1. Chapter 5: Electrons in Atoms
Chemistry

2. 5.1 Light and Quantized Energy
Rutherford’s Nuclear Atomic Model was incomplete, did not explain
How electrons were arranged
Why electrons were not pulled into the nucleus
The differences in chemical behavior

3. Wave Nature of Light
In 1900s, observation of light when elements were heated were observed
Analysis of light revealed an elements chemical behavior is related to the arrangement of electrons in its atoms

4. Electromagnetic Radiation
Form of energy that exhibits wavelike behavior as it travels through space
Includes visible light which is only a small portion of EMR
Visible Spectrum – ROY G BIV

5. Electromagnetic Spectrum Chart

7. Wave

8. Wavelength Symbol: λ (Lambda)
Shortest distance between equivalent points (Peak to Peak)
Units = m, cm, nm
nm=nanometers (10-9)

9. Frequency Symbol: ν (nu)
Number of waves that pass a given point per second
Units: hertz (Hz or 1/sec)
652 Hz=652 waves/sec=652/sec=652s-1

10. Waves (Continue) Wavelength and frequency have an inverse relationship
As wavelength increases, frequency decreases
Amplitude: Height of a Wave
Speed and Amplitude are not affected by wavelength

11. Light All electromagnetic waves travel at 3.00 x 108 m/s in a vacuum
This is known as the speed of light
Symbol is c
Speed of light is a product of wavelength and frequency
c = λν
All light travels at the same speed but with different wavelength and frequency

12. Practice Problems
A helium-neon laser emits light with a wavelength of 655 nm. What is the frequency of the light?
What is the wavelength of x-rays having a frequency of 4.80 x 1017 Hz?
An FM radio station broadcasts at a frequency of 98.5 MHz. What is the wavelength of the station’s broadcast signal?

13. Electromagnetic Spectrum
Electromagnetic Spectrum: classifies light based on wavelength and frequency
Encompasses all forms of electromagnetic radiation
Visible Spectrum: Roy G. Biv
Only difference in types of radiation is wavelengths and frequencies
Energy increases with greater frequency
Violet light has shorter wavelength, greater frequency than red light, therefore violet light has more energy that red light.

14. Particle Nature of Light
Wave model could not explain
Why heated objects emit light at given temperatures
Why some metals emit electrons when light shines on them
Max Planck studied light emitted from heated objects,
Discovered that matter can gain or lose energy in small, specific amounts called quanta
Quantum: the minimum amount of energy that can be lost or gained by an atom
Proposed emitted light (from glowing objects) is quantized
Equantum= energy
h= Planck’s constant= x J.s
J= Joules = Kg .m2/s2 (SI unit of energy)
ν= frequency

15. Plank’s Equation Equantum=hν
Planck’s Theory= for a given frequency matter can emit or absorb energy only in whole number multiples ( 1hν, 2hν, 3hν, 4hν)
As energy increases, frequency increases
Equation explains why violet light has more energy than red light

16. Photoelectric Effect
Einstein reasoned that light acted both wave like and particle like
Reasoned that light acted like a stream of particles, knocking electrons out of atoms
Photoelectric Effect = electrons called photo electrons are emitted from metals surface when light of a certain frequency shines on surface

17. Photoelectric Effect
Light of a certain minimum frequency ejects electrons from the surface of metal

18. Planck and Einstein
An energy increases, frequency increases

19. Practice Problems
Calculate the energy of a gamma ray photon whose frequency is 5.02 x 1020 Hz.
What is the difference in energy between a photon of violet light with a frequency of 6.8 x 1014 Hz and a photon of red light with a frequency of 4.3 x 1014 Hz?
Calculate the energy of a photon of ultraviolet light that has a wavelength of 49.0 nm.

20. Atomic Emission Spectra
Atomic Emission Spectra: when atoms of an element in the gaseous state are excited by energy they emit light
The light can be broken down into a spectrum consisting of discrete lines of specific frequencies, or colors
Atomic Emissions Spectra is unique for each element
Example: Neon signs

21. 5.2 Quantum Theory and the Atom
Bohr Model of the Atom
proposed quantum or planetary model of the atom
Correctly predicted frequencies of the lines in H atomic emission spectrum
Related H atomic energy states to motion of e-
assigned Quantum numbers n to each orbit (n = energy level)
Known as the principle quantum number
n = 1, 2, 3, 4, 5, 6, 7 (numbers correspond to energy level as well as quantum numbers)

23. Problems Electrons do not orbit in circular paths
Bohr’s model only worked for hydrogen
Electrons do not orbit in circular paths

24. Louis de Broglie Accounted for the fixed energy level of Bohr’s model
Proposed/thought that if waves have particle-like behavior then particles can have wave-like behavior.
De Broglie Equation
Predicts all moving particles have wave characteristics
Example: cars (have wavelengths to small to be seen even with sensitive equipment)

25. Werner Heisenberg
Heisenberg Uncertainty principle  it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.
We are certain about being uncertain
Bumping into an electron while trying to determine its position and movement transfers energy and disrupts the electron

26. Erwin Schrödinger Treated hydrogen atoms electrons as a wave
Applied well to other elements
Led to quantum mechanical model of an atom (atomic model in which electrons are treated as waves)
Wave equation was able to predict three-dimensional regions of an e- probable location (ATOMIC ORBITALS-electron cloud, fuzzy cloud)

27. Atomic Orbitals
Principle Quantum numbers indicate the relative sizes and energies of atomic orbitals
The atom’s major energy levels are called principal energy levels
Principal energy levels contain sublevels
The number of energy sublevels in a principal energy level increases as the quantum number increases

28. Principal Quantum Number n
Corresponds to energy level and contains 2n2 electrons
n = 1 contains 2e- [2(1)2]
n = 2 contains 8e- [2(2)2]
n = 3 contains 18e- [2(3)2]
n = 4 contains 32e- [2(4)2]
n = 5 contains 50e- [2(5)2]
n = 6 contains 72e- [2(6)2]
n = 7 contains 98e- [2(7)2

29. Energy Levels Contain Sub-Levels
Known as Quantum number l
s,p,d,f (Sublevels)
s = 0, p = 1, d = 2, f = 3
Number of orbitals in a sub-level = (2l +1)
Number of electrons in a sub-level = 2(2l +1)

31. Electron Configuration
Electron configuration-arrangement of electrons in an atom, usually lowest amount of energy is more stable.
Three Rules or principles that define how electrons can be arranged
The Aufbau Principle
The Pauli Exclusion Principle
Hund’s Rule

32. The Aufbau Principle
States that each electron occupies the lowest energy orbital available

33. Pauli Exclusion Principle
States that a maximum of 2 electrons may occupy a single orbital, but only if the electrons have opposite spins ↑↓

34. Hund’s Rule
States that a single electron with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital

35. Orbital Diagrams and Electron Configuration Notation
Orbital Diagram Method
Electron Configuration Notation
Noble Gas Configuration

36. Valence Electrons Electrons in the outer most orbitals
Determines the chemical properties of an element
Generally the electrons in the highest energy level
Elections involved in bonding

37. Electron-Dot Structures
element symbol that represents atomic nucleus and inner electrons surrounded by dots which represents valence electrons