Spencer L. Seager
Michael R. Slabaugh
www.cengage.com/chemistry/seager
Chapter 2:
Atoms and Molecules
SYMBOLS & FORMULAS
• A unique symbol is used to represent each element.
• Formulas are used to represent compounds.
• ELEMENTAL SYMBOLS
• A symbol is assigned to each element. The symbol is based
on the name of the element and consists of one capital letter
or a capital letter followed by a lower case letter.
• Some symbols are based on the Latin or German name of
the element.
CHEMICAL ELEMENTS & THEIR SYMBOLS
CHEMICAL ELEMENTS & THEIR SYMBOLS
(continued)
Elements from Group 7A
chlorine
bromine
iodine
COMPOUND FORMULAS
• A compound formula consists of the symbols of the elements
found in the compound. Each elemental symbol represents
one atom of the element. If more than one atom is
represented, a subscript following the elemental symbol is
used.
COMPOUND FORMULAS EXAMPLES
• Carbon monoxide, CO
• one atom of C
• one atom of O
• Water, H2O
• two atoms of H
• one atom of O
• Ammonia, NH3
• one atom of N
• 3 atoms of H
Example 2.4 & 2.7
ATOMIC STRUCTURE
• Atoms are made up of three subatomic particles, protons,
neutrons, and electrons.
• The protons and neutrons are tightly bound together to form the
central portion of an atom called the nucleus.
• The electrons are located outside of the nucleus and thought to
move very rapidly throughout a relatively large volume of space
surrounding the small but very heavy nucleus.
SUBATOMIC PARTICLES
• Protons are located in the nucleus
of an atom. They carry a +1
electrical charge and have a mass
of 1 atomic mass unit (u).
• Neutrons are located in the
nucleus of an atom. They carry no
electrical charge and have a mass
of 1 atomic mass unit (u).
• Electrons are located outside the
nucleus of an atom. They carry a
-1 electrical charge and have a
mass of 1/1836 atomic mass unit
(u). They move rapidly around the
heavy nucleus.
SUBATOMIC PARTICLE CHARACTERISTICS
ATOMIC STRUCTURE REVIEW
• Which subatomic particles are
represented by the pink spheres?
• Which subatomic particles are
represented by the yellow and blue
spheres?
• What structure do the yellow and
blue spheres form?
Example 2.9
ATOMIC & MASS NUMBERS
• ATOMIC NUMBER OF AN ATOM
• The atomic number of an atom is equal to the number of
protons in the nucleus of the atom.
• Atomic numbers are represented by the symbol Z.
• MASS NUMBER OF AN ATOM
• The mass number of an atom is equal to the sum of the
number of protons & neutrons in the nucleus of the atom.
• Mass numbers are represented by the symbol A.
ATOMIC & MASS NUMBERS APPLICATION
• Based on the information given above, what is the atomic
number of fluorine?
ATOMIC & MASS NUMBERS APPLICATION
• Based on the information given above, what is the atomic
number of fluorine?
• Answer: The atomic number of fluorine is 9.
• On the periodic table, the atomic number is written as a
whole number above the symbol F.
• In the written description, fluorine is said to have 9 protons
(the atomic number is the number of protons).
• In the symbol, the number 9 is written in the atomic number
or Z (lower left) position.
ATOMIC & MASS NUMBERS APPLICATION
• Based on the information given above, what is the mass
number of fluorine?
ATOMIC & MASS NUMBERS APPLICATION
• Based on the information given above, what is the mass
number of fluorine?
• Answer: The mass number of fluorine is 19.
• In the written description, fluorine is said to have 9 protons
and 10 neutrons (the mass number is the sum of the
numbers of protons and neutrons).
• In the symbol, the number 19 in written in the mass number
or A (upper left) position.
• Note: The periodic table does not show the mass number
for an individual atom. It lists an average mass number for
a collection of atoms!
ISOTOPES
• Isotopes are atoms that have the same number of protons in
the nucleus but different numbers of neutrons. That is, they
have the same atomic number but different mass numbers.
• Because they have the same number of protons in the
nucleus, all isotopes of the same element have the same
number of electrons outside the nucleus.
ISOTOPE SYMBOLS
, where Z is the
• Isotopes are represented by the symbol
atomic number, A is the mass number, and E is the
elemental symbol.
60
• An example of an isotope symbol is 28 Ni. This symbol
represents an isotope of nickel that contains 28 protons and
32 neutrons in the nucleus.
