Chem 1151: Ch. 2
Atoms and Molecules
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
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 heavier elements
http://www.lbl.gov/abc/wallchart/chapters/10/0.html
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
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
Elements You Should Know
ELEMENT
Aluminum
Argon
Arsenic
Barium
Bromine
Cadmium
Calcium
Carbon
Cesium
Chlorine
Chromium
Cobalt
Copper
SYMBOL
Al
Ar
As
Ba
Br
Cd
Ca
C
Cs
Cl
Cr
Co
Cu
ELEMENT
Fluorine
Gold
Helium
Hydrogen
Iodine
Iron
Lead
Lithium
Magnesium
Manganese
Mercury
Neon
Nickel
SYMBOL
F
Au
He
H
I
Fe
Pb
Li
Mg
Mn
Hg
Ne
Ni
ELEMENT
Nitrogen
Oxygen
Phosphorous
Potassium
Silicon
Silver
Sodium
Strontium
Sulfur
Tin
Zinc
SYMBOL
N
O
P
K
Si
Ag
Na
Sr
S
Sn
Zn
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
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.
Since all of the mass of atom comes from the protons and the
neutrons, the mass number is the total number of protons and
neutrons.
You can therefore determine the number of neutrons by
subtracting the atomic number from the mass number.
Use of 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
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
12
6
%Abundance
AMU
C
Carbon-12
98.892
12 u
13
6
C
Carbon-13
1.108
14
6
C
Carbon-14
1.0 x 10-10
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%10.01u  80.22%11.01 u)
AW 
100
198.0 u  883.2 u

 10.81u
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 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
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
Problems
1. What is the mass in g of 1.35 mol of S?
Mass S? Mol g
1.35 mol S
X
32.1 g S
1 mol S
=
43.3 g S
2. How many S atoms are in 98.6 g of S?
98.6 g S X 1 mol S X 6.022 x 1023 atoms S
32.1 g S
1 mol S
= 1.85 x 1024 atoms S
3. What is the mass in g of 1 atom of S?
1 atom S X 1 mol S
X 32.1 g = 5.33 x 10-23 g S
6.022 x 1023 atoms S
1 mol S
Moles of Molecules
1. What is the mass in g of 1.62 mol of O2 molecules?
MW of O2 = 2 x (atomic weight of O) = 2 x (16.0 u) = 32.0 u
1 mol O2 molecules = 6.022 x 1023 molecules O2
6.022 x 1023 molecules O2 = 32.0 g O2
1 mol O2 molecules = 32.0 g O2
1.62 mol O2 molecules X 32.0 g O2 = 51.8 g O2
1 mol O2 molecules
1 mol O2 molecules = 12.04 x 1023 atoms O
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 molecule of H2SO4 contains
2 atoms of H
1 atom of S
4 atoms of O
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
Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011
Finding the Molecular Weight
*Use the mole relationship of the compound to find the MW
Ex. 1 H2SO4
Atomic Weights
H = 1.008 g/mol
S = 32.06 g/mol
O = 16.00 g/mol
MW of 1 mol of H2SO4 = (2 * H = 2.016 g/mol) + (1 * S = 32.06 g/mol) + (4 * O =
64.00 g/mol) = 98.22 g/mol
Ex. 1 C3H8O (isopropyl alcohol)
Atomic Weights
C = 12.01 g/mol
H = 1.008 g/mol
O = 16.00 g/mol
MW of 1 mol of C3H8O = (3 * C = 36.03 g/mol) + (8 * H = 8.064 g/mol) + (1 * O =
16.00 g/mol) = 60.09 g/mol
Finding the Number of Atoms in a Compound
Ex. 1 How many C atoms in 5.2 g of C3H8O (isopropyl alcohol)
Atomic Weights
C = 12.01 g/mol
H = 1.008 g/mol
O = 16.00 g/mol
MW of 1 mol of C3H8O = (3 * C = 36.03 g/mol) + (8 * H = 8.064 g/mol) + (1 * O =
16.00 g/mol) = 60.09 g/mol
1 mol of C3H8O contains 3 mol of C
Percent Composition
Mass relationships can be used to determine percent compositions
1 mol of H2SO4
%H=
2.0 g
98.1 g
x
100
=
2.0 %
%S=
32.1 g
98.1 g
x
100
=
32.7 %
%O=
64.0 g
98.1 g
x
100
=
65.2 %
% Mass of element in compound
% mass N in HNO3?
N = 14.01 g/mol
H = 1.008 g/mol
O = 16.00 g/mol
HN03 % mass= part
total
X 100
Total= (1 x H) + (1 x N) + (3 + O) =
(1 x 1.008) + (1 x 14.01) + (3 x 16.00) = 63.018 g/mol
HN03 % mass=
14.01 g/mol X 100 = 22.23%
63.018 g/mol
% Mass of element in compound
% mass N in NaN3?
N = 14.01 g/mol
Na = 22.99 g/mol
NaN3 % mass= part
total
X 100
Total= (1 x Na) + (3 x N) =
(1 x 22.99 g/mol) + (3 x 14.01 g/mol) = 65.02 g/mol
NaN3 % mass=
42.03 g/mol X 100 = 64.64%
65.02 g/mol
Atomic Weight of Element with Multiple Isotopes
Element X has 3 isotopes
(10)X 70% 41.00 u
(11)X 20% 42.00 u
(12)X 10% 43.00 u
What is atomic weight of element X?
(70% x 41.00 m) + (20% x 42.00 u) + (10% x 43.00 u) =
(70 x 41.00 m) + (20 x 42.00 u) + (10 x 43.00 u) = 41.40 u
100
Ions
 Ion: Atom or molecule that has either lost or gained electrons from
valence shell resulting in a net charge (positive or negative)
compared to the number of protons.
 Ions of element have the same number of protons, but a different
number of electrons.
 For example compare the following:
44
20
Ca
vs.
44
20
Ca 2
Common atomic ions you should know:
 H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+,
 N3-, P3-, O2-, S2-, F-, Cl-, Br-
Exercises
What are the charge, mass (u) in the following nuclei?
1. 5 p+, 6 n
2. 10 p+, 12 n
3. 11 p+, 14 n, 1eWhat are the p+, n, e-, charge?
22
10
Ne
44
20
Ca
40
20
Ca 2
What are the MW values in u?
C6H12O6 (glucose)
O3 (ozone)
Which is denser, U-235 or U-238? Why?
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