Download Document

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)
Wavelength
10-11
gamma
10-9
X-ray
10-8
UV
10-6
Visible light
10-5
infrared
10-2
microwaves
1017
1015
1014
frequency
Wavelength increases
Frequency decreases
Energy decreases
1012
1010
(m)
radiowaves
FM
VIBGYOR
1018
10-1
109
short
AM
(s-1)
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
a. Definite frequency
b. Definite wavelength
656nm
Change in Energy (E) = Planck’s Constant(h) x Frequency (ν)
=6.626 x10-34 J s
=3.31 x 10-18J
x 5.00 x 1015 s-1
Frequency (ν ) = Speed of Light (c) / Wavelength(λ)
Therefore
ΔE
=
hv
=
h x c
λ
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
Infrared
Visible
Ultraviolet
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
energy
a. n=1 is closest to the nucleus, and lowest in
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)
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
s orbital
px orbital
dxy orbital
dxz orbital
py orbital
dyz orbital
pz orbital
dx2-y2 orbital
dz2 orbital
Principal
Quantum
Number
(n)
1
Sublevels
in main
energy
level
(n
sublevels)
s
Number of Number
orbitals per of
sublevel
electron
s per
sublevel
Number
of
electrons
per main
energy
level (2n2)
1
2
2
2
s
p
1
3
2
6
8
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
lowestenergy
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
2 # of Electrons
Principle Quantum
Number
Sublevel
1s
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 exceptions
And will steal one electron from an adjacent s sublevel
To 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 Worksheet
and
Review Sheet