Atoms and Elements
Visualizing Atoms
 Binnig & Rohrer –
development of the
Scanning Tunneling
Microscope (STM).
 Produce images on the
atomic level.
 Iodine atoms on the
surface of platinum
metal
Modern Atomic Theory
 Law of Conservation of Mass
 Antoine Lavoisier – 1789
 In a chemical reaction, matter is neither created nor
destroyed.
7.7 g Na
+ 11.9 g Cl2
 19.6 g NaCl
Modern Atomic Theory
 Law of Definite Proportions
 Proust – 1797
 All samples of a given compound will have the same
proportions of their constituent elements.
Mass of Water
Mass of Oxygen
Mass of Hydrogen
Ratio O:H
18.0 g
16.0 g
2.0 g
8:1
45.0 g
40.0g
5.0 g
8:1
Modern Atomic Theory
 The atomic theory of matter
 Dalton, 1808
 Four postulates (main themes)
1.
Each element is composed of tiny, indestructable
particles called atoms.
2. All atoms of a given element are identical; The atoms
of different elements are different and have different
properties.
Modern Atomic Theory
3. Atoms combine in simple, whole number ratios
to form compounds.
4. Atoms of one element cannot change into atoms
of another element. In a chemical reaction,
atoms only change the way they are bound
together.

The evidence for the existence of atoms is
overwhelming!
Discovery of Atomic Structure
 Discovery of the electron.
 J.J. Thompson - cathode ray tube - 1897 – determined
that an electron was negatively charged.
Discovery of Atomic Structure
 Robert Milliken - oil drop - 1909 – determined the
charge value of an electron as well as its mass.
Oil Drop Experiment
Discovery of Atomic Structure
 Radioactivity
 Curie – 1900
 Rutherford – 1905
 Three main types of radioactive particles: Alpha, Beta,
and Gamma.
 Alpha particles
 Essentially a helium nuclei.
 Sources – many, but Curie’s used radium.
Nuclear Model of the Atom
 Ernest Rutherford - 1911.
 Gold foil experiment.
 Gold can be smashed into very thin sheets that are only
a few atoms thick.
 Alpha particles, from an alpha source, beamed at the
gold foil.
 Photographic paper placed as a detector in front and
behind alpha source.
Gold Foil Experiment
Gold Foil Experiment
Nuclear Model of the Atom
 Only 1 in 8000 alpha particles is scattered.
 Scattering occurs when an alpha particle
encounters a massive gold nuclei.
 Rutherford proposed that:
1. Most of the atom’s mass and all of its positive
charge were found in the small core of the atom
called the nucleus.
2. Most of the volume of the atom is empty space.
Nuclear Model of the Atom
 Rutherford’s model still had one problem.
 H = 1 proton in nucleus.
 He = 2 protons in nucleus.
 He mass  4x mass of H mass.
 Final piece of the puzzle is the neutron.
 Neutrons have no charge and a mass of 1amu.
 Discovered in 1932 by James Chadwick.
Modern View of Atomic Structure
 Three subatomic
particles exist in
an atom.
 Protons,
neutrons, and
electrons.
Modern View of Atomic Structure
 Electrons have a negative charge of 1.602 x 10-19 C and a
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negligible mass.
Protons and neutrons reside in the nucleus which is
extremely small.
Protons have a positive charge, equal in magnitude to
an electron and a mass of about 1amu .
Neutrons have no charge and a mass of 1amu.
Over 99.9% of the mass of an atom resides in the
nucleus.
Modern View of Atomic Structure
 The atom is 100,000 times larger than the nucleus. If a
golf ball represented the size of the nucleus, the atom
would be about 3 miles in diameter.
 Diameter of nucleus  10-15 m.
 Diameter of atom  10-10 m.
 The density of the nucleus is roughly 1013 to 1014 g/cm3.
Atomic and Mass Numbers
 The Atomic Number (Z) is equal to the number of
protons in the nucleus.
 Each element has a unique atomic number and hence
a unique number of protons.
 The Mass Number (A) is the sum of the protons and
neutrons found in the nucleus.
Isotopes & Symbols
 Isotopes for an element occur when they have more
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than one mass number.
Isotopes of an element have the same number of
protons, but a different number of neutrons.
An isotope can be designated by its mass number in
the upper left corner – AX.
Or it can be designated after the symbol – X–A.
Example - 12C or C-12
Ions
 Atoms quite often will gain or lose electrons when
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

