Download Atomic Theory Lecture Notes

HC History of the atom.notebook
October 21, 2011
History & Development
of Atomic Theory
Aug 28­10:34 AM
History of Atomic Theory
Democritus ~ 460 to 370 BC
Experiment: None. He had beliefs that were disregarded.
Beliefs:
• Named the atom "Atomos" meaning indivisible.
• Matter is composed of atoms too small to be seen.
• Believed atoms could not be created, destroyed or
further divided.
• Atoms are solid and homogeneous with
empty space between them.
• Different atoms have different sizes and shapes
which determine the properties of matter.
Democritus
Atomic Model:
a sphere
Aug 20­1:50 PM
1
HC History of the atom.notebook
October 21, 2011
History of Atomic Theory
The Alchemists ~ 300 BC to 300 AD
Experiment: Pseudoscience influenced by Aristotle,
concerned with...
Using chemistry to change cheap metal to gold.
Finding an eternal life elixir.
Aristotle
Beliefs:
• All matter was made up of a combination of four elements;
fire, wind, earth, water
Aug 20­1:35 PM
History of Atomic Theory
Dalton ~ early 1803 English School Teacher
Developed the first "official" atomic theory.
Experiment: Studied chemical reactions, making observations
and measurements. His experimental results
demonstrated the following...
John Dalton
The Law of Conservation of Matter
total mass of reactants = total mass of products
Ex. 2 H2 + O2 --> 2 H2O
mass of hydrogen + mass of oxygen reacted = mass of water produced
The Law of Definite Proportions
regardless of starting amounts, a compound will always form with the
same mass ratio of elements
Ex. The formula for table sugar will always be C12H22O11, whether you
start with 5.0 grams or 20.0 grams of carbon, hydrogen & oxygen.
The Law of Multiple Proportions
different compounds can form from the same elements but with
different mass ratios
Ex. The formula for water is H2O (2:1 ratio of hydrogen to oxygen)
The formula for peroxide is H2O2 (1:1 ratio of hydrogen to oxygen)
Aug 20­1:58 PM
2
HC History of the atom.notebook
October 21, 2011
Five Principles-Dalton's Atomic Theory
1. All matter is made of indestructible atoms.
2. Atoms of the same element are identical in their physical
and chemical properties.
3. Atoms of different elements have different physical
and chemical properties.
4. Atoms of different elements combine in simple whole
number ratios to form chemical compounds.
5. In chemical reactions, atoms cannot be subdivided,
created, or destroyed. They are combined, separated, or rearranged.
Atomic Model:
Sphere-Billiard Balls
Which principles remain true today?
Aug 23­1:23 PM
History of Atomic Theory
Experiment:
Cathode Ray Tube
Electricity is passed through a glass tube filled
with gas. The gas beam can be bent with a
magnet.
Joseph John Thomson
J.J. Thomson ~ 1897 English Physicist
Discoveries:
• Atoms consist of charged particles
• The negatively charged particles are called electrons (1897)
• The positively charged particles are called protons (1920)
Model:
Plum Pudding or Chocolate Chip Cookies
The chocolate chips
represent negatively charged
electrons stuck in positively
charged dough.
Aug 20­3:07 PM
3
HC History of the atom.notebook
October 21, 2011
History of Atomic Theory
Robert Millikan ~ 1909
Experiment:
Oil Drop
Discoveries:
measured the charge on an electron
determined the mass of an electron
Aug 23­1:56 PM
History of Atomic Theory
Rutherford ~ 1911 English Physicist
(a student of Thomson)
Experiment: Gold Foil Experiment
Positively charged alpha particles are shot
at a piece of thin gold foil. Most alpha
particles had little deflection. Some were
deflected at large angles.
