1. Models of the atom & quantum theory

2. Niels Bohr (1913)
Previous research had concluded that light has a dual nature.
It could act as both particles and waves
Bohr proposed a model that explained why only certain frequencies of light were emitted from hydrogen.

3. Bohr’s model of the atom
Suggested that the electron in a hydrogen atom moves around the nucleus in a circular path
i.e. Like planets around a sun
Atoms can have multiple “orbits” that have a definite, fixed amount of energy
The lowest allowable energy state of an atom is the ground state
This is closest to the nucleus
The smaller the orbit, the lower the energy and vise versa
Orbits

4. Excited electrons In the ground state, atoms do not radiate energy
When energy is add from an outside source, electrons move to a higher-energy orbit
This movement raises the electron to an “excited” state
Once excited, the electron can drop from the higher- energy orbit to a lower-energy orbit
Results in the atom emitting a photon

5. Ladder analogy
Like the rungs of the ladder on the right, the energy levels in an atom are not equally spaced
The higher the energy level occupied by an electron, the less energy it takes to move from the energy level to the next highest energy level

6. Problem with bohr’s model
Bohr’s model had some limits
It could not explain the spectra of light emitted for any other element
It did not fully account for the chemical behavior of atoms

7. Electrons as waves
French scientist De Broglie pointed out the electron orbits in Bohr’s model were similar to the behavior of waves
Electrons, like waves, could be:
Diffracted (bent)
Bending a wave as it passes by the edge of an object
Interfere with each other
When waves overlap. Results in a reduction of energy in some areas and an increase of energy in others.

8. Uncertainty principle
Werner Heisenberg – 1927
Showed that it is impossible to take any measurement of an object without disturbing the object!
Detection of electrons
Detected by their interaction with photons
Have about the same mass as an electron
Found by “bumping” a photon into an electron
Doing so changes both wavelength of photon and the position and velocity of the electron
Heisenberg’s Uncertainty Principle:
It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.

10. Schrödinger's wave equations
Remember the problem with Bohr’s model?
Only relevant to hydrogen…
Erwin Schrodinger – 1926
Developed an equation that treated electrons in atoms as waves
His new model not only could be applied to hydrogen, but all the other elements as well!!! YAY
This new model, in which electrons are treated as waves is called:
The Quantum Mechanical Model of the Atom
Other names: Wave Mechanical Model or Charge Cloud Model

11. Schrödinger sound familiar?

12. Location of Electrons
Electrons described as waves only have a certain probability of being found within a particular volume of space around the nucleus
Schrödinger’s wave function predicts a 3D region around the nucleus that described the electron’s probable location
Atomic Orbital:
Description of the 3D region around the nucleus (looks like a fuzzy cloud)
Clouds show the region of probable electron locations
Size and shape of cloud depends on the energy of the electrons that occupy them

14. Activity
Probability of finding an electron around the nucleus activity

15. Electrons around the nucleus
Orbitals can differ based on size and shape
There are four types of orbitals
S, P, D and F
Each orbital can hold two electrons
An S orbital is shaped like a sphere
A 1s orbital is smaller than a 2s orbital 1s 2s

16. Orbital Shape

17. Electrons in the orbitals
Orbital type
Number of Types
Electrons that fit into this “shell” s 1
2(1) = 2 p 3
2(3) = 6 d 5
2(5) = 10 f 7
2(7) = 14
The right column tells us the total electrons that fit into all types of that one orbital

20. Rules of arrangement
Three rules, or principles define how electrons can be arranged in an atom’s orbitals
Aufbau Principle
Pauli Exclusion Principle
Hund’s Rule

21. Aufbau Principle
States: each electron occupies the lowest energy orbital available
First step in determining the ground-state electron configuration is to learn the sequence of atomic orbitals from lowest to highest energy
Each box in the Aufbau diagram represents an orbital

22. Electron configuration
Based on Aufbau’s Principle, the order of the orbitals should be: 1s – 2s – 2p – 3s – 3p – 4s – 3d – 4p – 5s – 4d – 5p – 6s… All elements follow this rule with few exceptions

23. Pauli exclusion principle
Electrons in orbitals can be represented by arrows in boxes
Each electron has an associated spin
Arrow pointing up represents electron spinning in one direction
Arrow point down represents electron spinning in opposite direction
Empty box represents an unoccupied orbital
Pauli Exclusion Principle:
A maximum of two electrons can occupy a single atomic orbital, but only if the electrons have opposite spins
= 0 e-
= 1 e-
= 2 e-

24. Hund’s Rule
Remember that negatively charged electrons repel each other
Hund’s Rule:
Single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same orbital

25. Electron configuration
The electron configuration on an atom is a shorthand method of writing the location of electrons by sublevel
The first number describes the principle energy level
The letter represents the sublevel
The sublevel is followed by a superscript with the number of electrons in the sublevel
If the 2p sublevel contains 2 electrons, it is written 2p2

26. Writing electron configurations
First, determine how many electrons are in the atom
Iron has 26 electrons
Arrange the energy sublevels to increasing energy
1s 2s 2p 3s 3p 4s 3d…
Fill each sublevel with electrons until you have used all the electrons in the atom
Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
The sum of the superscripts equals the atomic number of iron (26)

27. Example
Phosphorus, an element used in matches, has an atomic number of 15. Write the electron configuration of a phosphorus atom.
Answer:
1s2 2s2 2p6 3s2 3p3

28. Practice electron configurations
Write the complete electron configurations for: Be F S Ca
Things to double check:
All superscripts add up to total electrons
Orbital order follows the Aufbau Principle

29. answers Be: 1s2 2s2 F: 1s2 2s2 2p5 S: 1s2 2s2 2p6 3s2 3p3 Ca:
1s2 2s2 2p6 3s2 3p6 4s2