• Isotopes are also represented by the notation: Name-A, where
Name is the name of the element and A is the mass number of
the isotope.
• An example of this isotope notation is magnesium-26. This
represents an isotope of magnesium that has a mass number
of 26.
Example 2.15 & 2.21
RELATIVE MASSES
• The extremely small size of atoms and molecules makes it
inconvenient to use their actual masses for measurements or
calculations. Relative masses are used instead.
• Relative masses are comparisons of actual masses to each
other. For example, if an object had twice the mass of another
object, their relative masses would be 2 to 1.
ATOMIC MASS UNIT (u)
• An atomic mass unit is a unit used to express the relative
masses of atoms. One atomic mass unit is equal to 1/12 the
mass of a carbon-12 atom.
• A carbon-12 atom has a relative mass of 12 u.
• An atom with a mass equal to 1/12 the mass of a carbon-12
atom would have a relative mass of 1 u.
• An atom with a mass equal to twice the mass of a carbon-12
atom would have a relative mass of 24 u.
ATOMIC WEIGHT
• The atomic weight of an element is the relative mass of an
average atom of the element expressed in atomic mass units.
• Atomic weights are the numbers given at the bottom of the
box containing the symbol of each element in the periodic
table.
• According to the periodic table, the atomic weight of nitrogen
atoms (N) is 14.0 u, and that of silicon atoms (Si) is 28.1 u.
This means that silicon atoms
are very close to twice as
massive as nitrogen atoms.
Put another way, it means that
two nitrogen atoms have a total
mass very close to the mass of
a single silicon atom.
MOLECULAR WEIGHT
• The relative mass of a molecule in atomic mass units is called
the molecular weight of the molecule.
• Because molecules are made up of atoms, the molecular
weight of a molecule is obtained by adding together the
atomic weights of all the atoms in the molecule.
• The formula for a molecule of water is
H2O. This means one molecule of water
contains two atoms of hydrogen, H, and
one atom of oxygen, O. The molecular
weight of water is then the sum of two
atomic weights of H and one atomic
weight of O:
• MW = 2(at. wt. H) + 1(at. wt. O)
• MW = 2(1.01 u) + 1(16.00 u) = 18.02 u
MOLECULAR WEIGHT PRACTICE
• The clear liquid is carbon disulfide, CS2. It is composed of
carbon (left) and sulfur (right). What is the molecular weight for
carbon disulfide?
• Answer: MW = 1(atomic weight C) + 2(atomic weight S)
12.01 u + 2(32.07 u) = 76.15 u
ISOTOPES & ATOMIC WEIGHTS
• Many elements occur naturally as a mixture of several
isotopes.
• The atomic weight of elements that occur as mixtures of
isotopes is the average mass of the atoms in the isotope
mixture.
• The average mass of a group of atoms is obtained by dividing
the total mass of the group by the number of atoms in the
group.
• A practical way of determining the average mass of a group of
isotopes is to assume the group consists of 100 atoms and
use the percentage of each isotope to represent the number of
atoms of each isotope present in the group.
ISOTOPES & ATOMIC WEIGHTS (continued)
• The use of percentages and the mass of each isotope leads to
the following equation for calculating atomic weights of
elements that occur naturally as a mixture of isotopes.
• According to this equation, the atomic weight of an element is
calculated by multiplying the percentage of each isotope in the
element by the mass of the isotope, then adding the resulting
products together and dividing the resulting mass by 100.
ISOTOPES & ATOMIC WEIGHTS EXAMPLE
• A specific example of the use of the equation is shown below
for the element boron that consists of 19.78% boron-10 with a
mass of 10.01 u and 80.22% boron-11 with a mass of 11.01u.
• This calculated value is seen to agree with the value given in
the periodic table.
Example 2.36, 2.37, & 2.39
THE MOLE CONCEPT
• THE MOLE CONCEPT APPLIED TO ELEMENTS
• The number of atoms in one mole of any element is called
Avogadro's number and is equal to 6.022x1023 .
• A one-mole sample of any element will contain the same number
of atoms as a one-mole sample of any other element.
• One mole of any element is a sample of the element with a mass
in grams that is numerically equal to the atomic weight of the
element.
• EXAMPLES OF THE MOLE CONCEPT
• 1 mole Na = 22.99 g Na = 6.022x1023 Na atoms
• 1 mole Ca = 40.08 g Ca = 6.022x1023 Ca atoms
• 1 mole S = 32.07 g S = 6.022x1023 S atoms
THE MOLE CONCEPT (continued)
• THE MOLE CONCEPT APPLIED TO COMPOUNDS
• The number of molecules in one mole of any compound is
called Avogadro's number and is numerically equal to
6.022x1023.
