Download CHAPTER 5 NOTES – ELECTRONS IN ATOMS

CHAPTER 11 NOTES
MODERN ATOMIC THEORY
RUTHERFORD’S MODEL COULD
NOT EXPLAIN THE CHEMICAL
PROPERTIES OF ELEMENTS
The Bohr Model
• Bohr proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus
• Energy Levels – the fixed energies an
electron can have – like rungs of a ladder
• Quantum – the amount of energy required
to move an electron from one energy level
to another energy level
• Quantum Mechanical Model – the modern
description of the electron in atoms – from
the mathematical solutions to the
Schrödinger equation – determines the
allowed energies an electron can have
and how likely it is to find the electron in
various locations around the nucleus
ATOMIC ORBITALS
• Atomic orbital – a region in space in which there
is a high probability of finding an electron
• Different atomic orbitals are denoted by letters.
 s Orbitals
 p Orbitals
 d Orbitals
 f Orbitals
 g Orbitals
• Each energy sublevel corresponds to an orbital
of different shape describing where the electron
is likely to be found
Hydrogen Energy Levels
• The s and p types of sublevel
Principal Quantum Numbers
• (n) – always
equals the
number of
sublevels
within that
principal
energy level
Principal
Energy
Level
# of
Sublevel
s
Type of
Sublevel
s
Max # of
Electrons
n=1
1
1s
2
n=2
2
2s,2p
8
n=3
3
3s,3p,3d
18
n=4
4
4s,4p,4d, 32
4f
n=5
5
5s,5p,5d, 50
5f,5g
n=6
5
6s,6p,6d, 50
6f,6g
n=7
2
7s,7p
8
ELECTRON CONFIGURATIONS
•
•
Electron Configurations – the ways in which electrons are arranged into
various orbitals around the nuclei of atoms
3 RULES
AUFBAU PRINCIPLE
•
ELECTRONS OCCUPY THE ORBITALS OF LOWEST ENERGY FIRST
PAULI EXCLUSION PRINCIPLE
•
AN ATOMIC ORBITAL MAY DESCRIBE AT MOST 2 ELECTRONS
HUND’S RULE
•
ELECTRONS OCCUPY ORBITALS OF THE SAME ENERGY IN A WAY THAT
MAKES THE NUMBER OF ELECTRONS WITH THE SAME SPIN DIRECTION AS
LARGE AS POSSIBLE
EXAMPLE I
o Element = SODIUM
o Element Symbol = Na
o ATOMIC NUMBER = 11
o NUMBER OF ELECTRONS = 11
o LONG-HAND VERSION
1s2 2s2
2p6
3s1
EXAMPLE I continued
• SHORT-HAND VERSION
1s22s22p63s1
• NOBLE GAS CONFIGURATION
[Ne]3s1
EXAMPLE II
o Ion Name = SODIUM ION
o Ion Symbol = Na+
o ATOMIC NUMBER = 11
o NUMBER OF ELECTRONS = 10
o LONG-HAND VERSION
1s2 2s2
2p6
EXAMPLE II continued
• SHORT-HAND VERSION
1s22s22p6
• NOBLE GAS CONFIGURATION
[Ne]
EXCEPTIONAL ELECTRON
CONFIGURATIONS
• MEMORIZE THE FOLLOWING
Cr, Cu, Mo, Pd, Ag and Au
• Some actual electron configurations differ
from those assigned using the Aufbau
Principle because half-filled sublevels are
not as stable as filled sub-levels, but they
are more stable than other configurations
• Transition metals usually lose s orbital
electrons first.
EXAMPLE I
•
•
•
•
Element Name = Chromium
# of Electrons = 24
Short-Hand Version = 1s22s22p63s23p64s13d5
Noble Gas Configuration = [Ar] 4s13d5
EXAMPLE II
•
•
•
•
Element Name = Copper
# of Electrons = 29
Short-Hand Version = 1s22s22p63s23p64s13d10
Noble Gas Configuration = [Ar] 4s13d10
• Transition metal ions having partially filled d orbitals
usually have a color.
• Transition metals usually lose s orbital electrons first.
• Example A: Fe and Fe3+
• Fe (26 electrons) - 1s22s22p63s23p64s23d6
• Fe3+ (23 electrons) - 1s22s22p63s23p63d5
• Example B: Cu and Cu2+
• Cu (29 electrons) - 1s22s22p63s23p64s13d10
• Cu2+ (27 electrons) - 1s22s22p63s23p63d9
Physics and the Quantum
Mechanical Model
• Amplitude – wave’s height from zero to the crest
• Wavelength (λ) – the distance between the crests
• Frequency (ν) – the number of wave cycles to pass a
given point per unit of time
SI unit of frequency is a hertz (Hz) or
expressed as a reciprocal second (s-1 or 1/s)
• c = λν
• The wavelength and frequency of light are
inversely proportional to each other
• c = speed of light (3E8 m/s or 3 E10 cm/s)
• When atoms absorb energy, electrons move into
higher energy levels, and these electrons lose
energy by emitting light when they return to
lower energy levels.
•
•
•
•
E
E
h
ν
=
=
=
=
hν
energy measure in Joules (J)
Planck’s Constant = 6.6262E-34 Js
frequency (s-1)
•
•
•
•
•
E
E
h
c
λ
=
=
=
=
=
hc/λ
energy measure in Joules (J)
Planck’s Constant = 6.6262E-34 Js
speed of light = 3E8 m/s
wavelength measured in meters (m)
•
•
•
•
•
λ = h/mv
λ = wavelength measured in meters (m)
h = Planck’s Constant = 6.6262E-34 Js
m = mass measured in kilograms (kg)
v = velocity measured in meters per
second (m/s)
EXAMPLES
• WHAT IS THE FREQUENCY OF RADIATION
WHOSE WAVELENGTH IS 550 NM?
• WHAT IS THE ENERGY (IN J) OF A PHOTON
WHOSE FREQUENCY IS 3.2 E 14 HZ?
• WHAT IS THE WAVELENGTH (IN NM) OF
RADIATION WITH A FREQUENCY OF 6.50E14 S-1?
Electromagnetic Spectrum
• Atomic emission spectrum – frequencies of light emitted
by an element that separate into discrete lines
• Ground State – lowest possible energy of an electron (n
= 1). Excitation of the electron by absorbing energy
raises it from the ground state to an excited state (n =
2,3,4,5,6 or 7)
• The light emitted by an electron moving from a higher to
a lower energy level has a frequency directly
proportional to the energy change of the electron
• Heisenberg uncertainty principle – it is impossible to
know exactly both the velocity and the position of a
particle at the same time.
Atoms can give off light.
 They first must receive energy and become
excited.
 The energy is released in the form of a
photon.
 The energy of the photon corresponds exactly
to the energy change experienced by the
emitting atom.
• Atomic states
 Excited state – atom with excess energy
 Ground state – atom in the lowest
possible state
• When an H atom absorbs energy from
an outside source it enters an excited
state.
Quantized Energy Levels
• Since only certain energy changes occur
the H atom must contain discrete energy
levels.
Quantized Energy Levels
• The energy levels of all atoms are
quantized.