1. Chapter 11 - Arrangement of Electrons in Atoms
The “Puzzle” of the nucleus:
• Protons and electrons are attracted to each other because of opposite charges
• Electrically charged particles moving in a curved path give off energy
• Despite these facts, atoms don’t collapse

2. 11-1 The Development of a New Atomic Model
I. Properties of Light
A. Electromagnetic Radiation
1. Many types of EM waves
a. visible light
b. x-rays
c. ultraviolet light
d. infrared light
e. radio waves

3. 2. EM radiation are forms of energy which move through space as waves
a. Move at speed of light (c)
3.00 x 108 m/s
b. Speed is equal to the frequency times the wavelength
c = νλ
Frequency (ν) (greek letter – nu) is the number of waves passing a given point in one second
Wavelength (λ) (greek letter - lambda) is the distance between peaks of adjacent waves
c. Speed of light (c) is a constant, so νλ is also a constant
ν (nu) and λ (lambda) must be inversely proportional

5. B. Light and Energy - The Photoelectric Effect
a. Electrons are emitted from a metal when light shines on the metal
Increased light intensity = more electrons ejected
Higher frequency light = faster electrons being ejected

6. 2. Radiant energy is transferred in units (or quanta) of energy called photons
a. A photon is a particle of energy having a rest mass of zero and carrying a quantum of energy
b. A quantum is the minimum amount of energy that can be lost or gained by an atom

7. a. E = hν (h is Planck’s constant, 6.62554 x 10 -34 J sec)
3. Radiant Energy (E) of a photon is directly proportional to the frequency (ν) of radiation
a. E = hν (h is Planck’s constant, x J sec)
Wavelength
(m)
gamma
X-ray UV
Visible light
infrared
microwaves
radiowaves FM
short AM
VIBGYOR
(s-1)
frequency
Wavelength increases
Frequency decreases
Energy decreases

8. 4. Wave-Particle Duality
a. Energy travels through space as waves, but can be thought of as a stream of particles (Einstein)

9. Read Sections and Do
11.6 pg
Questions -5,6 pg , pg 267
11.7 pg
Questions -7 pg 260 , 29 pg 267
Small Scale Labs
16 - Design and construct a quantitative electroscope
17 – Visible Spectra and the Nature of Light Color

10. II. The Hydrogen Line Spectrum
A. Ground State
1. The lowest energy state of an atom
B. Excited State
1. A state in which an atom has a higher potential energy than in its ground state

11. C. Bright line spectrum
1. Light is given off by excited atoms as they return to lower energy states
2. Light is given off in very definite wavelengths
3. A spectroscope reveals lines of particular colors
410nm 434nm nm nm
a. Definite frequency
b. Definite wavelength

12. Change in Energy (E) = Planck’s Constant(h) x Frequency (ν)
=6.626 x10-34 J s x x 1015 s-1
=3.31 x 10-18J
Frequency (ν ) = Speed of Light (c) / Wavelength(λ)
Therefore
ΔE = hv = h x c λ

13. III. The Bohr Model of the Atom
A. Electron Orbits, or Energy Levels
1. Electrons can circle the nucleus only in allowed paths or orbits
2. The energy of the electron is greater when it is in orbits farther from the nucleus
3. The atom achieves the ground state when atoms occupy the closest possible positions around the nucleus
4. Electromagnetic radiation is emitted when electrons move closer to the nucleus

14. Ultraviolet
Visible
Infrared

15. B. Energy transitions
1. Energies of atoms are fixed and definite quantities
2. Energy transitions occur in jumps of discrete amounts of energy
3. Electrons only lose energy when they move to a lower energy state
C. Shortcomings of the Bohr Model
1. Doesn't work for atoms larger than hydrogen (more than one electron)
2. Doesn't explain chemical behavior
Replaced by a better model

16. Read 11.9 Pgs 262-263 Questions 8 pg 261 9 pg 264 33 pg 267
Worksheet 11 B

17. The Quantum Model of the Atom
I. Electrons as Waves and Particles
A. Louis deBroglie (1924)
1. Electrons have wavelike properties
2. Consider the electron as a wave confined to a space that can have only certain frequencies

18. B. The Heisenberg Uncertainty Principle
(Werner Heisenberg )
1. "It is impossible to determine simultaneously both the position and velocity of an electron or any other particle
a. Electrons are located by their interactions with photons
b. Electrons and photons have similar energies
c. Interaction between a photon and an electron knocks the electron off of its course

