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

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

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

B. Light and Energy - The Photoelectric Effect

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

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

3. Radiant Energy (E) of a photon is directly proportional to the frequency (ν) of radiation

a. E = hν (h is Planck’s constant, 6.62554 x 10 -

34 J sec)

gamma X-ray UV Visible light infrared microwaves

FM short AM

radiowaves

VIBGYOR

Wavelength increases

Frequency decreases

Energy decreases

10-11 10-9 10-8 10-6 10-5 10-2 10-1 (m)

Wavelength

1018 1017 1015 1014 1012 1010 109 (s-1)

frequency

4. Wave-Particle Duality

a. Energy travels through space as waves, but can be thought of as a stream of particles (Einstein)

•Read Sections and Do

•11.6 pg255-257

•Questions -5,6 pg 257-258 , 24-28 pg 267

•11.7 pg 260-261

•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

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

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 486nm 656nm

a. Definite frequency

b. Definite wavelength

ΔE = hv = h x c

λ

Change in Energy (E) = Planck’s Constant(h) x Frequency (ν)

Frequency (ν ) = Speed of Light (c) / Wavelength(λ)

Therefore

=6.626 x10-34 J s x 5.00 x 1015 s-1

=3.31 x 10-18J

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

Ultraviolet

Visible

Infrared

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

•11.8 pg 261-262

•Questions -30-32 pg 267

•Read 11.9 Pgs 262-263•Questions

•8 pg 261•9 pg 264•33 pg 267

Worksheet 11 B

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

B. The Heisenberg Uncertainty Principle

(Werner Heisenberg - 1927)

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

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

•Read 11.10 Pages 264-265

•Questions•34 pg 267

II. Atomic Orbitals and Quantum Numbers

Quantum Numbers specify the properties of atomic orbitals and the properties of the electrons in orbitals

A. 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 anenergy level is given by the formula 2n2, where n = the principal quantum number (1-7)

B. Angular Momentum Quantum Number (l)

1. 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

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

px orbital py orbital pz orbital

dxy orbital dxz orbital dyz orbital dx2-y2 orbital dz2

orbital

s orbital

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 1 2 2

2 s

p

1

3

2

68

3 s

p

d

1

3

5

2

6

10

18

4 S

P

d

f

1

3

5

7

2

6

10

14

32

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

•Read 11.3 Pgs 248-250

•Questions •1 pg 249•14 pg 266

Electron Configurations

I. Writing Electrons Configurations

A. Rules

1. Aufbau Principle

a. An electron occupies the lowest-energy orbital that can receive it

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.

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

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

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 1s2Principle Quantum Number

Sublevel

# of Electrons

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

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

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 exceptionsAnd will steal one electron from an adjacent s sublevelTo fill the d sublevel

½ filled and full sublevels are more stable

•1) Read 11.4 and 11.5 pg pg 251-254

•Worksheet -8 Electron Configurations and Periodicity

•2) Read 11.1 and 11.2 pg245-247

•Questions•10-13 pg 266

•3) Worksheet 11A

Objective Worksheetand

Review Sheet