• A one-mole sample of any compound will contain the same
number of molecules as a one-mole sample of any other
compound.
• One mole of any compound is a sample of the compound
with a mass in grams equal to the molecular weight of the
compound.
• EXAMPLES OF THE MOLE CONCEPT
• 1 mole H2O = 18.02 g H2O = 6.022x1023 H2O molecules
• 1 mole CO2 = 44.01 g CO2 = 6.022x1023 CO2 molecules
• 1 mole NH3 = 17.03 g NH3 = 6.022x1023 NH3 molecules
THE MOLE CONCEPT (continued)
• THE MOLE AND CHEMICAL CALCULATIONS
• The mole concept can be used to obtain factors that are
useful in chemical calculations involving both elements and
compounds.
One mole quantities of six
metals; top row (left to right):
Cu beads (63.5 g), Al foil (27.0
g), and Pb shot (207.2 g);
bottom row (left to right): S
powder (32.1 g), Cr chunks
(52.0 g), and Mg shavings
(24.4 g).
One mole quantities of four
compounds: H2O (18.0 g);
small beaker NaCl (58.4 g);
large beaker aspirin, C9H8O4,
(180.2 g); green (NiCl2 · 6H2O)
(237.7 g).
MOLE CALCULATIONS
• The mole-based relationships given earlier as examples for
elements provide factors for solving problems.
• The relationships given earlier for calcium are:
1 mole Ca= 40.08 g Ca = 6.022x1023 Ca atoms
• Any two of these quantities can be used to provide factors for
use in solving numerical problems.
• Examples of two of the six possible factors are:
and
MOLE CALCULATION EXAMPLE
• Calculate the number of moles of Ca contained in a 15.84 g
sample of Ca.
• The solution to the problem is:
• We see in the solution that the g Ca units in the denominator of
the factor cancel the g Ca units in the given quantity, leaving
the correct units of mole Ca for the answer.
MOLE CALCULATIONS (continued)
• The mole concept applied earlier to molecules can be applied
to the individual atoms that are contained in the molecules.
• An example of this for the compound CO2 is:
1 mole CO2 molecules = 1 mole C atoms + 2 moles O atoms
44.01 g CO2 = 12.01 g C + 32.00 g O
6.022x1023 CO2 molecules = 6.022x1023 C atoms +
(2) 6.022x1023 O atoms
• Any two of these nine quantities can be used to provide factors
for use in solving numerical problems.
MOLE CALCULATION EXAMPLES
• Example 1: How many moles of O atoms are contained in
11.57 g of CO2?
MOLE CALCULATION EXAMPLES
• Example 1: How many moles of O atoms are contained in
11.57 g of CO2?
• Note that the factor used was obtained from two of the nine
quantities given on the previous slide.
MORE MOLE CALCULATION EXAMPLES
• Example 2: How many CO2 molecules are needed to contain
50.00 g of C?
MORE MOLE CALCULATION EXAMPLES
• Example 2: How many CO2 molecules are needed to contain
50.00 g of C?
• Note that the factor used was obtained from two of the nine
quantities given on a previous slide.
MORE MOLE CALCULATION EXAMPLES
• Example 3: What is the mass percentage of C in CO2?
• The mass percentage is calculated using the
following equation:
• If a sample consisting of 1 mole of CO2 is used, the molebased relationships given earlier show that:
1 mole CO2 = 44.01 g CO2 = 12.01 g C + 32.00 g O
MORE MOLE CALCULATION EXAMPLES
(continued)
• Thus, the mass of C in a specific mass of CO2 is known, and
the problem is solved as follows:
MORE MOLE CALCULATION EXAMPLES
• Example 4: What is the mass percentage of oxygen in CO2?
• The mass percentage is calculated using the following
equation:
• Once again, a sample consisting of 1 mole of CO2 is used to
take advantage of the mole-based relationships given earlier
where:
1 mole CO2 = 44.01g CO2 = 12.01 g C + 32.00g O
MORE MOLE CALCULATION EXAMPLES
(continued)
• Thus, the mass of O in a specific mass of CO2 is known, and
the problem is solved as follows:
• We see that the % C + % O = 100% , which should be the case
because C and O are the only elements present in CO2.
Example 2.42
MOLE CALCULATIONS HELP
Example 2.44, 2.53, 2.54 & 2.60