Ernest Rutherford
Discoveries:
• A positively charged core of an atom called the nucleus
• Electrons surround the nucleus
• The rest of the atom is empty space
Model: Nuclear atom
Quote: "It was about as believable as if you had fired a 15 inch shell at a
piece of tissue paper, and it came back and hit you."- Rutherford
Aug 20­3:19 PM
4
HC History of the atom.notebook
October 21, 2011
History of Atomic Theory
James Chadwick ~ 1932 English Physicist
(Student of Rutherford)
Discoveries:
• The nucleus contains another particle with no charge
(neutral) and a mass equal to that of a proton.
• Called this particle the neutron.
Model:
Sir James Chadwick
Same as Rutherford
Aug 20­3:30 PM
Rutherford's Model Revisted
The Nuclear Atom:
• Atoms have a positively
charged nucleus containing
protons (+) and neutrons (0).
Electrons (-) are located outside
of the nucleus.
The Problem:
According to the laws of physics, with nothing holding the electrons
in place, the negatively charged particles would spiral into the
positively charged nucleus (- and + charges attract) and Rutherford's
atom would collapse.
Sep 16­7:33 AM
5
HC History of the atom.notebook
October 21, 2011
Niels Bohr ~ 1913 Danish Physicist
• student of both Thomson & Rutherford
The Experiment: He studied the light spectrum of
hydrogen gas in an attempt to more
accurately explain the role of electrons in
the atom.
Sep 16­7:45 AM
Electromagnetic Radiation:
Radiation that is produced by electrically charged particles (like electrons!)
All electromagnetic radiation exhibits wave like behavior.
Wavelength - the distance
from one peak to another
Amplitude - the height of
a peak
Frequency - the number of
waves that pass a particular
point per second
Sep 10­8:14 PM
6
HC History of the atom.notebook
October 21, 2011
The Electromagnetic Spectrum
Low Energy Light = longer wavelength & lower frequency
red light = low energy light
High Energy Light = shorter wavelength & higher frequency
blue light = high energy light
Sep 10­8:21 PM
Light Spectrums
Hydrogen Spectrum
Helium Spectrum
• Bohr observed when energy (light or heat) was added to a
sample of hydrogen gas a distinctive color was given off.
• When he looked at the light through a prism, only certain
color lines could be seen, instead of the whole rainbow like he
expected.
• Concluded that the reason for why hydrogen's spectrum only
contained a red, a blue, and a purple line involved the behavior
of the electrons in the hydrogen atoms.
Sep 16­7:59 AM
7
HC History of the atom.notebook
October 21, 2011
Bohr's Experimental Conclusions
• Thought the atom contained energy levels which kept the
electrons from falling into the nucleus.
• Individual elements each have a unique arrangement of electrons
in those energy levels.
• When an atom absorbs energy the electrons enter into an "excited
state" by jumping up to higher energy levels.
• The excited state is unstable, so the electrons release the excess
energy in the form of colored light.
• The color of the light observed is determined by the amount of
energy released by the electrons.
• Electrons in lower positions (closer to the nucleus) have lower
energy (red light)
• Electrons in higher positions (farther from the nucleus) have
higher energy (blue light)
Sep 16­7:27 AM
Important Vocabulary:
ground state- all electrons in their lowest possible
energy levels
excited state- electrons absorbed energy & jump
to a higher energy level
Sep 16­9:41 AM
8
HC History of the atom.notebook
October 21, 2011
Bohr's Model: Solar System
Electrons "orbit" around the nucleus
in certain paths (energy levels).
An electron must have a certain
amount of energy to be "allowed" in a
particular energy level.
Similar to the way the planets in
our solar system orbit around the
sun.
The energy levels prevent the
electrons from falling into the nucleus.
Sep 16­7:59 AM
Excited vs. Ground State Cont...
The Problem with Bohr's Model???
It worked well for hydrogen, but not for
elements containing more than 1 electron.
Sep 16­9:50 AM
9
HC History of the atom.notebook
October 21, 2011
History of Atomic Theory
Quantum Mechanics
Erwin Schrodinger ~ 1926 Austrian Physicist
Experiment: NONE! It is a mathematical model. It cannot be
represented by anything that exists in the real world.
