Download Atomic Structure

Atomic Structure
Part 1:
The History of the Atomic Model
Democritus
• About 400 BC:
•
Democritus thought the atom existed,
but had no proof
•
The word atom comes from “atomos”
which means indivisible
•
He said that matter must be made of
these indivisible particles
Aristotle
•
•
Aristotle taught that matter was
continuous – that it was not made of
particles, such as Democritus
suggested.
Aristotle was a Greek philosopher who’s
ideas were the standard at that time.
He was highly respected and thought to
be a very wise source of information.
John Dalton
•
•
1800’s
John Dalton has proof that atoms exist and
writes the first “Atomic Theory”
1. All elements are made up of atoms
2. Atoms of the same element are identical.
Different elements have different kinds of
atoms.
3. Atoms can form mixtures or chemically
combine in whole number ratios.
4. Atoms cannot be destroyed or changed into
other elements.
Dalton’s model
• He assumed the atoms were small, hard
spheres with different sizes and different
masses for different elements.
J.J. Thomson
• 1897:
• J.J Thomson discovered electrons by
using a cathode ray tube.
• By using magnets, he bent the cathode
rays and determined that the charge of the
particles in the ray are negative.
Animation - Cathode Ray
Video
Thomson’s model
• Theory at the time said that the whole
atom was neutral (had no charge), so if
the electrons had a negative charge, there
also must be a positively charged part of
the atom, thus the idea of the proton.
• Thomson suggested a model that had a
positive base, with negative parts
scattered throughout – “plum pudding”
model.
Ernest Rutherford
• 1911
• Rutherford and his associates discover
the nucleus.
• Gold Foil Experiment – positively charged
alpha particles are deflected by a dense
positively charged nucleus in gold atoms.
animation Gold Foil Experiment
Gold Foil 2
Rutherford’s model
• The nucleus must be very tiny.
• The nucleus has a positive charge.
• The nucleus contains most of the atom’s
mass
• The electrons are very lightweight and are
found in the area surrounding the nucleus.
• The atom is mostly empty space.
One question remained:
Why don’t the electrons
collapse into the nucleus?
Niels Bohr
• 1913
• Suggested that the electrons are held in
“fixed” positions around the nucleus.
• Electrons have a fixed amount of energy –
the energy is “quantized.”
• Planetary model – electrons orbit the
nucleus.
• In Search of a Giant w/ Brian Cox
Robert Millikan
Watch the Experiment
• 1916
• Oil Drop Experiment - Using charge ratios,
Millikan found the mass of the electron to
be 1/1837 of the mass of the atom.
• He also stated that the electron had one
unit of negative charge. Therefore
suggesting the idea of a second particle (a
proton) with one unit of positive charge.
Eugen Goldstien
• canal rays – opposite cathode rays
• Goldstein's work with anode rays of H+
was apparently the first observation of the
proton
James Chadwick
• 1932
• Confirmed the existence of a third particlethe neutron.
• The neutron has no charge.
• The neutron has the same mass as a
proton.
Irwin Schrödinger
• 1926
• Quantum Mechanical Model of the AtomSchrödinger used mathematical equations
to formulate the probability of finding
electrons located at fixed energy levels
around the nucleus.
• Electrons are not contained to orbits, but
are inside an “electron cloud.”
http://en.wikipedia.org/wiki/Schr%C3%B6di
nger_equation
The Modern Quantum Model
• An electron cloud shape is called an
“orbital.”
• The orbital shapes are probable locations
of the electrons based on the amount of
energy of the electron.
http://www.orbitals.com/orb/orbtable.htm
Atomic Structure
Part 2:
The Parts of the Atom
Major Subatomic Particles
• Proton – positively charged +
found inside the nucleus
• Electron – negatively charged –
found outside of the nucleus
• Neutron – neutral charge 0
found inside the nucleus
Atomic Mass
• The total mass of the atom is found inside
the nucleus. This includes the protons and
the neutrons.
