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Chapter 2:
Atoms Molecules and Ions
The investigation and
understanding of the atom is
what chemistry is all about!
Anyone who says that they can contemplate quantum
mechanics without becoming dizzy has not understood
the concept in the least.
-Niels Bohr
©Bires, 2002
Slide 1
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Atomic History
• The Greek philosopher Democritus (400BC)
coined the term atomon which means “that
which cannot be divided.”
• Idea of an indivisible thing that made up all
matter was refined by colorblind chemist John
Dalton in 1803. Among his interests, Dalton was
very interested in a scientific explanation for his
colorblindness and the behavior of gases.
• In his A New System of Chemical Philosophy,
Dalton published five principles of matter.
©Bires, 2002
Slide 2
Dalton’s Top Five
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• All matter is made of indestructible and
indivisible atoms. (atoms are hard,
unbreakable, and the smallest thing there is)
• Atoms of a given element have identical
physical and chemical properties. (all atoms of
X will behave the same anywhere) These two are
usually combined
• Different atoms have different properties. (X
behaves differently than Y)
• Atoms combine in whole-number ratios to form
compounds. (two H’s and one O = Water (H2O)
• Atoms cannot be divided, created or destroyed,
(just rearranged) in chemical reactions.
©Bires, 2002
Slide 3
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The Laws:
• Constant Composition: Ratios of atoms in a
compound is constant for that compound.
• Conservation of Mass: Mass is not created or
destroyed in a chemical reaction.
• Multiple Proportions: Since atoms bond in
small, whole number ratios to form compounds,
their masses are small whole number ratios.
©Bires, 2002
Slide 4
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The CRT (cathode ray tube)
Electrostatics!
• A new invention, the cathode ray tube, (c1850s)
suggested the presence of charges – areas of
positive and negative…charge.
• This suggested that atoms must be divisible,
and Dalton’s theory had to be modified.
• In 1897, English Physicist J. J. Thompson
proposed that the atom is a sphere of positive
charge with small areas of negative charge.
• This theory became known as the “plum
pudding” model after an English “dessert” of
purple bread and raisins.
©Bires, 2002
Slide 5
Millikan’s Oil
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• In 1897, Thompson used electrostatics
experiments to determine the charge-to-mass
1 C = 1 Amp•1 second
ratio of 1.76 x 108 C/g.
• Millikan’s classic oil-drop experiment allowed
the charge of a single electron to be
determined: 1.60 x 10-19 C.
• Using these two numbers, we can calculate the
1.60 x1019 C
mass of an electron: m

 9.10 x1028 g
electron
•
•
©Bires, 2002
1.76 x108 C / g
The mass of an electron is about
1/2000 of the mass of a proton!
Slide 6
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Ernest Rutherford
• While studying radioactive elements, New
Zealander Physicist Ernest Rutherford found
that radioactive alpha particles deflected when
fired at a very thin gold foil.
• This was known as the gold foil experiment,
and it suggested that the atom was not a hard
sphere as thought, but was mostly space, with
a small concentration of mass.
• This concentration of mass became known as
the nucleus.
• Link to experiment…
©Bires, 2002
Slide 7
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The Bohr Model
• A Danish physicist, Niels Bohr, (a student of
Rutherford) rebuilt the model of the atom placing
the electrons in energy levels (the “solar system”
model).
• Bohr was one of the founders of quantum physics – a
discipline that states that energy can be given off in
small packets or quanta of specific size.
• Energy levels closer to the nucleus were lower in
energy than those farther away.
• When a specific amount of energy was added to an
atom, an electron could jump into a higher energy
level.
©Bires, 2002
Slide 8
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Adding the Neutrons
• Although theorized by Rutherford, British
physicist, James Chadwick, proved the
existence of massive, neutral particles.
• These particles came to be called neutrons…
• and their discovery in 1932 opened the door for
more in depth investigations into radioactive
materials and gave the WWII Allies the ability to
enrich and purify fissionable uranium, a
necessity in the production of nuclear weapons.
©Bires, 2002
Slide 9
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The Modern Model
Democritus and
Dalton’s atom
electron
Thompson’s electrons
neutron
Rutherford’s space and
nucleus
proton
Bohr’s energy levels
(not to scale)
©Bires, 2002
Chadwick’s neutrons
Most atoms ~ 1-5 Å
(Angstroms) = (1x10-10m)
Slide 10
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Elements
8
O
• We currently know of about 110
elements, 92 of which are
naturally occurring.
OXYGEN
• We illustrate an element with an
15.9994
atomic symbol.
• We add the number of protons,
the atomic number…
• And the atomic mass, the total
mass of the protons plus the
neutrons. Notice that the atomic mass is not a round number,
©Bires, 2002
even though protons and neutrons each have a mass
of 1. This is due to natural abundance.
Slide 11
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Natural Abundance - Isotopes
• There is not just one type of each atom, there
are several. When a nuclide has more or less
neutrons than another nuclide of the same
element, we call them isotopes.
