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Atomic Theory
Chapter 2 Pg: 41-57
2.1.1 State the position of protons, electrons and neutrons in the atom
2.1.2 State the relative masses and relative charges of protons, neutrons and
electrons
2.1.3 Define the terms mass number (A), atomic number (Z) and isotopes of an
element
2.1.4 Deduce the symbol for an isotope given its mass number and atomic number
2.1.5 Calculate the number of protons, neutrons and electrons in atoms and ions
from the mass number, atomic number and charge.
2.1.6 Compare the properties of the isotopes of an element
2.1.7 Discuss the uses of radioisotopes
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
atoms
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 compound always
has the same relative number
and kind of atoms
Atoms are divisible!
• By the 1850s,
scientists began to
realize that the atom
was made up of
subatomic particles
• Thought to be positive
and negative
• How would we know
this if we can’t see it or
touch it?
Cathode Rays and Electrons
• Mid-1800’s scientists began to study electrical
discharge through cathode-ray tubes. Ex: neon
signs
• Partially evacuated tube in which a current passes
through
• Forms a beam of electrons which move from cathode to
anode
• 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
• http://www.chem.uiuc.edu/clcwebsite/video
/Cath.mov
• 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
(nucleus of helium) were
scattered as they passed
through a thin gold foil
• http://www.mhhe.com/
physsci/chemistry/ess
entialchemistry/flash/r
uther14.swf
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
enough
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
Name
Symbol
Charge Relative Mass Actual Mass (g)
(amu)
Electron
e-
-1
1/1840
9.11x10-28
Proton
p+
+1
1
1.67x10-24
Neutron
no
0
1
1.67x10-24
• 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
nucleus
• 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
A
elemental symbol
atomic number
Z
E
• Find the
• number of protons
• number of neutrons
• number of electrons
• atomic number
• mass number
19
9
F
80
35
Br
184
74
W
Ions
• Cation is a positively charged particle.
Electrons have been removed from the
element to form the + charge.
ex: Na has 11 e-, Na+ has 10 e-
• Anion is a negatively charged particle.
Electrons have been added to the atom to
form the – charge.
ex: F has 9 e-, F- has 10 e-
Isotopes
• 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
element
1
1
H
Hydrogen-1
2
1
H
Hydrogen-2
3
1
H
Hydrogen-3
• 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
isotopes)
• 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) + …
Problem:
Silver has two naturally occurring isotopes, 107Ag with a
mass of 106.90509 amu and abundance of 51.84 %
,and 109Ag with a mass of 108.90476 amu and
abundance of 48.16 % What is the average atomic
mass?
Average = (0.5184)(106.90509) + (0.4816)(108.90476)
= 107.87 amu
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
value
• Remember, the atomic masses are the decimal
numbers on the periodic table
Properties of Isotopes
• Chemical properties are primarily determined by
the number of electrons
• All isotopes has the same number of electrons,
so they have nearly identical chemical
properties even though they have different
masses.
• Physical properties often depend on the
mass of the particle, so among isotopes they will
have slightly different physical properties such
as density, rate of diffusion, boiling point…
More Practice 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?
Radioisotopes
• Isotopes of atoms that have had an extra
neutron attached to their nucleus.
• Carbon-14 radioactive decay is used to
measures the date of objects.
– After 5700 years the amount of 14C will be half
its original value.
• Iodine-125 or 131 is used to monitor the
activity of the thyroid gland (b/c the thyroid
tends to absorb iodine)
• Cobalt-60 produces gamma rays (intense
radioactivity) and is used in radiation
treatment of cancer.
• Note: gamma rays are the shortest
wavelength on the electromagnetic
spectrum. They are the most dangerous
and difficult to shield from.