Atomic Mass = Number of Protons + the number of neutrons
Atomic Number
• The number of protons in the nucleus of an
atom.
• Also equal to the number of electrons in a
neutral atom.
Atomic number = number of protons = number of electrons
Isotopes
• Atoms of the same element which have a
different number of nuetrons, and
therefore a different atomic mass.
• Example: Hydrogen has three isotopes.
•
http://education-portal.com/academy/lesson/isotopes-and-calculation-of-aweighted-average.html
Representation
• What do these mean? (Examples of isotopes)
Average Atomic Mass
• Masses recorded on the periodic table of
elements are the “weighted average” of
the various isotope masses that exist for
that element.
• Therefore the element’s mass is generally
not a whole number.
Calculating average atomic mass
Quantum Mechanical model
http://www.youtube.com/watch?v=45KGS1
Ro-sc
• http://www.youtube.com/watch?v=B7pACq_xWyw
http://www.youtube.com/watch?v=ARuyAX
mjjVI&list=PL69CA6EEFAEF15D1D
Chapter 5
Quantum Mechanical Model of
the Atom
• This model suggests that electrons are
located at definite energy levels inside an
electron cloud
• According to this theory, the electrons do
not travel in paths around the nucleus, but
can be found anywhere within a defined
region of space around the nucleus. We
call this area the electron cloud.
Electrons are “quantized”
• The lowest level is called ground state. Here, the
electrons have enough energy to maintain their
position and not collapse into the nucleus.
• Electrons can move from one level to another by
gaining or losing energy or “quanta” (tiny
bundles of energy)
• The more energy the electron has, the farther
away from the nucleus it will be found.
Schrödinger's Probability
• The exact location of an electron is
unknown.
• Mathematically, we can predict the most
probable areas where the electrons can be
found.
https://images=_&imgdii=_&images
Electron Density Diagram
The real picture of an orbital/
The famous equation
Orbitals
• The most probable areas where the
electrons can be found are called orbitals.
s orbital
p orbital
d orbital
Quantum Numbers
• Quantum numbers are assigned to
describe the location and energy of
electrons.
• The lowest energy level of the electrons is
called ground state.
Quantum Numbers
• Principal quantum numbers – these
designate the major energy levels, and
thus the distance from the nucleus.
• n = 1 - 1st and lowest energy level
(closest to the nucleus)
n=3
n=2
n = 2 - next closest
n=1
n = 3, n = 4, n = 5, etc….
• Orbital quantum numbers - describe the
sublevels of energy and show the shape of the
orbitals.
•
s – spherical (lowest sublevel)
•
p – dumbbell shaped (three kinds)
py
px
pz
• d – clover shaped ( five kinds)
dxy
dxz
dx2 – y2
dxy
dz
• f - complicated patterns (seven kinds)
http://www.uky.edu/~holler/html/orbitals_2.html
http://www.youtube.com/watch?v=A6DiVsp
oZ1E
Electron Configurations
• Electron Configuration – the way that
electrons are arranged in their most stable
pattern around the nucleus.
• Three rules apply:
Aufbau principle
Pauli Exclusion Principle
Hund’s Rule
Aufbau Principle
• Aufbau principle – electrons enter the
orbitals of lowest energy first.
• Follow this pattern: p 133
1s
2s
2p
3s
3p
3d
4s
4p
5s
4d
5p
4f
5d
5f
Pauli Exclusion Principle
•
Pauli exclusion principle – each orbital
can hold a maximum of two electrons.
These electrons must have opposite
spins. Electrons are represented by up
and down arrows.)
1s
Hund’s Rule
• Hund’s Rule – When electrons enter
orbitals of equal energy, one electron
enters each orbital until all the orbitals
contain one electron with spins in parallel.
Then the second electrons are entered
with opposite spins.
px
py
pz
Writing Orbital Notations
Another Example
• Cr has 24 electrons
Writing Electron Configurations
Shortcut Method