• For instance, the element carbon has 6 protons,
but it could have 5, 6, 7, or 8 neutrons, to form
12C
Carbon-11, Carbon-12, Carbon-13, and
13C
Carbon-14. Each has a different mass.
14C
• In nature, there is a mix of different natural
isotopes. We use this mix to calculate average
atomic mass…
©Bires, 2002
Slide 12
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Calculating Average Atomic Mass
• To find average atomic mass, we multiply the
relative abundance of an isotope by the mass of
the isotope. We then add each of the products
for each isotope.
• Example: The isotopes of element Bob are
found below:
• Bob-18, 25%
0.25x18  0.60x19  0.15x20
• Bob-19, 60%
• Bob-20, 15%
• What is the average atomic mass of naturally
occurring Bob?
1amu = 1.66x10-24 grams
18.90amu
©Bires, 2002
Slide 13
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Isotopes, Ions, and Allotropes (Oh my)
• Isotopes are atoms of the same element with
different numbers of neutrons.
• Ions are atoms of the same element with
different numbers of electrons.
• Ions are easy to create; adding or removing
electrons can be done with electric current.
• Allotropes are forms of the same element,
bonded in different structures.
• Diamond and pencil graphite are allotropes.
They are both pure carbon, but in different
structures.
©Bires, 2002
Slide 14
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Isotopes and allotropes
©Bires, 2002
Slide 15
Balancing Chemical Equations
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General Rules:
Compounds and Diatomics
Never change subscripts (H2SO4)
Balance groups first (H2SO4)
Balance water, oxygen, and hydrogen last
Diatomic Atoms
Hydrogen = H2
Fluorine = F2
Bromine = Br2
Nitrogen = N2
Chlorine = Cl2
Iodine = I2
• Recall that compounds are collections of atoms.
• Compounds can be made by:
Ionic Bonding: Giving and taking electrons
(Metal-nonmetal bond)
Covalent Bonding: (Molecules) Sharing
electrons (nonmetal-nonmetal bond)
• There are gases that are bound to themselves
as elements:
• Hydrogen = H2 Nitrogen = N2 Oxygen = O2
• Fluorine = F2
Chlorine = Cl2 Bromine = Br2
• Iodine = I
Oxygen = O2
©Bires, 2002
Slide 16
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Formulas
• We can represent molecules with molecular
and empirical formulas.
• Molecular Formula = shows actual number of
atoms in a molecule.
C6 H12O6
H 2O2
• Empirical Formula = shows ratio of atoms in a
molecule.
CH2O
HO
• Compounds are represented with empirical
formulas.
©Bires, 2002
Slide 17
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Ions
• Recall that an ion is an atom that has
gained or lost one or more electrons.
• Recall that the octet rule predicts that Cl
atoms try to achieve zero or eight
electrons in their outer (valence) shell.
Cl-
Na
Na+
• When an atom bonds with another
atom, it seeks to gain electrons or lose
them. For instance:
Positive ions are
• Cl has 7 and will gain one electron
called “cations”
• Na has 1 and will lose one electron
Negative ions are
called “anions”
©Bires, 2002
Slide 18
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Ionic attraction
• Just like opposite poles of a
magnet attract, oppositely
charged ions attract.
• Here, our positively charged
sodium ion and negatively
charged chlorine ion are
attracted.
• They remain ions, but stick
together in a lattice (3D grid
pattern) as other ions join
them.
©Bires, 2002
Cl
Na+
Cl- Na+
Na+
Cl-
+
Na
Cl
Slide 19
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Formula Units
• In our example, the Na+ and Clcombined to form NaCl.
• Here, NaCl is one formula unit – Cl- Na+
the smallest unit that has the
correct formula for our compound.
• NaCl is also our compound’s
empirical formula – the smallest
ratio of atoms that make up our
compound.
• Although there may be several
Na+ ions and Cl- ions in this
lattice, the empirical unit is still
just NaCl.
©Bires, 2002
Slide 20
•
•
•
•
Inorganic
Nomenclature
General Rules:
Cations are listed first
Anions/polyatomics last
Binary compounds get the “–ide”
ending
ex: (sodium chloride)
• Acids:
Monatomic anion : “hydro-” + name
+ ic “acid” (ex: hydrochloric acid)
Polyatomic anion “-ite” = name + “ous acid”
ex: (sulfite = sulfurous acid)
• Polyatomic anion “-ate” = name +
“-ic acid”
(ex: sulfate = sulfuric acid)
• Compounds:
Use Greek prefixes to denote
number of atoms of each
©Bires, 2002
Back
Naming Compounds Flow Chart
Does the Formula begin with H
NO
YES
Does it begin with a metal which has more than
one oxidation number? (e.g.; Fe, Ni, Cu, Sn, Hg)
NO
YES
It is an Acid (must be aqueous)
Does the acid contain a polyatomic ion?
Name the first element
followed by its
oxidation number
(Roman numeral)
YES
NO
Ending of polyatomic ion
Does the formula contain a polyatomic ion?