Mathematical Equation:
Problem:
Erwin Schrodinger
Heisenberg's Uncertainty Principle: you can never know how fast
an electron is moving and where an electron is at the same time.
In other words, you can find out where the electron started and
you can see where the electron ended up but how it got there
WE DON'T KNOW!
Discoveries:
• Mathematical model that deals with the probability of finding an electron within a
given space around the nucleus
• The probability is 90%
• The given space are called orbitals (or electron clouds)
• There are four orbitals with different shapes: s p d f
• These orbitals can be related to the periodic table
• Electrons behave like waves.
Model:
Aug 20­4:21 PM
The Quantum Mechanical Model
Quantum Numbers:
Numerical values that mathematically describe the energy and
location of electrons in atoms.
From these values you get orbitals - probability maps that show
the possible areas around the nucleus where electrons may be
located.
Four Quantum Numbers:
1. Principal Quantum Number (n)- Refers to the energy level (the distance of the
orbital from the nucleus) where an electron is
located.
• When n=1 the electron is closest to the nucleus and has the least energy.
• n = 1,2,3,4,5, etc.
2. Angular Quantum Number (l)- Refers to the shape of the orbital; s, p, d or f
• l = 0, 1, 2 ... (n-1)
• l = 0 (s-orbital), l = 1 (p-orbital), l = 2 (d-orbital)...
3. Magnetic Quantum Number (m)- Orientation of orbital(s)
• m = any integer from -l to +l
4. Spin Quantum Number (s)- electron spin
• electrons can have either "up spin" (+) or "down spin" (-)
• s = +1/2 or -1/2
Aug 24­7:18 PM
10
HC History of the atom.notebook
October 21, 2011
Orbital Shapes
p - orbitals are dumbbell shaped
s - orbitals are spherical
Aug 25­8:00 PM
Orbital Shapes
d-orbitals are clover shaped
f-orbitals
Sep 7­6:48 PM
11
HC History of the atom.notebook
October 21, 2011
Electron Configurations:
Represent the arrangement of electrons in an atom.
The Rules:
1. Aufbau Principle: An electron will occupy the lowest possible energy level.
2. Pauli Exculsion Principle: Each orbital cannot contain more than two
electrons, and those electrons must have opposite spin.
3. Hund's Rule: Orbitals of equal energy are each occupied by one electron
before any orbital is occupied by a second electron, and all single electrons must
have the same spin
Sep 10­7:54 PM
Sep 10­8:32 PM
12
HC History of the atom.notebook
October 21, 2011
Using the Periodic Table to Write Configurations.
• Each row on the periodic table represents an energy level.
• The s, p, d and f-orbitals are represented by different
"blocks" on the periodic table.
• The d-orbitals are located 1 energy level lower than
the s & p orbitals in the same row.
• The electrons configuration for a particular element is
determined by the location of that element on the table.
Aug 23­4:09 PM
The Orbital Blocks
Note: When writing configurations, the "d" orbitals
are 1 energy level lower than the "s" & the "p"
around them.
Aug 23­4:15 PM
13
HC History of the atom.notebook
October 21, 2011
Examples...
Hydrogen:
coefficient = energy level
(row on periodic table)
H
1s1
letter = orbital type
(periodic table block)
superscript = number of
electrons in that orbital
(# of spaces over on table)
He
C
Mg
Fe
Br
Ag
Configurations of Ions: Add or remove the number of
electrons gained or lost to the end of the configuration.
Examples:
Mg:
Mg+2:
N:
N-3:
Aug 23­4:16 PM
Final Few Notes on Configurations:
• Most ions have electron configurations identical to a noble gas.
• Noble gases are stable. Therefore when ions form they lose or gain
electrons in order to have stable electron configurations as well.
• Valence electrons - electrons in the highest energy level of an atom.
Example: How many valence electrons does phosphorus have?
• f-orbitals... 1 energy level lower than the "d"
• Example: Au
Aug 23­4:42 PM
14
Attachments
Pictures of atom
Chadwick apparatus
Time line of scientists
Schrodinger equation
Conversation with science
orbitals