1. Unanswered Questions
Rutherford’s model did not address the following questions:
What is the arrangement of electrons in the atom?
What keeps the electrons from converging on the nucleus?
What accounts for the differences in chemical behavior among elements?

2. Bohr Model of the Atom Lowest allowable energy level = ground state
Electrons move around the nucleus in only certain allowable circular orbit
Each orbit is associated with a certain amount of energy
The larger the orbit (farther away from the nucleus) the greater the energy

3. Quantum Mechanical Model
Like Bohr, electrons energies are limited to a certain value
Unlike Bohr, electrons do not travel in prescribed paths.
Electrons travel within a particular volume of space surrounding the nucleus

4. Organization of Electrons in an Atom
Energy Levels n=1,2,3,4,5,6,7
Energy levels increase as distance from the nucleus increases
Sublevels (s,p,d,f)
Volume of space defined by a collection of orbitals
Orbitals
Shapes

5. Ground-State Electron Configuration
Arrangement of electron in an atom at rest
The lowest energy arrangement is the most stable
Three Rules that govern how electrons can be arranged in an atom:
Aufbau Principle (each electron occupies the lowest energy orbital available)
Pauli Exclusion Principle (A maximum of two electrons can occupy a single orbital)
Hund’s Rule (Electrons with the same spin must occupy each orbital before electrons with opposite spins can occupy each orbital)

6. Orbital Filling Sequence
(subshell)
(Energy Level)

7. Orbital Filling Diagram
Fluorine (atomic #9)
Orbital Diagram 2 2 5
Electron Configuration

9. Nobel Gas Configuration
To shorten the amount of writing required to represent the configuration of elements with large numbers of electrons, the symbol for the noble gas that directly precedes an element can represent all of the electrons up to that point.
Long Hand Short Hand
1s2 2s2 2p6 3s2 3p6 4s2
[Ar] 4s2

10. Chemical Behavior Related to the Arrangement of Electrons
In the early 1900’s scientist observed that when they heated elements in a flame, the elements emitted colored light. Analysis of the light led scientist to determine that the arrangement of electrons in an atom is related to the elements chemical behavior.

11. Line Spectra

12. Particle Nature of Light
Concept developed to explain why only certain frequencies of light are given off by heated objects and why metals will eject electrons from their surface when certain frequencies of light is shine on them.

13. Max Planck and the Quantum
By examining the light emitted by glowing objects, Planck concluded that matter can gain and lose energy in only small specific amounts called quanta
Quantum – the minimum amount of energy that can be gained or lost by an atom.
Temperature
Energy
Temperature
Energy

14. Planck Continued Equantum=hv
E = energy
h = Planck’s constant=6.626 x 10-34J
V = frequency
Relationships between wavelength, frequency and Energy
Energy
Energy
Frequency
Wavelength
Wavelength
Frequency

15. Photoelectric Effect
Electrons are emitted from a metal’s surface when light of a certain minimum frequency of 1.14 x 1015 Hz shines on the surface.
(Conversion of light energy to electrical energy)

16. Heisenberg Uncertainty Principle
Premise: It is impossible to make any measurement on an object without disturbing the object.
The position of electron can be determined by shooting photons at electrons in an atom
It is impossible to know precisely both the velocity and position of a particle at the same time

17. Einstein’s Contribution
Explained that electromagnetic radiation behaves both like waves and particles.
Ephoton=hv
Photon – a particle of electromagnetic radiation with no mass that carries a quantum of energy.

18. Electrons as Waves (Louis de Broglie)
Each energy level represents multiples of a whole wavelength
de Broglie proposed the idea that all particles have wavelengths

19. Understanding Light – Characteristics of Waves
Using the word bank below, match the words to the lettered parts of the diagram A D C
B = 3/s
1 sec E
Frequency Amplitude Peak Trough Wavelength

20. Waves Characteristics
Wavelength (l) the shortest distance between two equivalent points on a wave.
(unit of measure: meters, centimeters, or nanometers)
Frequency (v) the number of waves that pass a given point per second.
(unit of measure: Hz, waves/second, /s, s-1) c l v
Waves
c=lv
Speed of Light ( c )

21. Relationship Between Wavelength and Frequency
c=lv
Frequency
Wavelength

22. Electromagnetic Spectrum
Electromagnetic radiation – Form of energy that exhibits wave-like behavior.
Electromagnetic spectrum – encompasses all forms of electromagnetic radiation. The only differences between each form is their characteristic wavelengths and frequencies.
White Light – a continuous spectrum of colors, each with a unique wavelength and frequency.

23. Which of the following waves has the longest wavelength
Which of the following waves has the longest wavelength? Which has the greatest frequency?
5.1 Quiz tomorrow, Atomic Emission Spectroscopy Lab on Thursday

24. Problems
What is the wavelength of a wave that has a frequency of 7.8 x 1010Hz? What is the frequency of a wave that has a wavelength of 6.5 x 10-4cm?