19. C. The Schroedinger Wave Equation
1. Proved quantization of electron energies and is the basis for Quantum Theory
a. Quantum theory describes mathematically the wave properties of electrons and other very small particles
2. Electrons do not move around the nucleus in "planetary orbits"
3. Electrons exist in regions called orbitals
a. An orbital is a three-dimensional region around the nucleus that indicates the probable location of an electron

20. Read Pages
Questions
34 pg 267

21. II. Atomic Orbitals and Quantum Numbers
Quantum Numbers specify the properties of atomic orbitals and the properties of the electrons in orbitals
Principal Quantum Number (n)
1. Indicates the main energy levels occupied by the electron (1-7)
2. Values of n are positive integers
a. n=1 is closest to the nucleus, and lowest in energy
b. The principal quantum number (n) always equals the
number of sublevels within that principal energy level
c,. The maximum number of electrons that can occupy an
energy level is given by the formula 2n2, where n = the
principal quantum number (1-7)

22. B. Angular Momentum Quantum Number (l)
Indicates the shape of the orbital
2. Number of orbital shapes = n
a. Shapes are designated s, p, d, f
sharp, principal, diffuse, and fundamental.
Actual 3 D spaces where electrons can be found. The principal quantum number is at the right of each row and the azimuthal quantum number is denoted by letter at top of each column

23. C. Magnetic Quantum Number (m)
1. The orientation (shape) of the orbital around the nucleus
s -orbitals have only one possible orientation
m = circular
p -orbitals have three
m= dumbbell
d -have five
m= dumbbell and donut
f -have 7 possible orientations
m = weird

24. s orbital
px orbital py orbital pz orbital
dxy orbital dxz orbital dyz orbital dx2-y2 orbital dz2 orbital

25. 1 2 8 18 4 32 Principal Quantum Number (n)
Sublevels in main energy level
(n sublevels)
Number of orbitals per sublevel
Number of electrons per sublevel
Number of electrons per main energy level (2n2) 1 s 2 p 3 6 8 d 5 10 18 4 S P f 7 14 32

26. D. Spin Quantum Number (m s)
1. Indicates the fundamental spin states of an electron in an orbital
2. Two possible values for spin, +1/2, -1/2
3. A single orbital can contain only two electrons, which must have opposite spins

27. Read 11.3 Pgs
Questions
1 pg 249
14 pg 266

28. Electron Configurations
I. Writing Electrons Configurations
A. Rules
1. Aufbau Principle
a. An electron occupies the lowest-energy orbital that can receive it

29. 2. Pauli Exclusion Principle
a. No two electrons in the same atom can have the same set of four quantum numbers or an atomic orbital can contain at most 2 electrons.

30. 3. Hund's Rule
a. Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron, and all electrons in singly occupied orbitals must have the same spin
____ ____ ____ ____ ____ ____ ____ ____ ____
2p 2p 2p

31. B. Electron Orbital Notation
1. Unoccupied orbitals are represented by a line, _____
a. Lines are labelled with the principal quantum number and the sublevel letter see page 251
2. Arrows are used to represent electrons
a. Arrows pointing up and down indicate opposite spins see page 252

32. 1s2 C. Electron Configuration Notation
1. The number of electrons in a sublevel is indicated by adding a superscript to the sublevel designation
Hydrogen = 1s1
Helium = 1s2
Lithium = 1s22s1
Use Periodic table method to identify electron configuration
1s2
# of Electrons
Principle Quantum
Number
Sublevel

33. II. Survey of the Periodic Table
A. Elements of the Second and Third Periods
1. Highest occupied energy level
a. The electron containing energy level with the highest principal quantum number
2. Inner shell electrons
a. Electrons that are not in the highest energy level

34. 3. Octet
a. Highest energy level s and p electrons are filled (8 electrons)
b. Characteristic of noble gases, Group 18 (s2p6)
4. Noble gas configuration
a. Outer main energy level fully occupied, usually (except for He) by eight electrons (s2p6)
b. This configuration has extra stability

35. B. Elements of the Fourth Period
1. Irregularity of Chromium
a. Expected: 1s22s22p63s23p64s23d4
b. Actual: 1s22s22p63s23p64s13d5
2. Several transition and rare-earth elements borrow from smaller sublevels in order to half fill larger sublevels
Configurations with either d4 or d9 are the exceptions
And will steal one electron from an adjacent s sublevel
To fill the d sublevel
½ filled and full sublevels are more stable

36. 1) Read 11.4 and 11.5 pg pg
Worksheet -8 Electron Configurations and Periodicity
2) Read 11.1 and 11.2 pg
Questions
10-13 pg 266
3) Worksheet 11A

37. Objective Worksheet and Review Sheet