forming compounds.
When Lithium metal reacts, it loses one electron
forming a +1 charge. Li+1.
When Fluorine gas reacts, it gains one electron
forming a -1 charge. F-.
Positive ions are called cations.
Negative ions are called anions.
Periodic Law
 Mendeleev (1869) – first to group the elements by similar
properties.
 First, he listed them in order of increasing atomic mass.
 He then started a new row when elements had similar
properties.
 Thus, they are arranged according to horizontal rows
which highlight the repeating properties of the
elements in the vertical columns.
 periods - the elements in a horizontal row constitutes a
period.
 groups - the elements in a vertical column constitutes a
group.
The Periodic Table
 Groups are numbered and labeled with A and B's.
 Different conventions of numbering are used,
however, we will use the traditional N.A. method with
the A’s for the first two and last six groups.
The Periodic Table
 The A groups are called main group elements and the
B groups are called transition elements.
 Some groups have special names:
 Group 1A = Alkali metals
 Group 2A = Alkaline Earth metals
 Group 7A = Halogens
 Group 8A = Noble gases or inert gases.
The Periodic Table
The Periodic Table
 Metals - a substance that has a characteristic of luster
or shine, and is a good conductor of heat and
electricity.
 Metals are solids at room temperature and tend to lose
electrons easily.
 Metals are found to the left and below the diagonal
line that runs through the right side of the main group
elements.
 The greatest majority of elements are metals.
The Periodic Table
 Non-metals - either a gas at room temperature or a brittle
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solid, they are non-conductors of heat and electricity.
Non-metals tend to gain electrons easily.
Non-metals are found to the right and above the diagonal
line.
Metalloids - are elements that have the properties of both
metals and non-metals.
These fall along the diagonal line and include B, Si, Ge, As,
Sb, and Te
Hydrogen – the one oddball on the periodic table.
Describing An Element
 How would you describe the element:

C

Al

Sr

Fe
Ions
 Metals form cations.
Ex) Na  Na+ + 1e Non-metals form anions.
Ex) Cl + 1e-  Cl Predicting the charge of a species based on the
periodic table.
Ions and Ionic Compounds
 Main group elements usually form only one charge
(valence)
 Transition metals usually form many different charges.
 Polyatomic ions are special groups of atoms that
chemically combine to form a charged species.
Atomic Weights
 The Atomic Mass Unit Scale is based on the 12C atom.
 1 amu = 1/12 the mass of the 12C atom.
 Average atomic masses are based on the masses of
each type of isotope a well as their abundance in
nature.
 Relationship of an amu to grams:
 1 amu = 1.66054 x 10-24 g
 1 g = 6.012214 x 1023 amu
Determining An Average Mass
 A.W. = S(fract. abundance) x (isotopic mass)
 Example
 Chlorine has two isotopes, Cl-35 and Cl-37. Cl-35 has a
mass of 34.969amu and an abundance of 75.78% Cl-37
has a mass of 36.966amu and an abundance of 24.22%.
What is the atomic weight of Chlorine?
Determining An Average Mass
A.W. = (34.969amu) x (0.7578)
+ (36.966amu) x (0.2422)
A.W. = 35.45amu
Molar Mass
 The Mole
 Unit of quantity used in Chemistry.
 Not convenient to count atoms.
 1 atom of C-12 = 12 amu (exact)
 One mole of C-12 = 12 grams (exact)
 Avogadro’s Number
 Represents the number of C-12 atoms in 12 grams.
 =6.02 x 1023 atoms.