1. CHAPTER 5: ELECTRONS IN ATOMS
Rutherford’s model of the atom did not explain how the electrons occupy the space surrounding the nucleus.
Recall:
Opposites attract….right??
In the early 1900’s, investigations using the relationships that exist between light and electrons will reveal a more accurate model of the atom.
Model of
The Atom

2. Emission Spectrum
atoms can exist in low energy states called ground states or at high energy states called excited states.
when excited atoms return to ground states, it gives off energy in the form of electromagnetic radiation.
all elements have a characteristic emission spectra that corresponds to the amount of energy given off.
Ground &
Excited
States
Emission
Spectra

3. Probability of Finding an Electron
in a Hydrogen Atom.

4. The Bohr Model ---it’s really not boring
Niels Bohr linked photon emission with an atom’s electrons.
electrons exist in fixed orbits and never in between.
his model only explains the emission of Hydrogen atoms.
Niels
Bohr
Flame Test
Emission
Flame Test
Explanation
The Quantum Model of the Atom
Louis deBroglie suggested that electrons be considered waves confined to the space around the nucleus.

5. electrons, like waves, are capable of interference.
investigations also confirmed that electrons can be bent or diffracted.
electrons, like waves, are capable of interference.
Wave Properties
Of Electrons
Heisenberg Uncertainty Principle: it is impossible to determine simultaneously both the position and velocity of an electron.
Heisenberg
Erwin Schrodinger’s wave equation and Heisenberg’s Uncertainty Principle laid the foundation for the quantum model of the atom.
electrons do not travel in specified orbits.
instead electrons exist in regions called orbitals.
Quantum Model
of the Atom

6. Electron Configurations
“The arrangement of electrons in an atom”
Rules Governing Electron Configurations
Aufbau Principle: an electron occupies the lowest-energy orbital that can receive it.
2. Pauli Exclusion Principle: no two electrons in the same atom can have the same set of four quantum numbers.
Quantum
Subshells

7. 3. Hund’s Rule: orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron.
all electrons in singly occupied orbitals must have the same spin.
Orbital Notation
a way to illustrate what orbitals the electrons occupy in an atom.
ie: Nitrogen – 7 electrons

8. Electron Configurations
electrons configurations (shorthand method) are a method of noting how many electrons are in an atom and where they are located.
Order of Filling
Maximum amount of electrons:
s-orbital  2 electrons
p-orbital  6 electrons
d-orbital  10 electrons
f-orbital  14 electrons
Quantum
Subshells

9. Example:
Determine the electron configuration for the following elements:
Lithium  3 electrons
1s22s1
Nitrogen  7 electrons
1s22s22p3
Neon  10 electrons
1s22s22p6

10. Noble Gas Notation
in order to write long configurations much easier we can use a noble gas configuration.
In order to write a noble gas notation simply include the previous noble gas (in brackets) and list the electron configuration from the next element from the noble gas to the element you are determining the configuration for.
Write the noble gas notation for Sodium
[Ne]3s¹
Write the noble gas notation for Sulfur
[Ne]3s²3p⁴

11. Light behaves both as a wave and a particle.
Properties of Light
Light behaves both as a wave and a particle.
Light’s Wavelike Properties
visible light is one kind of electromagnetic radiation.
Light as
a Wave
Electromagnetic Spectrum
Prism

12. all electromagnetic radiation moves at the speed of light (3
all electromagnetic radiation moves at the speed of light (3.00 x 108 m/s)
Properties of Waves
Wavelength: (l) the distance between corresponding points on adjacent waves.
2. Frequency: (n) the number of waves that pass a given point in one second.
Frequency is measured in Hertz (Hz).

13. Speed of Light Equation
What is the wavelength of radiation with a frequency of
1.5 x 1013 Hz?
Wavelength x 1.50 x 1013Hz = 3.00 x 108m/s
Wavelength = 3.00 x 108m/s / 1.50 x 1013Hz
Wavelength = 2.00 x 10-5m

14. The Photoelectric Effect
a phenomena involved with the emission of electrons from a metal when light shines on a metal.
light of only a certain frequency enabled this event to occur  light behaves like streams of particles!
Photoelectric
Effect
Max Planck
suggested that energy could only exist in discrete lumps called quantum.
E = h n
h = x Js
Planck and
Blackbody
Radiation
Einstein also proposed that light delivers its energy in chunks called photons.
Therefore, in order for an electron to be ejected (photoelectric effect) , the electron must be struck by a single photon, of certain energy, to knock the electron loose.

15. One more question……….Are You Ready?
The threshold photoelectric effect in
tungsten is produced by light of a
wavelength 260 nm. Give the energy
of a photon of this light in joules.
Find frequency!
λ x ν = C
260 nm x 1m/109 nm = 2.6 x 10-7 m
2.6 x 10-7 m x ν = 3.0 x 108 m/s
v = 1.15 x 1015/sec
E = h x ν
E = x 10-34Js x 1.15 x 1015/sec
E = 7.6 x J