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Chemistry Section III A World of Particles
 A “theory” indicates that an explanation is supported by overwhelming evidence.
 A model is a simplified representation of something you want to explain (so a model that represents the
structure of an atom is called an atomic model).
 An Atom- the smallest unit of an element that retains the chemical properties of that element and can exist as
a separate particle.
 Atomic theory: All matter is made up of atoms, and helps us make accurate predictions about the behavior of
matter.
 Models of the atom, and the scientists that led us there
o Democritus (solid sphere model)
 His question: Democritus wondered how many times it would be
possible to break a piece of matter in half, if the pieces would just
keep getting smaller forever. He thought that if he could just keep
breaking matter in half he would eventually end up with the smallest
bit of matter possible. This led to what we know as the atom.
 His hypothesis: atoms are eternally unchanging and indivisible (he
was not able to prove his thoughts due to lack of technology)
o John Dalton
 Father of Atomic Theory; First to show proof of atoms
 Experiment: He observed elements combine in whole number ratios to form compounds;
Matter is NOT created or destroyed in chemical reactions.
 4 Postulates of Theory:
1) all matter is made of atoms. Atoms are indivisible and indestructible.
2) All atoms of a given element are identical in mass and properties
3) Compounds are formed by a combination of two or more different elements
4) A chemical reaction is a rearrangement of atoms
o J.J. Thomson (plum-pudding model)
 Experiment: Zapped atoms with electricity
 Conclusion: atoms consist of negatively charged electrons found
inside positively-charged spheres; this is called the “plum pudding”
model
o Rutherford (discovered the nucleus)
 His experiment: gold-foil experiment using alpha particles
 Conclusion: atoms contain a small, massive, and positively-charged particle called a
nucleus.
o Bohr (solar system model)
 Experiment: Observed light is given off when elements
are exposed to flame or electric fields
 Conclusion: an atom consists of a dense, positivelycharged nucleus, containing nearly all the atom’s mass,
surrounded by electrons traveling in specific allowed
orbits, like the planets around a star—sometimes called
the “planetary” model.
o Heisenberg (cloud model, named the Uncertainty Principle)
 Experiment: “Thought experiment” He imagined a microscope that could an electron and to
measure its position. He found that the electron's position and momentum did indeed obey
the uncertainty relation he had derived mathematically
 Conclusions: Electrons are located in clouds, not neat orbits; tells you where the electron is
most likely to be found (a matter of probability).
o Chadwick
 Discovered the neutron.
 Experiment: He followed up on
the work performed by Ernest
Rutherford. Chadwick bombarded
alpha rays at beryllium. When
struck, the beryllium emitted mysterious neutral rays.

Conclusion: He reasoned that neutrons were important in holding the positively charged
protons together
o
Review
o
 Atomic number, mass, chemical symbol
How can we determine the number of Electrons?
 Valence electrons: electrons in the outer most energy level; determines reactivity of element
Example: Use the previous notes to fill in the blanks and answer the questions.
# of protons = _____
# of electrons = _____
# of neutrons = _____
Identity of atom = ______
Are the # of protons &
electrons equal? ______

Isotope: atoms of the same elements with different number of
neutrons
 Isotopes have different masses
o Mass Number: # of p+ + # of n O
 specific to the isotope; may be different from mass on
periodic table
o Isotope notation
Example: Write the isotope notation for the three isotopes of Helium shown below.
# of protons = _____
# of electrons = _____
# of neutrons = _____
Identity of atom = ______
Are the # of protons &
electrons equal? ______


Average atomic mass
o The average atomic mass is a weighted average mass of all the naturally occurring isotopes of an
element based on the abundance of the element in nature.
Calculating average atomic mass
Example: Hydrogen is 99% 1H, 0.8% 2H, and 0.2% 3H.

Calculating Average Atomic Mass by converting % to decimal
first
Example: calculate the average atomic mass of magnesium.