NO
-ite
YES
Are both elements nonmetals?
NO
-ate
Name the polyatomic
ion, replacing the –ate
ending with –ic. Add
the word acid.
YES
Name the first element,
Then the second element
With the –ide ending.
Name the first element
using the proper prefix.
(never mono-). Name the
second element with the
proper prefix (including
mono- and –ide ending.
1 mono2 di3 tr4 tetra5 penta6 hexa7 hepta8 octa9 nano10 decaName the first element, then the
polyatomic ion. If two elements
are present, name both, then the
polyatomic ion. (e.g.; NaHCO3
Is sodium hydrogen carbonate)
Name the polyatomic
ion, replacing the –ite
ending with –ous. Add
the word acid.
Write the prefix hydro,
then the name of the
second element with –ic
ending. Add the word
acid.
PRACTICE!
PRACTICE!
PRACTICE!
PRACTICE!
Slide 21
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Molecular and Empirical Formulas
• The formulas for compounds that exist
as molecules are called molecular
formulas.
Ex: Hydrogen peroxide, H2O2, consists
of two H atoms and two O atoms.
• The atomic ratio of hydrogen peroxide
is one to one.
• The simplest formula that would
indicate this ratio would be HO. This
simple formula is an empirical formula.
©Bires, 2002
Slide 22
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Molecular and Empirical Formulas
• Formulas for ionic compounds are
almost always empirical.
Ex: NaBr is empirical because it takes
one
sodium atom to combine with one
bromine atom to obtain a neutral
charge.
©Bires, 2002
Slide 23
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Molecular and Empirical Formulas
• C6H12O6
• CH2O
is molecular formula for glucose
is empirical because we could
factor out a six from each
subscript.
• CH3COOH is molecular for acetic acid
• CH2O
is empirical (factor out a 2).
* Note: glucose and acetic acid are very
different molecules yet they have the same
empirical formula.
©Bires, 2002
Slide 24
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Analysis by Combustion
• Formulas for organic substances can be
determined by combustion analysis. Carbon is
recovered as CO2, hydrogen as water. Oxygen
can be determined by mass difference if only C,
H and O are present. 2B discussed more
later…
©Bires, 2002
Slide 25
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Steps for Determining an Empirical
Formula:
1. Start with the number of grams of each element,
given in the problem.
*If percentages are given, assume that the total mass is
100 grams so that
the mass of each element = the percent given.
2. Convert the mass of each element to moles using
the molar mass from the periodic table.
3. Divide each mole value by the smallest number of
moles calculated.
4. Round to the nearest whole number. This is the
mole ratio of the elements and is represented by
subscripts in the empirical formula.
©Bires, 2002
Slide 26
What if you cannot round off to whole
numbers?
Back
– If the number is too far to round (#.10 to
#.90), then multiply each solution by the
same factor to get the lowest whole number
multiple.
• e.g. If one solution is 1.5, then multiply
each solution in the problem by 2 to get 3.
• e.g. If one solution is 1.25, then multiply
each solution in the problem by 4 to get 5.
©Bires, 2002
Slide 27
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Molecular Formulas
• Calculations that are based on % composition
always lead to the empirical formula of a
compound (because it is already in the most
reduced ratio).
• Once the empirical formula is found, the
molecular formula for a compound can be
determined if the molar mass of the compound
is known.
• The molecular formula differs from the empirical
formula in that it is a whole number multiple of
the empirical formula. The empirical formula is
the most reduced ratio of the molecular formula.
©Bires, 2002
Slide 28
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Calculating Molecular Formulas
STEPS:
1. Calculate the mass of the empirical
formula and divide the molar mass of the
compound by the mass of the empirical
formula to find the ratio between the
molecular formula and the empirical
formula.
2. Multiply all the atoms (subscripts) by this
ratio to find the molecular formula.
©Bires, 2002
Slide 29
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Formula for a hydrate
• Finding the empirical formula for a hydrate
involves finding the mass difference of the
hydrate and anhydrous salt by driving off the
excess water:
• DEMO:
http://www.chem.iastate.edu/group/Greenbowe/
sections/projectfolder/flashfiles/stoichiometry/e
mpirical.html
©Bires, 2002
Slide 30
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THE END
©Bires, 2002
Slide 31
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The Periodic Table
• The periodic table is a collection of all the
known elements into a model that groups
elements with similar properties.
• Vertical columns represent Groups
of elements with similar properties.
• Horizontal Periods represent elements with
similar electron configurations.
©Bires, 2002
Slide 32
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The Periodic Table
• The periodic table was created by Russian
chemist Dmitri Mendeleev, who ordered the
known elements according to atomic mass.
• The modern periodic table is ordered according
to atomic number.
• The elements in the periodic table are arranged
with Periodic Law which states that the
chemical and physical properties of elements
are periodic functions of their atomic numbers.
• Periodic Law shows certain trends in the
properties of elements (more in C6 and C7)
©Bires, 2002
Slide 33