1. Chapter 5 Notes
Electrons

2. Chapter 5 Section 1
Models of the Atom

3. Bohr Model
Bohr proposed that electrons were found in specific circular paths around the nucleus called orbits.
Each possible orbit had a specific energy level.
A quantum is the amount of energy required to move from one level to another.

4. Orbitals
Atomic orbital- a region of space where it is likely to find an electron.
Energy levels are represented by principal quantum numbers (n).
Levels are made up of sublevels that contain different shaped orbitals.
The principal quantum number tells you how many sublevels make each level. n=1 has 1 sublevel, n=2 has 2

5. Orbitals There are four types of orbitals n=1 has 1s n=2 has 2s, 2p
In order of increasing energy =s, p, d, f
n=1 has 1s
n=2 has 2s, 2p
n=3 has 3s, 3p, 3d
n=4 has 4s, 4p, 4d, 4f

6. S orbitals
The s orbital is a perfect sphere and has only one possible arrangement.

7. The p orbital is a dumbell shape and can have 3 possible arrangements.

8. Orbital arrangements s=1, p=3, d=5, f=7
Each orbital can hold two electrons
Each energy level can hold a different number of electrons.
The formula 2n2 allows you to find the number of electrons at each level

9. Electron Arrangement in Atoms
Chapter 5 Section 2
Electron Arrangement in Atoms

10. Electron Configuration
Three rules tell you how to find the electron configuration of atoms.
Aufbau principle- Electrons occupy the orbitals of lowest energy first.
Pauli Exclusion Principle- An orbital can hold at most two electrons.
Hund’s rule- electrons will occupy orbitals to make the number of electrons with the same spin as large as possible.

11. Electron Arrangement
We will illustrate electrons with up and down half arrows and we will illustrate the orbitals with a box.
Arrows will be in opposite direction.

12. Aufbau Principle
Electrons fill orbitals from lowest to highest energy.

13. Pauli Exclusion Principle
A full orbital contains two electrons.
The paired electrons will have opposite spins. The opposite spin is illustrated with up and down half arrows. ↑↓

14. Hund’s Rule
Electrons occupy orbitals of the same energy so that the number of electrons with the same spin direction will be as large as possible.
Example – Oxygen would be 1s22s22p4. The 4 p electrons would fill the three p orbitals to allow the highest number with the same spin. Which would be ↑↓ ↑ ↑

15. Example Electron configurations

16. Writing Out Configuration
We can write out the electron configuration for elements as follows.
Hydrogen – 1s1, Helium = 1s2
Titanium – 1s22s22p63s23p64s23d2
You can abbreviate from the last group 18 element. Put that element in brackets, then write the configuration that follows.
For Titanium – [Ar] 4s23d2

17. Chapter 5 Section 3
Light

18. Light Light travels in waves. Waves travel in cycles.
A wave will start at zero, increase to its highest point, pass through zero and go to its lowest point, and return to zero.

19. Amplitude- wave’s height from zero to the crest (high point)
Wavelength (λ) – The distance between crests or cycles.
Frequency (v) – Number of wave cycles per unit of time.

20. The SI unit of frequency is hertz, which is a reciprocal second (s-1).
The frequency times the wavelength is equal to the constant c (the speed of light) 3 x 108 m/s
This is represented by the formula c= λv
Frequency and wavelength are inversely proportional to each other.

21. Light consists of electromagnetic waves
Light consists of electromagnetic waves. This includes: radio, radar micro, infrared, visible light, ultraviolet, X rays, and gamma rays.
The range of light that we see (visible light) is actually a very small part of the electromagnetic spectrum.

22. When light travels through a prism, the different frequencies of light separate into a spectrum of colors.
Each color blends into the next in the order of red, orange, yellow, green, blue, indigo, violet.
Red light has the longest wavelength and the lowest frequency.
Violet has the shortest wavelength and the highest frequency.

23. Atomic spectra
Atoms absorb energy that raises electrons into higher energy levels.
Atoms lose this energy by emitting light when electrons return to lower energy levels.
This is called an electronic transition.
Ground state- the lowest possible energy of an electron. (n=1)

24. Each line in an emission spectrum corresponds to one exact frequency of light emitted by an element.
This is the element’s fingerprint. No two elements have the same emission spectrum.
We study the atomic spectra of the stars to gain knowledge of the composition of the universe.

26. Light Light can behave as both a particle and a wave.
Heisenberg uncertainty principle- it is impossible to know both the velocity and position of a particle at the same time.