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Chem 1151: Ch. 2
Atoms and Molecules
“It’s a theory not a fact”
• Hypothesis: A possible explanation.
• Theory: A proposed explanation that is consistent with all
available evidence that has been repeatedly tested.
• A theory evolves from interpretation of data (facts).
• “Theory” is frequently misunderstood by the general public.
• Gravity (Newton and Einstein) is “just a theory”, but I predict
that if you jump from the roof of your house, you will fall.
• The scientific method is a sequence of steps for systematically
analyzing the natural world.
The origin of atoms
Following the big bang
 Expansion of space
 Cooling
 Formation of fundamental particles
 Formation of low mass nuclei (H and He)
 Star formation
 Fusion reactions forming elements up to Iron-56
http://www.lbl.gov/abc/wallchart/chapters/10/0.html
Structure of the Atom
Mass (g)
Mass (u)
Proton (p+)
1.67 x 10-24
1
Neutron (n)
1.67 x 10-24
1
Electron (e-)
9.07 x 10-28
1/1836
Most of the mass is actually in the nucleus
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
4 Fundamental Forces in the Universe
These forces generally considered manifestations of the same force. Unified.
http://hyperphysics.phy-astr.gsu.edu/hbase/forces/funfor.html#c4
Synthesis of Matter
• Nucleosynthesis: Protons and
neutrons join to form nuclei.
• Fusion: Multiple nuclei join to
form heavier nucleus.
• Formation of heavier elements:
• Two protons collide
– Releases positron, neutrino
• Nucleus with proton and neutron
collides with another proton
– releases gamma ray
• Two He-3 atoms collide
– Produces He-4
– Releases two protons
http://en.wikipedia.org/wiki/Image:FusionintheSun.png
Abundance of Elements in the Universe
Mg
K
Cd
• Fuse 3 4He atoms together  12C, basic building block of all organic
compounds  life.
• Elements with FW or AMU multiples of 4 are more abundant: (C, O, Si,
Ca, etc.)
• Fusion Process continues up through Fe.
http://en.wikipedia.org/wiki/File:SolarSystemAbundances.png
Periodic Table of the Elements
• All matter in our
universe categorized in
periodic table.
– Based on atomic
number (number of
protons).
• Arranged in columns
(groups) and rows
(periods).
– Groups have similar
properties.
– Periods correspond
to filling of quantum
shells by electrons.
Elements from Group 7A
chlorine
bromine
iodine
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Elemental Notation
http://hyperphysics.phy-astr.gsu.edu/HBASE/pertab/pertab.html
Applications of Atomic and Mass Numbers
– 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.
– 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!
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Elemental Notation
Q: How to represent element X with 4
p+ and 5 n.
Q: How to represent lead-208?
Q: How many p+, e-, n?
82, 82, 126
9
4
208
82
X
Pb
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.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Isotope Symbols
A
• Isotopes are represented by the symbol Z
E , where Z is the
atomic number, A is the mass number, and E is the elemental
symbol.
• 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.
Relative Masses and Mass Units
• 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.
• 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 because carbon12 has 6 protons and 6 neutrons.
Proton (p+)
Neutron (n)
Electron (e-)
Mass (g)
1.67 x 10-24
1.67 x 10-24
9.07 x 10-28
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Mass (u)
1
1
1/1836
Determining Mass
12
However, carbon has an actual
mass listed of 12.011 u, not 12 u.
Why are these values different?
C
6
12.011
Mass (g)
Mass (u)
Proton (p+)
1.67 x 10-24
1
Neutron (n)
1.67 x 10-24
1
Electron (e-)
9.07 x 10-28
1/1836
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
How Isotopes Determine Atomic Weight
• The atomic weight of an element is the relative mass of an
average atom of the element expressed in atomic mass units.
• Many elements have more than 1 isotope (e.g. – 12C, 13C, 14C).
• Abundance of isotopes are not evenly distributed.
• Weighted atomic mass of Carbon (12C, 13C only) = (0.98882*12u)
+ (0.01108 * 13.300335u) = 12.011u.
12
12
CC
66
12.011
12.011
AMU
12 u
13.300335 u
isotope % isotope mass 
Atomic weight 
100
Determining Atomic Weight
• 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.

19.78% 10.01 u  80.22%11.01 u)
AW 
100
198.0 u  883.2 u

 10.81 u
100
• This calculated value is seen to agree with the value given in
the periodic table.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
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
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Molecular Weight
• 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
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
The Mole
• 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
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
The Mole
• 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
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
What these numbers mean?
Mass of 1 atom of Mg = 4.037 x 10-23 g
Mg atomic weight = 24.31 u
How many atom of Mg in24.31 g Mg?
24.31 g Mg x 1 atom Mg
=
4.037 x 10-23 g
6.022 x 1023 atoms Mg
Mass of 1 atom of C = 1.994 x 10-23 g
C atomic weight = 12.01 u
12.01 g C
x
1 atom C
=
1.994 x 10-23 g
6.022 x 1023 atoms C
The number of atoms of an element, in a mass equivalent to it’s atomic weight,
is equal to Avogadro’s number (mole).
or
1 u = 1 g/mole
Relationships: Mass, Moles, Molecular Weight
1 mol C atoms = 6.022 x 1023 atoms C
6.022 x 1023 atoms C = 12.01 g C
1 mol C atoms = 12.01 g C
C has atomic weight of 12.01 u or 12.01 g/mol
1 mol S atoms = 6.022 x 1023 atoms S
6.022 x 1023 atoms S = 32.1 g S
1 mol S atoms = 32.1 g S
S has atomic weight of 32.1 u or 32.1 g/mol
Compound (Molecular) Formulas
Compound formula: all elements and number of each in a compound
Examples:
Urea
Hydrofluoric acid
Sodium bicarbonate
Sodium Azide
1C, 4H, 2N, 1O
1H, 1F
2H, 1C, 3O
1Na, 3N
The compound (molecular) chemical formula represents the numerical
relationships that exist between atoms in a compound. This also applies
to moles.
1 mol of H2SO4 contains
2 mol of H
1 mol of S
4 mol of O
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.
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011