Transcript Electron-Config

ELECTRONS IN ATOMS
Electrons in Atoms
 The chemical properties of atoms,
ions, and molecules are related to
the arrangement of the
ELECTRONS within them.
 Over time, the atomic model has
evolved regarding where electrons
are located.
Atomic Theory / Electrons
 Democritus and Dalton
 Atom is indivisible—no electrons
 Thompson’s Model (Plum Pudding)
 Described the atom as a ball of positive
charge containing randomly placed
electrons.
 Rutherford’s Model
 Electrons are still randomly placed, but
now there’s a positive nucleus.
Atomic Theory
The Next Step
 Max Planck’s Quantum Energy Theory
 Energy is always absorbed or emitted in
“packets” which he called quanta.
 Photon= quantum of light
 10th grader = High school student
 Bohr’s Model (Planetary)
 Bohr used the quantum energy theory and
proposed that electrons move in particular
paths with fixed energy in “orbits” around the
nucleus.
Atomic Theory / Electrons
 In Bohr’s Model…
 Electrons cannot exist between energy levels,
but they can move from one level to another.
 Think of a ladder, and the electrons can only jump
from rung to rung.
 To move from one energy level to another one,
an electron must gain or lose just the right
amount of energy.
 This amount of energy is called a quantum of
energy.
Atomic Theory / Electrons
 Bohr’s Model was only partly successful.
 Two other contributions led to a new
atomic theory.
 Louis de Broglie illustrated that
electrons have a dual nature with properties
of both particles and waves.
 Heisenberg’s Principle of Uncertainty
stated that it was impossible to know both
the momentum and the position of an
electron.
Modern Atomic Theory
 Schrodinger’s Quantum Mechanical
Model is the atomic model in use today.
 It was able to explain many previously unexplainable
phenomena.
 Restricts energy of electrons to certain values
 Same as Bohr
 Does NOT define an exact path around the nucleus
 Unlike Bohr
 Estimates the probability of finding an electron in a
certain position (Heisenberg).
Quantum Numbers: Where is the electron??
 In Schrodinger’s model, there are four
“quantum” numbers that tell us where an
electron is likely to be located.
 Principal (n), 1-7, gives the energy level
 Subshell (l), s-p-d-f, gives the shape of region
 Orbital (m), gives the orientation in space of
the shapes
 Spin (s), clockwise or counterclockwise
 No two electrons have the same 4
quantum numbers.
Principal Quantum Number (n)
 Designates the energy level
 Tells how far away from the nucleus an electron
is likely to be.
 Assigned values in order of increasing energy:
n = 1, 2, 3, 4, 5, 6, or 7
 Corresponds to the rows/period on the periodic
table.
Principal Quantum Number (n)
Lower value of n means the electron is
closer to nucleus and has less energy.
+
n=1
n=2
n=3
Subshell (l), s-p-d-f (Shapes)
 Within each principal energy level there are a
certain number of sublevels
 Energy level 1 has 1 sublevel; energy level 2 has 2
sublevels; energy level 3 has 3 sublevels, etc.
 These sublevels are called s, p, d, or f.
 Each sublevel has a unique shape
 These shapes describe the region of space
where the electron is likely to be located
s
p
d
Subshell, (l), s-p-d-f
Principal
Energy
Level
Number
of
Subshells
Type of
Subshell
n=1
1
1s
n=2
2
2s, 2p
n=3
3
3s, 3p, 3d
How many sublevels will principal energy
level 4 have, and what shapes will they be?
Orbitals (m)
(The orbital names s, p, d, and f stand for names
given to groups of lines in the spectra of the alkali
metals. These line groups are called sharp, principal,
diffuse, and fundamental—This is not important.)
•Within each sublevel there are a certain number
of orbitals.
•s sublevel—1 orbital
•p sublevel—3 orbitals
•d sublevel—5 orbitals
•f sublevel—7 orbitals
Each orbital
can only
hold 2
electrons.
•Orbitals within each sublevel have similar
shapes, but are oriented differently in space.
s - orbitals
 Spherical shape
 Hold a maximum
of 2 electrons
 First found in the
1st energy level
 Pear or dumbbell
p-orbitals
shaped
 3 different
orientations
 Each orientation
holds a max of 2 e The p sublevel as a
whole holds a
maximum of 6 e First found in the
2nd energy level
d - orbital
 5 different
orientations
 Each holds 2 e-’s
 The sublevel as a
whole holds a
maximum of 10 e First found in the
3rd energy level
Do not appear until the 4th shell (energy level)
and higher.
Maximum eֿs Allowed
The maximum # of electrons that can
occupy a principal energy level is given by
the formula 2n².
How many eֿs are allowed in energy level 3?
18
Let’s Review
How many orbitals are in the following sublevels?
a. 3p 3
b. 2s 1
c. 4f 7
d. 4p 3
e. 3d 5
Electron Configuration:
•How are the electrons in any given element
arranged?
 In an atom, electrons are positioned around
the nucleus in the most stable arrangement
possible.
 This is the arrangement that has the lowest
energy.
 There are 3 important rules to help us find
the configuration of electrons in atoms.
Aufbau’s Principle
Electrons will occupy orbitals having the
lowest energy FIRST and then in order of
increasing energy.
The "Ground State" of an atom is when
every electron is in its lowest energy
orbital.

Pauli’s Exclusion Principle
An orbital can be empty, have one
electron, or AT MOST have two
electrons.
Electrons occupying the same orbital
must have opposite spins!
One clockwise, and one counter
clockwise
Hund’s Rule
When electrons occupy orbitals
of equal energy, one electron
enters each orbital until all the
orbitals contain one electron with
parallel spins.
Chemistry Riddle
Q: What happens when electrons
lose their energy?
A: They get Bohr’ed!