1. Electrons in Atoms Big Idea #2 Electrons and the Structure of Atoms
Chemistry I Notes Ch. 5
Electrons in Atoms
Big Idea #2
Electrons and the Structure of Atoms

2. 5.1 Revising the Atomic Model
The Bohr Model of the Atom
Electrons occupy specific energy levels in atoms. These energy levels are identified by a principle quantum number – n (1-7)
When electrons occupy the lowest energy levels available the atom is in its ground state.
When electrons occupy higher energy levels than the ground state the atom is in an excited state.
When atoms absorb energy the electrons jump to higher energy levels (excited state). When they drop back down to the ground state these atoms emits light of specific wavelength corresponding to the energy change of the electron (emission spectra)

3. 5.1 Revising the Atomic Model
Louis De Broglie (1924) proposed that particles of matter have wavelike character
Heisenberg’s Uncertainty Principle – The position and speed of an electron cannot be measured simultaneously. If you measure one you can’t know the other.
So electrons are located based on the probability
of finding them in a region of space around an atom (orbital).

4. 5.2 Electron Arrangement in Atoms
Quantum Mechanical Model – Treats the electron as a wave with quantized energy and explains the behavior of atoms.
Electrons occupy a cloud where the probability of finding an electron is highest where the cloud is densest.
Orbitals – a region around the nucleus where an electron with a given amount of energy is likely to be found. Maximum capacity is 2 electrons.

5. Orbitals and Energy
Electrons are found in primary energy levels sublevels and orbitals
Primary energy level – quantum number n where n = 1-7 (floor of hotel)
Sublevels – Each primary energy level has the number of sublevels equal to its n quantum number they are labeled s, p, d, & f (wing of hotel)
s sublevels contain only 1 s orbital that can hold 2 electrons
p sublevels contain 3 p orbitals that can hold 6 electrons
d sublevels contain 5 d orbitals that can hold 10 electrons
f sublevels contain 7 f orbitals that can hold 14 electrons

6. 1s Orbital

7. Representation of the 4f Orbitals in Terms of Their Boundary Surfaces

8. The Boundary Surfaces of All of the 3d Orbitals

9. The Boundary Surface Representations of All Three 2p Orbitals

10. Orbitals and Energy cont…
Orbitals contain only 2 electrons
s orbital – spherical shape 1 per energy level
p orbital – dumbell shaped 3 per energy level
d orbital – double dumbell shaped 5 per energy level
f orbital – very complex shapes 7 per energy level
Electrons within an orbital have opposite spin – Pauli Exclusion Principle

11. Orbital Energies

12. Electron Capacities of Shells and Orbitals
Primary energy Level 1 2 3 4 5 6 7
Orbital Type s
s p
s p d
s p d f
Electron Capacity
2 6
2 6 10
Total
Electrons 8 18 32

13. Quantum Numbers and Electrons
If the primary energy level is given as n=1-7
Then the number of sublevels in that primary energy level = n
The number of orbitals in that energy level = n2
The number of electrons in that energy level= 2n2

14. 5.2 Electron Arrangement in Atoms cont…
Electron configurations are determined by distributing the atom’s electrons into energy levels, sublevels , and orbitals based on a set of principles.
Aufbau Principle –Electrons are added one at a time to the lowest energy orbitals available until all electrons are accounted for.
Pauli Exclusion Principle – Orbitals hold a maximum of 2 electrons with opposite spin.
Hund’s Rule – Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results (college dorm rule). Fill electrons into sublevels 1 per orbital until all orbitals have 1 then double up.
Remember the exceptions at Copper and Chromium

15. Energy Level Positions
For The First
Twenty Elements 1S

16. 5.3 Atomic Emission Spectra
Light has properties of both waves and particles.
Light is a form of electromagnetic radiation – electric and magnetic fields oscillating at right angles to each other.
Wave Characteristics
Amplitude – height of wave from origin to peak (m)
Wavelength -  -Distance between successive crests (m)
Frequency –  - cycles per second (Hz)
Speed – c- distance per unit time Light= 3x108 m/s
Speed = frequency x wavelength c=  

17. The Nature of Waves

18. Electromagnetic Waves
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19. Problem
Calculate the wavelength of ultra violet light with a frequency of 2.00X1016 Hz?

21. 5.3 Atomic Spectra cont… LIGHT: What Is It? Light Energy
Atoms
As atoms absorb energy, electrons jump out to a higher energy level.
Electrons release light when falling down to the lower energy level.
Photons - bundles/packets of energy released when the electrons fall.
Light: Stream of Photons

22. 5.3 Atomic Spectra cont…
Quantum – Restriction on the amount of energy an object emits or absorbs (Max Planck)
E= h where
E is the energy of a photon in joules (j)
h is Planck’s constant 6.62x10-34 j-s
 is the frequency in hertz (Hz) or 1/s
Energies absorbed or emitted by atoms are quantized
Photoelectric Effect- electrons ejected from the surface of a metal when light shines on it.
For every metal there is a minimum frequency of light needed to release electrons. This quantum of energy is called a photon. (Albert Einstein)
Thus light has particle wave duality

23. Problem
What is the energy of a photon of Infrared light with a wavelength of 1.00x10-5 m ?
First calculate frequency using c=λv then calculate energy using E=hv

24. 5.3 Atomic Spectra cont… Types of spectra
Line Spectra – contain only specific wavelengths or colors – atoms emit these when their electrons are excited.
Continuous Spectrum – white light broken into all its colors R-O-Y-G-B-I-V (rainbow)
Absorption spectra – Black lines on the continuous spectra of star light that represent fingerprints of the elements present.

25. Refraction of White Light

26. The Line Spectrum of Hydrogen