbital fills after the 6s orbital.

60. Noble Gas Notation
There is a shorthand way of writing electron configurations
Write the core electrons corresponding to the filled Noble gas in square brackets.
Write the valence electrons explicitly.
Example, P: 1s22s22p63s23p3
but Ne is 1s22s22p6
Therefore, P: [Ne]3s23p3.

61. Practice Problem
Determine the expected electron configurations for each of the following: S Ba
Ni2+ Eu
Ti+

62. Effective Nuclear Charge - revisited
Many properties of atoms depend on electron configurations and how strongly valence electrons are attracted to the nucleus.
Coulomb’s Law – strength of the interaction between 2 electrical charges depends on the size of the charges and the distance between them.
Zeff = Z – S where Z is # protons in nucleus and S is number of core electrons
Explains differences in sublevel energies but also describes periodic trends.

63. Effective nuclear charge
The effective nuclear charge increases as we move across any row (period) of the periodic table (Z gets larger while S stays the same)
The effective nuclear charge also increases as we go down a column (group) of the periodic table, but the effect is far less than going across a row.

64. Atomic Radius Simple diatomic molecule
The distance between the two nuclei is called the bond distance.
If the two atoms which make up the molecule are the same, then half the bond distance is called the covalent radius of the atom.

65. Atomic Radius
Atomic size varies consistently through the periodic table.
As we move down a group, the atoms become larger.
As we move across a period, atoms become smaller.
There are two factors at work:
principal quantum number, n (down a group)
the effective nuclear charge, Zeff (across a period)

66. Atomic Radius

67. Ionization Energy
Ionization energy – minimum amount of energy required to remove an electron from the ground state of an isolated gas atom or ion.
Na(g)  Na+(g) + e- First ionization energy
Na+(g)  Na2+(g) + e- Second ionization energy
The greater ionization energy, the more difficult it is to remove the electron.

68. Ionization Energy
Ionization energy increases for each additional electron removed from an atom.
There is a sharp increase in ionization energy when a core (non-valence) electron is removed.

69. Ionization Energy Trend
Same factors influence ionization energy – effective nuclear charge & distance of electron from nucleus.
Increasing effective charge or decreasing distance from nucleus increases attraction between electron & nucleus – more difficult to remove an electron so ionization energy increases. (Both happen when move across row)
As we move down group, the atomic radius increases (due to larger n) while effective nuclear charge only increases slightly. Attraction between nucleus & electron decreases, so ionization energy decreases.