1. Chapter 5 Electrons in Atoms
Mr. Samaniego
Lawndale High School

2. Summary of Atomic Theory
Year
Event
400BC
Democritus proposes idea of atom
1808
Dalton develops Atomic Theory
1897
Thomson uses cathode ray to discover electron
1916
Millikan measures the mass of an e-
1919
Rutherford uses gold foil experiment to discover nucleus

3. Structure Of An Atom
So by this point, we know that protons and neutrons are located in the nucleus and electrons are around the outside of the nucleus

4. Section 5.1 – Models of the Atom
In 1897 J. J. Thomson discovered the electron
Observed that a magnet deflected the straight paths of the cathode rays

5. Atoms were known to be electrically neutral which meant that there had to be some positively charged matter to balance the negative charges

6. Ernest Rutherford’s experiment disproved the plum pudding model of the atom and suggested that there was a positively charged nucleus because most of the alpha particles went straight through the gold foil
BUT, Rutherford’s atomic model could not explain the chemical properties of elements

7. The Bohr Model (Niels Bohr)
In 1913, Niels Bohr came up with a new model (Bohr was a student of Rutherford)
He noticed that light given out when atoms were heated always had specific amounts of energy, so Niels Bohr proposed a model that electrons in an atom must be orbiting the nucleus and can reside only in fixed energy levels

8. Energy Levels
Each of the electrons in Bohr’s model has a fixed amount of energy called energy levels
This is similar to steps of a ladder (can climb up the ladder, but cannot step in between the steps)
Quantum is the amount of energy required to move an electron from one energy level to another
The further away from the nucleus, the more energy the electron has

10. The Quantum Mechanics Model (Erwin Schrodinger)
While Rutherford’s model described the path the electron moves, Erwin Schrodinger solved mathematical equations to describe the behavior of electron
Similar to Bohr’s model, Schrodinger describes the energy of electrons with certain values but does not involve an exact path the electron takes around the nucleus

11. The Quantum Mechanics View of the Atom (Schrodinger)
The quantum mechanical model does not describe the exact path an electron takes around the nucleus, but determines the probability of finding an electron in a certain area

12. Quantum Mechanical Model
In this model, electrons move similar to a rotating propeller blade
You cannot tell its precise location at any instant because it’s a blurry region, but you have information regarding the probability of finding an electron within a certain volume of space
Similar to a fuzzy cloud…the probability of finding an electron is higher where the cloud is more dense

13. Atomic Orbitals
An Atomic Orbital is a region of space where there is a high probability of finding an electron
Each energy sublevel corresponds to an orbital of different shape describing where the electron is likely to be found (there are 4 different types of shapes)

14. Shapes of Orbitals

15. Shapes of Orbitals

16. Shapes of Orbitals

17. Chapter 5.2 – Electron Arrangement in Atoms
Each electron in an atom is assigned a set of four quantum numbers. These help to determine the highest probability of finding the electrons.
Three of these numbers (n, l, m) give the location of the electron
The fourth (s) describes the orientation of an electron in an orbital.

18. Quantum letters can be thought of like the numbers and letters on a concert ticket

19. Labeling Electrons in Atoms
Probable Location of electron
Probability
Probable location of Finding Beyonce or Taylor
Energy level (n)
High Probability
Hotel Name
Sublevel (l)
Higher Probability
Hotel Floor Number
Orbital (m)
Highest probability
Hotel Room Number

20. Electron Configuration
Sublevel
# of Orbitals Available
# of Electrons Available s 1 2 p 3 6 d 5 10 f 7 14

21. Electron Configurations
Filling order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Example He = 2 electrons
1s2
Example Li = 3 electrons
1s22s1
Example B = 5 electrons
1s22s22p1

23. Practice Problems Li 5. P N 6. Si Be 7. Mg C 8. Al
Write electron configurations for the following atoms
Li P
N Si
Be Mg
C Al

24. Electron Configurations can be written in terms of noble gases
To save space, configurations can be written in terms of noble gases
Example 1: Ne = 1s22s22p6
S = 1s22s22p63s23p4
Or S = [Ne] 3s23p4
Example 2: Ar = 1s22s22p63s23p6
Mn = 1s22s22p63s23p64s23d5
Mn = [Ar] s23d5

25. Reading the Periodic Table

26. Locating Electrons in Atoms
So far we have discussed 3 quantum numbers
n= principal quantum level (principal energy level)
l= Sublevel
m = magnetic quantum number (shape of orbitals)
1s2 n
Number of electrons in sublevel l

27. s = spin When an electron moves, it generates a magnetic field.
s describes the direction an electron spins
They must spin in opposite directions
Spin= up down
There are two values of s: +1/2 and -1/2

28. Orbital Diagrams
The electron configuration gives the number of electrons in each sublevel but does not show how the orbital of a sublevel are occupied by the electrons

29. Orbital Diagrams
Used to show how electrons are distributed within sublevels
Each orbital is represented by a box and each electron is represented by an arrow
Notice that each box is drawn higher than the last set because it has more energy
Example: Boron 1s22s22p1 2p 2s 1s

30. Orbital Diagrams 2p 2s 1s 2p 2s 1s
Steps to writing orbital diagrams:ex F (Z=9)
Write the electron configuration
1s22s22p5
2. Construct an orbital filling diagram using boxes for each orbital
3. Use arrows to represent the electrons in each orbital. 2p 2s 1s 2p 2s 1s

31. Rule #1 - Aufbau Principle
Electrons must occupy the orbital with the lowest energy first
Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

32. Rule #2 - Pauli Exclusion Principle
Orbitals can only have two electrons max
The 2 electrons must have opposite spins
Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

33. Rule #3 - Hund’s Rule
Orbitals of equal energy are each occupied by one electron before any pairing occurs
Example: Oxygen 1s22s22p4 2p 2p 2s 2s 1s 1s

34. Draw orbital diagrams for these elements
Li P
N Si
Be Mg
C Al

35. Section 5.3 - Atomic Spectra
When atoms absorb energy, electrons move into higher energy levels (excited state)
When these electrons return to their lower energy levels, they lose energy by emitting light
Atomic Emission Spectrum – the discrete lines representing the frequencies of light emitted by an element

36. Calculating Wavelength of Light
c = speed of light (3 x 108 m/s2)
 = wavelength (called lambda)
 = frequency (called nu)

37. Practice
1. Calculate the wavelength of a yellow light if the frequency is 5.10 x 1014 sec-1 or Hz.
Answer = 5.88 x 10-7m
2. What is the wavelength of 1.50 x 1013 sec-1?
Answer = 2.00 x 10-5m
3. What frequency is radiation with a wavelength of 5.00 x 10-8m?
Answer = 6.00 x 1015 sec-1 or Hertz

38. Atomic Spectra
Each discrete line in an emission spectrum corresponds to one exact frequency of light emitted by the atom
Ground State – lowest possible energy of the electron in the Bohr model
The light emitted by an electron moving from higher to a lower energy level has a frequency directly proportional to the energy change of the electron

39. Chapter 5 Assessment Page 148
Homework
Chapter 5 Assessment Page 148
#’s 22-24, 27, 29, 30-39,
50-53, 57, 60, 68, 70-72