1. Periodicity and Atomic Structure
Chapter 5
Periodicity and Atomic Structure

2. Development of the Periodic Table
The periodic table is the most important organizing principle in chemistry.
Periodic table powerpoint – elements of a group have similar properties
Chapter 2 – elements in a group form similar formulas
Predict the properties of an element by knowing the properties of other elements in the group

3. Light and the Electromagnetic Spectrum
Radiation (light) composed waves of energy
Waves were continuous and spanned the electromagnetic spectrum

4. Light and the Electromagnetic Spectrum

5. Light and the Electromagnetic Spectrum

6. Light and the Electromagnetic Spectrum
Speed of a wave is the wavelength (in meters) multiplied by its frequency in reciprocal seconds.
Wavelength x Frequency = Speed
 (m) x  (s–1) = c (m/s–1)
C – speed of light x 108 m/s–1

7. Electromagnetic Radiation and Atomic Spectra
Classical Physics does not explain
Black-body radiation
Photoelectric effect
Atomic Line Spectra

8. Particlelike Properties of Electromagnetic Radiation: The Plank Equation
Blackbody radiation is the visible glow that solid objects emit when heated.
Max Planck (1858–1947): Developed a formula to fit the observations. He proposed that energy is only emitted in discrete packets called quanta.
The amount of energy depends on the frequency:

9. Particlelike Properties of Electromagnetic Radiation: The Plank Equation
A photon’s energy must exceed a minimum threshold for electrons to be ejected.
Energy of a photon depends only on the frequency.

10. Electromagnetic Radiation and Atomic Spectra
Atomic spectra: Result from excited atoms emitting light.
Line spectra: Result from electron transitions between specific energy levels.

11. Electromagnetic Radiation and Atomic Spectra
1/λ = R [1/m2 – 1/n2]

12. Quantum Mechanics and the Heisenburg Uncertainty Principle
Niels Bohr (1885–1962): Described atom as electrons circling around a nucleus and concluded that electrons have specific energy levels.
Erwin Schrödinger (1887–1961): Proposed quantum mechanical model of atom, which focuses on wavelike properties of electrons.

13. Quantum Mechanics and the Heisenburg Uncertainty Principle
Werner Heisenberg (1901–1976): Showed that it is impossible to know (or measure) precisely both the position and velocity (or the momentum) at the same time.
The simple act of “seeing” an electron would change its energy and therefore its position.

14. Wave Functions and Quantum Mechanics
Erwin Schrödinger (1887–1961): Developed a compromise which calculates both the energy of an electron and the probability of finding an electron at any point in the molecule.
This is accomplished by solving the Schrödinger equation, resulting in the wave function, .

15. Wave Functions and Quantum Mechanics
Wave functions describe the behavior of electrons.
Each wave function contains three variables called quantum numbers:
• Principal Quantum Number (n)
• Angular-Momentum Quantum Number (l)
• Magnetic Quantum Number (ml)

16. Wave Functions and Quantum Mechanics
Principal Quantum Number (n): Defines the size and energy level of the orbital. n = 1, 2, 3, 
As n increases, the electrons get farther from the nucleus.
As n increases, the electrons’ energy increases.
Each value of n is generally called a shell.

17. Wave Functions and Quantum Mechanics
Angular-Momentum Quantum Number (l): Defines the three-dimensional shape of the orbital.
For an orbital of principal quantum number n, the value of l can have an integer value from 0 to n – 1.
This gives the subshell notation: l = 0 = s orbital l = 1 = p orbital l = 2 = d orbital l = 3 = f orbital l = 4 = g orbital

18. Wave Functions and Quantum Mechanics
Magnetic Quantum Number (ml): Defines the spatial orientation of the orbital.
For orbital of angular-momentum quantum number, l, the value of ml has integer values from –l to +l.
This gives a spatial orientation of: l = 0 giving ml = 0 l = 1 giving ml = –1, 0, +1 l = 2 giving ml = –2, –1, 0, 1, 2, and so on…...

19. Wave Functions and Quantum Mechanics

20. Problem Why can’t an electron have the following quantum numbers?
(a) n = 2, l = 2, ml = 1 (b) n = 3, l = 0, ml = 3
(c) n = 5, l = –2, ml = 1
Give orbital notations for electrons with the following quantum numbers:
(a) n = 2, l = 1, ml = 1 (b) n = 4, l = 3, ml = –2
(c) n = 3, l = 2, ml = –1

21. The Shapes of Orbitals
s Orbital Shapes:

22. The Shapes of Orbitals
p Orbital Shapes:

23. The Shape of Orbitals
d and f Orbital Shapes:

24. Orbital Energy Levels in Multielectron Atoms

25. Orbital Energy Levels in Multielectron Atoms
Zeff is lower than actual nuclear charge.
Zeff increases toward nucleus ns > np > nd > nf
This explains certain periodic changes observed.

26. Orbital Energy Levels in Multielectron Atoms
Electron shielding leads to energy differences among orbitals within a shell.
Net nuclear charge felt by an electron is called the effective nuclear charge (Zeff).

27. Wave Functions and Quantum Mechanics
Spin Quantum Number:
The Pauli Exclusion Principle states that no two electrons can have the same four quantum numbers.x

28. Electron Configurations of Multielectron Atoms
Pauli Exclusion Principle: No two electrons in an atom can have the same quantum numbers (n, l, ml, ms).
Hund’s Rule: When filling orbitals in the same subshell, maximize the number of parallel spins.

29. Electron Configurations of Multielectron Atoms
Rules of Aufbau Principle:
Lower n orbitals fill first.
Each orbital holds two electrons; each with different ms.
Half-fill degenerate orbitals before pairing electrons.

30. Electron Configurations and Multielectron Atoms
Li   s2 2s1
1s 2s
Be   s2 2s2
B    s2 2s2 2p1
1s 2s 2px 2py 2pz
C     s2 2s2 2p2
1s 2s 2px 2py 2pz

31. Electron Configurations and Multielectron Atoms
N      1s2 2s2 2p3
1s 2s 2px 2py 2pz
O      1s2 2s2 2p4
1s 2s 2px 2py 2pz
Ne      1s2 2s2 2p5
S [Ne]     [Ne] 3s2 3p4
3s 3px 3py 3pz

32. Problems Give the ground-state electron configurations for:
Ne (Z = 10) Mn (Z = 25) Zn (Z = 30)
Eu (Z = 63) W (Z = 74) Lw (Z = 103)
Identify elements with ground-state configurations:
1s2 2s2 2p4 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s2 4d6
1s2 2s2 2p6 [Ar] 4s2 3d1 [Xe] 6s2 4f14 5d10 6p5

33. Electron Configurations and the Periodic Table

34. Some Anomalous Electron Configurations
Anomalous Electron Configurations: Result from unusual stability of half-filled & full-filled subshells.
Chromium should be [Ar] 4s2 3d4, but is [Ar] 4s1 3d5
Copper should be [Ar] 4s2 3d9, but is [Ar] 4s1 3d10
In the second transition series this is even more pronounced, with Nb, Mo, Ru, Rh, Pd, and Ag having anomalous configurations (Figure 5.20).

35. Electron Configurations and Periodic Properties: Atomic Radii

36. Optional Homework
Text – 5.24, 5.26, 5.28, 5.30, 5.32, 5.34, 5.44, 5.56, 5.58, 5.66, 5.68, 5.70, 5.72, 5.76, 5.78, 5.82, 5.84, 5.94, 5.98, 5.108
Chapter 5 Homework online

37. Required Homework
Assignment 5