Transcript Electrons in Atoms Part 2 – Quantum Mechanical - chem30-wmci

ELECTRONS IN ATOMS PART
2 – QUANTUM
MECHANICAL MODEL
QUANTUM MECHANICAL
MODEL
• The Bohr Model of the atom has several short comings
• Cannot work beyond element 21 very well
• Doesn’t describe certain behaviour of atoms or
molecules
• The model currently used to describe the atom is the
Quantum Mechanical Model of the atom
• This is the current theoretical framework that is used to
describe all of the information we have about atoms
and how they function
BASIC DEFINITIONS
• Quantum (plural ‘quanta’)
• A finite amount of energy
• i.e. – an energy level in an atom
• The amount of energy required to move an electron
from its present energy level to the next higher one
• Mechanical
• Movement of parts in relation to a whole
• i.e. – electrons in an atom
• Hence the Quantum Mechanical Model deals
with the movement and location of electrons
in an atom
UNCERTAINTY PRINCIPLE
• We cannot know where an electron is and where it is
going at the same time
• The more precisely the position of an electron is known,
the less precisely its momentum is known and vice versa
• Because of this, we use probability to determine where
an electron is most likely to be
ELECTRON CLOUDS
• Using the electron
probabilities, we find areas
where electrons are most
likely to be
• These areas are called
electron clouds where the
probabilities of finding
electrons is very high
• The shapes and distance
from the nucleus of these
electron clouds depends
on several factors
QUANTUM NUMBERS
Principal Quantum
Number
•
•
•
Energy level
Distance from the nucleus
Represents the PERIOD on the periodic table
Angular Quantum
Number
•
The shape of the orbital; represented the
letters s, p, d, and f.
Magnetic Quantum
Number
•
Determines the orientation of the orbital in
space
•
Spin Quantum
Number
•
•
Which axis it lies on
Specifies the value for spin
Electrons in the same orbital must spin in
opposite directions
QUANTUM NUMBERS
• Principal Quantum Number
• Energy level
• Distance away from the nucleus
• As # increases, distance from the nucleus also increases
• As the number increases, so does the energy of the
electrons in those orbitals
• Represented by integers 1,2,3,4,5,6,7 that correspond to
the seven horizontal rows on the periodic table
• Determined by counting as you move down (top to
bottom) the periodic table
QUANTUM NUMBERS
• Angular Quantum Number
• Also known as “sub-shells”
• Refer to the shape of the orbital
• There are four (4) different shapes
• S, P, D, F
• These correspond to the s, p, d, f blocks on the periodic
table
QUANTUM NUMBER
“S” Sub-shell
 Spherical shape
Only one (1) orbital per
energy level
 The 1 sub shell can
hold 2 electrons
One with +1/2 spin
One with -1/2 spin
QUANTUM NUMBERS
“P” Sub-shell
 Dumbbell shape
 Three (3) orbitals per
energy level
mL and be from –L to +L
 Each shell can hold 2
electrons
3 orbitals mean the pshell can hold up to 6
electrons
QUANTUM NUMBERS
• “D” Sub-shell
• Tend to have a
clover-leaf shape
• Five (5) orbitals per
energy level
• Each can hold a
maximum of two (2)
electrons
• Can hold a max of
10 electrons
QUANTUM NUMBERS
• “F” Sub-shell
• Shape contains 6
lobes for the most
part
• Seven (7) orbitals per
energy level
• Each can hold a
maximum of two (2)
electrons
• Fourteen (14)
electrons total at
each energy level
TO SUMMARIZE
Sub-shell
Number of
sub-shells
at a level
S
Energy
level (n) in
which it is
first found
1
1
Number of
electrons in
these subshells
2
P
2
3
6
D
3
5
10
F
4
7
14
QUANTUM NUMBERS
• Spin Quantum
Number
• Remember, in each
sub-shell there can
be two (2) electrons
• These electrons must
have spins that go in
opposite directions
• Represented by
arrows pointing in
opposite directions
REPRESENTING ELECTRONS
USING THE QUANTUM
MECHANICAL MODEL
• There are two (2) different types of notation used to
represent the quantum mechanical model:
• Orbital Notation
• Electron Configuration Notation
ORBITAL NOTATION
• Illustrates the following quantum numbers: principal,
second (shape), and spin
• Use the template to draw and “fill” the sub-shells with
electrons
• Order of filling electrons is governed by three (3) rules:
• Aufbau Principle
• Pauli Exclusion Principle
• Hund’s Rule
ORBITAL NOTATION
Aufbau Principle:
 Electrons enter sub-shells of lowest energy first
 1st energy level fills up before the next
Pauli Exclusion Principle:
 All atomic sub-shells contain a maximum of two (2)
electrons. Each MUST have a different spin
Hund’s Rule:
 when electrons occupy sub-shells of equal energy,
ONE electron enters EACH sub-shell until all the subshells contain one electron with identical directions
 Electrons are added to sub-shells so that a maximum
number of unpaired electrons result
EXAMPLES
• Oxygen
• Titanium
• Strontium
ELECTRON
CONFIGURATION
• Illustrates the following quantum numbers: principal,
and angular (shape)
• Does not indicate the spin, but it does indicate the
number of electrons
• Easier to use because there is less to draw
ELECTRON CONFIGURATION
EXAMPLES:
• Nitrogen
• Chlorine
• Copper
• Tin
NOBLE GAS NOTATION
• An even more simplified and shorthand
method for representing electron
configuration.
• Emphasizes the outermost energy level only
• Instead of listing every energy level and
amount of electrons individually, it utilizes
the nearest noble gas element of the
energy level below as a representation of
the inner energy levels
NOBLE GAS NOTATION
• For Example: Sulfur
• Electron configuration would be:
• 1s22s22p63s23p4
• Its Noble Gas Notation would be:
• [Ne] 3s23p4
• …this is because we know that the electron configuration
of Ne is: 1s22s22p6, therefore there is no need to write it all
out.