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Introduction to Atomic Theory
History of the atom
•Democritus (400 BC) suggested that the material world was
made up of tiny, indivisible particles
• atomos, Greek for “uncuttable”
• Aristotle believed that all matter was made up of 4
elements, combined in different proportions
• Fire - Hot
• Earth - Cool, heavy
• Water - Wet
• Air - Light
• The “atomic” view of matter faded for centuries, until early
scientists attempted to explain the properties of gases
Re-emergence of Atomic Theory
John Dalton postulated that:
1. All matter is composed
of extremely small,
indivisible particles called
2. All atoms of a given
element are identical
(same properties); the
atoms of different
elements are different
3. Atoms are neither
created nor
destroyed in
chemical reactions,
only rearranged
4. Compounds are
formed when atoms
of more than one
element combine
• A given
always has the
same relative
number and kind
of atoms
Make-up of the Atom
• By the 1850s,
scientists began to
realize that the atom
was made up of
subatomic particles
• Thought to be positive
and negative
Cathode Rays and Electrons
• Mid-1800’s scientists began to study electrical
discharge through cathode-ray tubes. Ex: neon
• Partially evacuated tube in which a current passes
• Forms a beam of electrons which move from cathode to
• Electrons themselves can’t be seen, but certain
materials fluoresce (give off light) when energised
Oh there you are!
• JJ Thompson observed that when a
magnetic or electric field are placed
near the electron beam, they influence
the direction of flow
• opposite charges attract each other,
and like charges repel.
• The beam is negatively charged so it
was repelled by the negative end of
the magnet
• Magnetic field forces the beam to
bend depending on orientation
• Thompson concluded that:
• Cathode rays consist of beams of particles
• The particles have a negative charge
• Thompson understood that all matter was
inherently neutral, so there must be a counter
• A positively charged particle, but where to put it
• It was suggested that the negative charges were
balanced by a positive umbrella-charge
• “Plum pudding model” “chocolate chip cookie model”
Rutherford and the Nucleus
• This theory was
replaced with another,
more modern one
• Ernest Rutherford
(1910) studied angles
at which a particles
were scattered as they
passed through a thin
gold foil
Rutherford expected …
• Rutherford believed that the mass and positive
charge was evenly distributed throughout the atom,
allowing the a particles to pass through unhindered
a particles
Rutherford explained …
• Atom is mostly empty space
• Small, dense, and positive at the center
• Alpha particles were deflected if they got close
a particles
The modern atom is composed of two regions:
• Nucleus: Containing
protons and neutrons, it is
the bulk of the atom and
has a positive charge
associated with it
Electron cloud:
Responsible for the majority
of the volume of the atom, it
is here that the electrons
can be found orbiting the
nucleus (extranuclear)
Major Subatomic Particles
Charge Relative Mass Actual Mass (g)
• Atoms are measured in picometers, 10-12 meters
• Hydrogen atom, 32 pm radius
• Nucleus tiny compared to atom
• If the atom were a stadium, the nucleus would be a marble
• Radius of the nucleus is on the order of 10-15 m
• Density within the atom is near 1014 g/cm3
Elemental Classification
• Atomic Number (Z) = number of protons (p+) in the
• Determines the type of atom
• Li atoms always have 3 protons in the nucleus, Hg always 80
• Mass Number (A) = number of protons + neutrons
[Sum of p+ and nº]
• Electrons have a negligible contribution to overall mass
• In a neutral atom there is the same number of
electrons (e-) and protons (atomic number)
Nuclear Symbols
• Every element is given a corresponding symbol
which is composed of 1 or 2 letters (first letter upper
case, second lower), as well as the mass number
and atomic number
mass number
elemental symbol
atomic number
• Find the
• number of protons
• number of neutrons
• number of electrons
• atomic number
• mass number
• If an element has an atomic number of 34 and a
mass number of 78 what is the:
• number of protons in the atom?
• number of neutrons in the atom?
• number of electrons in the atom?
• complete symbol of the atom?
• If an element has 91 protons and 140 neutrons
what is the:
• atomic number?
• mass number?
• number of electrons?
• complete symbol?
• Atoms of the same element can have different
numbers of neutrons and therefore have different
mass numbers
• The atoms of the same element that differ in the
number of neutrons are called isotopes of that
• When naming, write the mass number after the name of
the element
How heavy is an atom of oxygen?
•There are different kinds of oxygen atoms (different
• 16O, 17O, 18O
• We are more concerned with average atomic
masses, rather than exact ones
• Based on abundance of each isotope found in nature
• We can’t use grams as the unit of measure
because the numbers would be too small
• Instead we use Atomic Mass Units (amu)
• Standard amu is 1/12 the mass of a carbon-12 atom
• Each isotope has its own atomic mass
Calculating Averages
Average = (% as decimal) x (mass1) +
(% as decimal) x (mass2) +
(% as decimal) x (mass3) + …
Try Again…THINK!!
Your test and quiz mark is 30%. 10% is
made up of quiz, 20% is made up of tests
If you made 85%, 78% and 79% on your
If you made 87% on your only test…
What would be your test quiz mark for your
report card?
Calculating Averages
• Calculate the atomic mass of copper if copper has
two isotopes
• 69.1% has a mass of 62.93 amu
• The rest (30.9%) has a mass of 64.93 amu
• Magnesium has three isotopes
78.99% magnesium 24 with a mass of 23.9850 amu
10.00% magnesium 25 with a mass of 24.9858 amu
The rest magnesium 26 with a mass of 25.9826 amu
What is the atomic mass of magnesium?
Average Atomic Masses
• If not told otherwise, the mass of the isotope is
the mass number in amu
• The average atomic masses are not whole
numbers because they are an average mass
• Remember, the atomic masses are the decimal
numbers on the periodic table