Transcript Electrons in Atoms - Mrs. Haug's Website

Chapter 5





From Democritus to Rutherford, models of the
atom have changed due to new experiments.
As technology develops, a more complete model
of the atom is developed.
Rutherford’s model identified the nucleus
surrounded by electrons.
His model DID NOT explain why some things
glow when heated.
His model DID NOT explain the chemical
properties of elements.



Niels Bohr (1885-1962) was one of
Rutherford’s students. He added to the
model.
Proposed that an electron is found only in
specific circular paths, or orbits, around the
nucleus.
Studied how energy of an atom changes
when it absorbs or emits light.

Energy levels are fixed “paths” or energies an
electron can have when orbiting the nucleus.





Close to nucleus = less energy
Further from nucleus = more energy
e- must reside on an energy level
Moving from one level to another is possible if the
right energy is lost or gained.
A quantum of energy is the amount of
energy required to move an electron from
one level to another.


Energy gained or lost in an atom is
not always the same.
Energy levels are not evenly
spaced.
 Higher levels are closer together.
Therefore less energy needed to
move levels further from nucleus.



Erwin Schrodinger (1887-1961) devised and
solved a mathematical equation describing
the behavior of the electron in a hydrogen
atom.
This model comes from his mathematical
solutions.
Determines the allowed energies an electron
can have and how likely it is to find the
electron in various locations around the
nucleus.



Based on his equations, Schrodinger was also
able to explain atomic orbitals.
Orbitals explain the probability of finding an
electron at various locations around the
nucleus.
Often thought of as the region of space in
which there is a high probability of finding an
electron.

Used to describe the region of space with the
highest probability of finding an electron.
 Energy level (n)
▪ n = 1,2,3,4,5,6,7
 Sublevels correspond to an orbital of a different shape
▪ Each energy level has an equal number of sublevels
▪ Denoted by letters (s,p,d,f)
 Orbital
▪ Describes highest probability
▪ Contained within sublevels
▪ Shapes describe probability
Principal
Energy Level
Number of
sublevels
Type of
Sublevel
Orbitals
n=1
1
1s
1
n=2
2
2s
2p
1
3
n=3
3
3s
3p
3d
1
3
5
n=4
4
4s
4p
4d
4f
1
3
5
7





Each energy sublevel corresponds to an
orbital of a different shape, which describes
where the electron is likely to be found.
These “shapes” are based on mathematical
probability experiments.
Click Here
“s” orbital is spherical
“p” orbital is dumbbell shaped
“d” orbital is clover shaped or dumbbell in a
donut.

Three Rules for Electron Configuration
 Aufbau Principle – e- occupy the orbitals of lowest
energy first
▪ Orbitals of any sublevel are always equal energy
 Pauli Exclusion Principle – atomic orbital may
describe, at most, two e- of opposite spin
 Hund’s Rule – One electron enters each orbital
until each orbital has an electron. Then orbitals
get partners.

Use the websites linked to this page to help
determine electron configuration.
https://www.caymanchem.com/app/template/chem
Assistant,Tool.vm/itemid/4001;jsessionid=EEAEFB4
09423347FDE326280AABDD091
 http://www.chem1.com/acad/webtut/atomic/Orbita
lPT.html




Now that you know the rules… you may
MEMORIZE and use this cheat sheet. I’ll
show you how to use it.
http://www.mpcfaculty.net/mark_bishop/co
mplete_electron_configuration_help.htm
Now practice!




Like everything else, there are exceptions to
the rules.
Some atoms are more stable when their
outer shells “break the rules.”
Half-filled sublevels are not as stable as filled
sublevels, but they are more stable than
other configurations.
Examples are copper and chromium