1. Chapter 5
Electrons In Atoms

2. Revising the Atomic Model
5.1

3. Energy Levels in atoms
After discovering the nucleus, Rutherford used existing information about the atom to create the nuclear model of the atom (electrons rotate around the nucleus similar to how planets orbit the sun)

4. Limitations to Rutherford’s ATOMIC MODEL
Rutherford’s model could not explain the chemical behavior of elements
Example: why metals or compounds of metals give off characteristic colors when heated
The explanation of what leads to the chemical properties of elements required a model that would show the behavior of electrons in atoms

5. The Bohr Model
Neils Bohr was a Danish physicist and a student of Rutherford
In he developed a new atomic model that incorporated how the energy of an atom changes when the atom absorbs or emits light

6. Key Question What did Bohr propose in his model of the atom
Key Question What did Bohr propose in his model of the atom? That an electron is found only in specific circular paths, or orbits, around the nucleus.

7. The bohr model-continued
The electron orbits in Bohr’s model has a fixed energy
The fixed energies an electron can have are called energy levels
The energy levels increase from bottom to top
Electrons can move from one energy level to another
Electrons cannot exist between energy levels
To move from one energy level to another the electron must gain or lose the right amount of energy

8. These energy levels are not equally spaced
A quantum of energy is the right amount of energy required to move an electron from one energy level to another; therefore, the energy of an electron is said to be quantized
The amount of energy that an electron gains or loses in an atom may vary, it is not always the same
These energy levels are not equally spaced
The higher energy levels are closer together
The higher the energy level occupied by an electron, the less energy it takes the electron to move from that energy level to the next higher energy level
Disadvantage
Bohr only studied the simplest atom (Hydrogen); therefore, his model did not explain the energies absorbed and emitted by atoms with more than one electron

9. The Quantum mechanical model
Rutherford and Bohr both related the movement of electrons to a large moving object
However, later calculations and experimental results proved otherwise
Erwin Schrodinger, an Austrian physicist, used calculations and results to devise and solve a mathematical equation that describe the movement of an electron in the hydrogen atom
The quantum mechanical model came from the mathematical solutions to the Schrodinger Equation, which is the modern description of the electrons in atoms

10. The Quantum mechanical model
Similarities
Both the Bohr and the Quantum Mechanical Model restricts the energy of electrons to certain values
Differences
The QMM does not specify an exact path the electron takes around the nucleus

11. Key Question What does the QMM determine about the electrons in an atom? It determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus of an atom.

12. The Quantum mechanical model
Concerning the QMM, probability describes how likely it is to find an electron in a particular area surrounding the nucleus
The probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloudlike region
The cloud is more dense where there is a high probability of finding an electron, and less dense where there is a low probability of finding an electron
There is no outer boundary to the cloud because there is a slight chance of finding the electron at a considerable distance from the nucleus (opposites attract)
Therefore, attempts to show probabilities as a fuzzy cloud are usually limited to the volume in which the electron is found 90 percent of the time
It is impossible to know the exact location of an electron at any given time
The fuzzy cloud is what you have previously learned as the electron cloud. Imagine stepping inside of the electron cloud and trying to get as close as you can to an electron…that is what we are doing on the next couple of slides (the probability of finding an electron!).

13. Atomic Orbitals
Solutions to the Schrodinger equation give the energies, or energy levels, an electron can have
The Schrodinger equation leads to a mathematical expression called an atomic orbital
An atomic orbital is a mathematical representation that describes the probability of finding an electron at various locations around the nucleus; which is represented as a region where there is a high probability of finding an electron

14. Atomic Orbitals-Continued
The energy levels of electrons in the QMM are labeled by principal quantum numbers (n), which are assigned n = 1, 2, 3, 4, and so on
The principal energy levels that are higher than 1 have several orbitals with different shapes and at different energy levels
These energy levels within a principal energy level constitute energy sublevels

15. Key Question How do sublevels of principal energy levels differ
Key Question How do sublevels of principal energy levels differ? Each energy sublevel corresponds to one or more orbitals of different shapes. The orbitals describe where an electron is likely to be found.

16. Atomic Orbitals-Continued
Different atomic orbitals are distinguished by letters
s orbitals are spherical in shape
The probability of finding an electron in this orbital does not depend on direction because it is spherical in shape
p orbitals are dumbbell-shaped
These orbital have different orientations in space
px, py,pz
d orbital has five orbitals, four of the five are clover-leaf shaped
f orbitals are more complicated
See page 131

17. Atomic Orbitals-Continued
The number and types of atomic orbitals depend on the principal energy level
The number of principal energy levels equals the number of sublevels in that principal energy level
The number of orbitals in a principal energy level is equal to n2
Only 2 electrons can occupy an orbital
The total number of electrons that can occupy a principal energy level is given by the formula 2n2

18. Atomic Orbitals-Continued
Table 5.1 page 132 Summary of Principal Energy Levels and Sublevels

19. Electron Arrangement In atoms
5.2

20. Aufbau Diagrams
This diagram may not always be provided for you; therefore, you must know the shorthand way to create this diagram.

21. Electron COnfigurations
Key Question What are the three rules for writing electron configurations of elements? Aufbau principle, Pauli exclusion principle, and Hund’s rule are used to determine electron configurations of atoms.

22. AUFBAU pRINCIPLE
Aufbau principle states that electrons occupy the orbitals of lowest energy first
Orbitals for any sublevel of a principal energy level are always of equal energy
The s sublevel is always the lowest in energy within a principal energy level
The range of energy levels within a principal energy level can overlap the energy levels of another principal energy level
Aufbau Diagram
Each box represents an orbital
The energy increases from the bottom to the top
See Section 5.2 of textbook to see the Aufbau Diagram.

23. Pauli Exclusion Principle
Pauli exclusion principle states that an atomic orbital may describe at most two electrons
The two electrons must have opposite spins (paired spins)
Spin is a quantum mechanical property of electrons and may be thought of as clockwise or counterclockwise ( or )
Paired spins are represented as

24. Hund’s Rule
Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible
Example: Diagram on Board
(Will be shown in class!)
The electrons will then occupy each orbital so that their spins are paired with the first electron in the orbital

25. Writing Electron configurations
Extended Electron Configurations The sublevels within the same principal energy level are written together Model N 1s22s22p3 Sc 1s22s22p63s23p63d14s2