1. Electrons in Atoms

2. Evolution of Atomic Models
For about 50 years after Dalton, the atom was considered a solid indivisible mass
JJ Thomson proposed the plum-pudding model

3. Evolution of Atomic Mdels
Rutherford proposed the nuclear atom, with a dense nucleus
Niels Bohr propsed that electrons move in orbits around the nucleus

4. Evolution of Atomic Models
Bohr proposed that electrons have a fixed energy
The energy level of an electron is the region where the electron is likely moving
Energy levels are like stairs. To move to the next stair you have to move the correct amount, or you won’t move

5. Evolution of Atomic Models
A quantum of energy is the amount of energy required to move an electron to the next higher energy level
In general, the higher the energy level of the electron, the farther it is from the nucleus

6. Quantum Mechanical Model
Schrodinger’s equation describes the location and energy of electrons  quantum mechancial model
Estimates the probability of finding an electron in a certain position

7. Quantum Mechanical Model
In this model, the probability of finding an electron is represented by a fuzzy cloud
The cloud is more dense where the probability of finding electrons is high and less dense where the probability of finding electrons is low

8. Atomic Orbitals
The Q-M Model designates energy levels with principal quantum numbers (n)
They are assigned in order of increasing energy (greater distance from nucleus)
n – 1, 2, 3, 4, …..

9. Atomic Orbitals Within each energy level, there are sublevels
The number of energy sublevels is the same as the principal quantum number
Describes the SHAPE of the orbital
s, p, d, f

10. Atomic Orbitals
Within each sublevel, there are regions where electrons are likely to be found, called atomic orbitals
Atomic orbitals are denoted by letters
s is spherical in shape, p is dumbbell-shaped, d is cloverleaf-shaped, and f is a complex shape pg 103 in book

11. Atomic Orbitals
The numbers and kinds of atomic orbitals depend on the energy sublevel
n = 1: 1 sublevel – 1s (1 orbital, s)
n = 2: 2 sublevels – 2s, 2p (4 orbitals, s and 3 p)
n = 3: 3 sublevels – 3s, 3p, 3d (9 orbitals, s, 3 p, and 5 d)
n = 4: 4 sublevels – 4s, 4p, 4d, 4f (16 orbitals, s, 3 p, 5 d, and 7 f)

12. Atomic Orbitals There are a maximum of two electrons per orbital
Each sublevel also has a maximum number of orbitals
Sublevel
Max # of orbitals
Max # of electrons s 1 2 p 3 6 d 5 10 f 7 14

13. Atomic Orbitals
Because of the maximum allowed electrons per orbital, and the maximum orbitals per type of orbitals, it can be determined how many electrons are allowed in each energy level
Energy level n 1 2 3 4
Max # of electrons
2 (s)
6 (s,p) 18
(s, p, d)
32 (s,p,d,f)

14. Development of the New Atomic Model

15. Light and Atomic Spectra
Isaac Newton proposed that light was made of particles
By 1900, most scientists had accepted that light was a wave

16. Light and Atomic Spectra
According to the wave model, light consists of electromagnetic waves
Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, x-rays, and gamma rays

17. Light and Atomic Spectra
Four properties for measuring waves:
Amplitude – wave’s height from the origin to the crest
Wavelength (λ) – distance between crests

18. Light and Atomic Spectra
Frequency (ν) – the number of wave cycles to pass a given point per unit time
The units of frequency are cycles per second. The SI unit is the Hertz (Hz) which can also be expressed as s-1
Speed (c) – the rate at which the wave travels

19. What do you think the relationship is between speed (c), frequency (ν), wavelength (λ)?

20. Light and Atomic Spectra
There is a relationship between the speed, frequency, and wavelength of light
c = λν
c is constant that equals 3.0 x 108 m/s

21. Light and Atomic Spectra
When light passes through a prism, it produces a spectrum of colors.
ROYGBIV (red-orange-yellow-green-blue-indigo-violet) is the range of visible light (380 nm – 700 nm)

22. Light and Atomic Spectra
Every element emits light when it is excited by the passage of an electric discharge through its gas or favor
Passing this light through a prism results in the atomic emission spectrum of the element

23. Quantum Concept
Classical physics does not explain the line spectra of the elements
Max Planck attempted to describe why iron changes color as it is heated

24. Quantum Concept
Planck showed that the amount of radiant energy (E) absorbed or emitted is proportional to the frequency of the radiation.
E = hν (Planck’s constant, h = x Js)

25. Quantum Concept The energy absorbed or emitted is a quantum
Quanta are so small, that you are unaware that energy is quantized

27. Photoelectric Effect
Albert Einstein proposed that light could be described as quanta of energy that behave like particles, called photons
This led to the theory of the dual nature of light

28. Photoelectric Effect
In the photoelectric effect metals eject electrons when light shines on them
It depends on the frequency of the light , contrary to the ideas of classical physics

29. Atomic Spectra
Bohr’s application of quantum theory to electron energy levels in atoms resulted in an explanation of the hydrogen spectrum

30. Atomic Spectra
When an electron is in the lowest energy level it is in the ground state
When an electron is excited, it raised to higher energy level. As it comes back to ground state, energy is emitted in the form of light

31. Quantum Mechanics
Louis de Broglie proposed that matter could also have a dual nature
λ = (h/mν) (de Broglie’s equation)
Objects with measurable wavelengths cannot be seen by the unaided eye

32. Quantum Mechanics
Werner Heisenberg stated that it is impossible to know exactly both velocity and the position of a particle at the same time (Heisenberg Uncertainty Principle)

33. Electron Configurations
In atoms, electrons and the nucleus interact to make the most stable arrangement possible.
The ways in which electrons are arranged around the nucleus are called electron configurations.

34. Electron Configurations
Three rules tell you how to find electron configurations:
Aufbau Principle – Electrons enter orbitals of lowest energy first. (see diagram on page 367 of your book)

35. Electron Configurations
Pauli Exclusion Principle – An atomic orbital may have a maximum of two electrons. These electrons must have opposite spins (paired)

36. Electron Configurations
Hund’s Rule – When electrons occupy orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with parallel spins

37. Electron Configurations
Electron configurations are written in the following way:
Carbon (6 electrons) – 1s22s22p2
1, 2 = energy level
s, p = sublevel
Superscript 2 = number of electrons in the sublevel
** Note that the sum of the supersrcipts equals the number of electrons in the atom

38. Electron Configurations
Electron configurations can also be abbreviated by substituting for the electron configuration of the previous noble gas:
Carbon – 1s22s22p2
Abbreviated – [He]2s22p2
1s2 is the electron configuration for He, so the symbol He replaced its electron configuration

39. Electron Configurations
There are some exceptions to the rules, such as copper and chromium