H CHEM - WED, 9/7/16
Do Now
Be ready for notes.
Sigfig review problem
Homework
Error Analysis
Possibly atomic theory paragraph
Agenda
Atomic Theory
THE ATOM
DEFINITION TO START
Atom – smallest particle of an element that
retains its identity
They are tiny!
Electron microscope – allows us to observe
individual atoms
DEMOCRITUS THE PHILOSOPHER
~490 BC
First credited with proposing the existence of an
atom
Indivisible and indestructible
Shortcomings?
Did Aristotle agree?
ARISTOTLE
~380BC
Did not agree!
Earth, water, air and fire
THEN CAME JOHN DALTON
1766-1844
Experimental methods -> scientific theory
Dalton’s atomic theory
Matter is composed of indivisible atoms
Atoms of same element are identical
Combine in whole number ratios to form compounds
Rxns occur when atoms separate, bond, or rearrange.
Atoms of one element never become atoms of another
thru chemical rxns
JJ THOMSON’S EXPERIMENT -1897
Cathode Ray Experiment
Gas filled glass tube fitted with electrodes
Electricity -> cathode ray (travels from cathode to anode)
Conclusion: electrons - negatively charged subatomic particles
Further tests:
Mass to charge ratio
Different gases
Conclusion: Electrons are part of atoms of all elements.
THOMSON’S MODEL
Thomson – ‘plum pudding’ (1897)
Chocolate chips in cookie dough
MILIKAN - 1909
Oil Drop Experiment
Determined the charge of an electron
OTHER SUBATOMIC PARTICLES?
We know things are neutral…so where’s the positive
Actually discovery is disputed….
Goldstein: anode ray (1886)
Detected rays that contained positively charged particles
Rutherford: 1920
Coined proton after his work with H nuclei
Conclusion: protons – positively charged subatomic
particles
1840x mass of an electron
THOMSON’S MODEL TO RUTHERFORD’S
MODEL
Gold Foil Experiment (1911)
Alpha particles thru gold foil
Predictions: only slight deflection
Results:
Most: straight thru or slight deflection
Some: large deflection or ‘bounced’ back
toward the source
RUTHERFORD MODEL/NUCLEAR MODEL
Atom is mostly empty space
explains why the alpha particles could pass straight thru
All the positive charge and most of mass is located in a small region
explains the large deflections
Nucleus - protons and neutrons
Electrons are around the nucleus and account for most of the volume
Still not quite right!
IT’S ALL ABOUT COLOR…
In terms of atomic models, so far: Dalton (1803) = Tiny, solid particle
Thomson (1897) = “Plum Pudding” model –Electrons stuck on the outside of a big positive charge
Rutherford (1911) = Positively-charged nucleus with electrons moving around it
Rutherford’s model of the atom not quite right Could not explain chemical properties of elements
Could not explain color changes when metal is heated
BOHR MODEL OF THE ATOM-1913
Niels Bohr’s model of the atom
Electron found only on specific, circular paths around nucleus
Each orbit has fixed energy levelHypothesis: When electrons are excited (added
energy), jump into higher energy levels. When they moved back into lower energy levels - gave off light.
Electrons do not exist between levels (think of rungs on a ladder)
Electrons absorb and emit only certain quanta (amounts) of energy
Quantum of energy = fixed amount of energy required to move from one energy level to another energy level
Chapter 5
BOHR’S MODEL
Nucleus
Electron
Orbit
Energy Levels
BOHR’S PLANETARY MODEL OF THE ATOM
Electrons must have enough energy to keep moving around the nucleus
Electrons orbit nucleus in defined energy levels, just like planets orbit the sun
Each energy level assigned a principal quantum number n.
Lowest energy level called ground state (n=1)
Higher energy levels (n=2, 3, 4...) excited states
Model worked OK for hydrogen but not so good for other elements
Nucleus
n = 1n = 2
Ch
ap
ter 5
BOHR’S MODELIn
crea
sin
g e
ner
gy
Nucleus
First
Second
Third
Fourth
Fifth Further away from the
nucleus means more
energy.
There is no “in
between” energy
Energy Levels
Nucleus
Lowest energy level = ground state
Higher energy levels = excited states
Energy
Level 1
Energy
Level 2Energy
Level 3
Electron starts on lowest energy level (ground state)
Add energy to electron –moves to excited state
Energy levels are not evenly spaced
Nucleus
Electron returns to lower state –emits/gives off quantum of energy
Energy
Level 1
Energy
Level 2
Energy
Electron starts on lowest energy level (ground state)
Lowest energy level = ground state
Higher energy levels = excited states
Add energy to electron –moves to excited state
Energy levels are not evenly spaced
Energy
Level 3
Ch
ap
ter 5
Bohr used this
theory to explain
the lines in the
atomic emission
spectra for
hydrogen
Ch
ap
ter 5
410 nm
434 nm
486 nm
656 nm
Each of these lines
corresponds to different
energy changes
So is the Bohr model correct?
For H, it does a wonderful job explaining things
For everything else….not so much
CHADWICK – THE LAST SUBATOMIC
PARTICLE
Chadwick (1932)
Chased evidence for a particle that Rutherford
predicted to exist
Neutrons – neutral subatomic particles
~same mass as proton
History lesson over….for now
WHAT CAN THE PERIODIC TABLE TELL US?
Atomic number – the number of protons in an
element
Elements are defined by their atomic number
Mass number – total number of neutrons and
protons
ISOTOPES!
Same number of protons
Different number of neutrons
Used to calculate the atomic weight of an element
Mass of an atom is tiny – so knowing the actual
mass is a bit impractical
ATOMIC MASS UNIT
Isotope Carbon 12 was assigned the mass of 12
atomic mass units.
So then, 1/12 of the mass of a carbon 12 atom is 1
amu.
So then helium 4 has 1/3 the mass of carbon 12.
Carbon has 6 protons and 6 neutrons…this
accounts for the bulk of the mass…so one proton
or 1 neutron has a mass of ~1 amu.
ATOMIC WEIGHT
Atomic weight – weighted average of the atomic
masses of the isotopes of an element in a
naturally occurring sample
Example problem….Grades!