Download Subatomic Particles,Average atomic Mass

• After various molecular models of the atoms had
been tested, it was determined that three
subatomic particles made up the atom
• Protons
• Neutrons
• Electrons
• Protons are found in the nucleus
• Protons have an actual charge of +1.6 x 10-19 C
and a relative charge of +1
• The actual mass of a proton is 1.67 x 10-24 g
• The relative mass of a proton is 1 atomic mass
unit (amu)
• The symbol is p+
• Neutrons are found in the nucleus
• Neutrons have an actual charge of 0 C and a relative
charge of 0
• The actual mass of a neutron is 1.67 x 10-24 g
• The relative mass of a neutron is 1 atomic mass unit
(amu)
• The symbol is n0
• Electrons are found outside the nucleus
• Electrons have an actual charge of -1.6 x 10-19 C
and a relative charge of -1
• The actual mass of an electron is 9.11 x 10-28 g
• The relative mass of a electron is 1/2000 atomic
mass unit (amu)
• The symbol is e-
• Each element has a certain number of protons in its
nucleus
• The number of protons in the nucleus is called the
atomic number
• Each element has its own atomic number because
each element has its own, unique number of protons
• Note: That an element only changes if the atomic
number changes. Which means that if the number of
protons changes it changes elements
• Which element has
• a) 87 protons
• b) 35 protons
• c) 50 protons
• d) 92 protons
• e) 8 protons
• f) 19 protons
• On Your Own:
• Determine the number of protons in the following
atoms as well as each atom’s identity
• a) 6 electrons
b) 14 electrons
• c) 72 electrons
d) 55 electrons
•Mass Number: the number of
protons and neutrons in an atom
added together
•Mathematically
•Mass Number = Protons +
Neutrons
• What is the mass number of an
atom with 16 protons and 16
neutrons?
•Determine the Mass Number for the
following atoms
•a) 17 protons and 18 neutrons
•b) 11 protons and 12 neutrons
•c) 1 proton and NO neutrons
•d) 3 protons and 4 neutrons
• Mass Numbers (This is represent by an A)
are written in the upper left preceding the
chemical symbol
• Atomic Number (This is symbolized by a
letter Z) is written directly under the mass
number
 Example: Write the correct nuclide
symbol for an element with
51 p+ and 71 n0
Practice
• Draw the following Nuclide Symbols for
the following:
• Sr
• Os
• Zn
•I
• Cs
•K
Practice
• Write the following elements in proper Nuclide
Symbols:
• Pt
• Fr
• At
•K
-Write the correct Nuclide Symbol for
the following elements:
*19 p+, 20 n
*82 e-, 125n
*Mass # 238, neutrons= 146
• Atoms of the same elements have the
same number of protons
• HOWEVER there may be different
numbers of neutrons and different mass
numbers
• When an element’s atom has different
numbers of neutrons, it is said to have
isotopes
• Hydrogen has the following isotopes:
• Protium-a hydrogen atom with one proton and
NO neutrons
• Deuterium-a hydrogen atom with one proton
and only one neutron
• Tritium-a hydrogen atom with one proton and
two neutrons
•Atomic mass is the mass of
an atom expressed in
atomic mass units or amu
•The atomic mass unit is
based on the relation of
standard carbon-12
•Carbon-12 has a mass of 12.000 00 amu
Example:
If an atoms weighs half as much as
carbon-12, its atomic mass will be 6.000
amu
Example:
If an atom weighs four times as much as
carbon-12, it will have a mass
of 48.000 00 amu
What is Atomic Mass?
•The atomic mass that is reported
in the periodic table is a weighted
average based on the relative
abundance of each element
•Relative abundance
refers to how common
the isotope occurs in
nature
•Percent Abundance which
refers to how many of
each isotope are in every
hundred
• 1) First convert relative abundance (%)
to decimal equivalent
• 2) Multiply mass (in amu) by decimal
equivalent
• 3) Add the numbers together
• 4) The sum (in amu) is the average
atomic mass
•For example, an element has two
naturally occurring isotopes. One
isotope has a relative abundance of
19.91% and a mass of 10.012 amu. A
second isotope has a relative
abundance of 80.08% and a mass of
11.009 amu. Calculate the atomic
mass
• For example, an element has two
naturally occurring isotopes. One
isotope has a relative abundance of
92.58% and a mass of 7.02 amu. A
second isotope has a relative
abundance of 7.42% and a mass of
6.02 amu. Calculate the atomic mass
Calculate the average atomic masses for the following:
Isotope:
Rel. Abund.
Rel. Mass
hydrogen-1
99.985%
1.008
hydrogen-2
0.015%
2.014
Practice
•Titanium has five common
isotopes: If the abundance of Ti46 is 8.0%, Ti-47 is 7.8 %, Ti-48
is 73.4 %, Ti-49 is 5.5% and Ti50 is 5.3 %. What is the
average atomic mass of
titanium?
Determine Avg. Atomic Mass for oxygen:
• Isotope Rel. Abund.
Actual Mass
• O-16
99.762
15.995
• O-17
0.038
16.999
• O-18
0.200
17.999
•
Review 2/11/2015
• What is an Isotope?
• What are hydrogen’s three isotopes
names?
• What is the difference between average
atomic mass and mass number?
Review Practice
• Rubidium is a soft, silvery-white metal that
has two common isotopes, 85Rb and 87Rb. If
the abundance of 85Rb is 72.2% and the
abundance of 87Rb is 27.8%, what is the
average atomic mass of rubidium?
Element Atomic Mass p
#
#
Br
+
n
0
e
-
85
80
121
39
19
Practice
Element Atomic Mass # p+
#
Ag
n0
e-
110
19
40
127
53
Practice
Finding % Abundance
Silver( Atomic Weight 107.868) has two naturallyoccurring isotopes with weights of 106.91 and
108.90. What is the percentage abundance of the
lighter isotope?
Practice
Oxygen comes in two stable isotopes, Oxygen- 16
and Oxygen- 18. The molar mass of Oxygen- 16 is
15.99 amu. The molar mass of Oxygen-18 is 17.99
amu. Determine the percent abundance of each
isotope.
Finding Percent Abundance
The element indium exists naturally as two
isotopes. 113In has a mass of 112.90 amu, and
115In has a mass of 114.90 amu. The average
atomic mass of indium is 114.82 amu. Calculate
the percent relative abundance of the two
isotopes of indium.
Review
Antimony has two naturally occurring
isotopes. The mass of antimony-121 is
120.90 amu and the mass of antimony-123
is 122.90 amu. Using the average mass
from the periodic table, find the abundance
of each isotope.
Practice
Europium has two stable isotopes: Eu-151
with a mass of 150.92 amu and Eu-153 with
a mass of 152.92. If elemental Europium is
found to have an average mass of 151.964
amu on earth, calculate the percent of each
of the two isotopes