Transcript Counting Atoms

Counting Atoms
Section 3.3
Atomic Number
Atomic number: the number of protons
of each atom
Represented by “Z”
The atomic number identifies the element
Isotopes
Isotopes: atoms of the same element that
have different masses
Isotopes have the same number of
electrons and protons, but a different
number of neutrons
Most elements consist of a mixture of
isotopes
Mass Number
Mass number: the total number of protons
and neutrons that make up the nucleus of
an isotope
Mass number = p+ + n0
Isotopes are usually identified by specifying
their mass number
Designating Isotopes
There are two methods to specify isotopes
Method 1 is called hyphen notation: write
the name of the element, then a hyphen,
then the mass number
Method 2 is called the nuclear symbol: in
front of the symbol, superscript the mass
number and subscript the atomic number
Hydrogen
Hydrogen is special in that it also has
specific names for each isotope
Hydrogen with only 1 p+ and no n0 is
called protium, hydrogen-1,
Hydrogen with 1 p+ and 1n0 is called
deuterium, hydrogen-2,
Hydrogen with 1 p+ and 2 n0 is called
tritium, hydrogen-3,
Practice: name the isotope both
ways
helium with 1 neutron
helium-3,
carbon with 7 neutrons
carbon-13,
oxygen with 8 neutrons
oxygen-16,
uranium with 142
neutrons
 uranium-234,
More Practice
How many protons, electrons, and neutrons
are there in an atom of chlorine-37?
17 electrons, 17 protons, 20 neutrons
How many protons, electrons, and neutrons
make up an atom of bromine-80
35 protons, 35 electrons, 45 neutrons
Relative Atomic Masses
If we express the mass of atom in grams, it
would be extremely small
Relative mass is used instead
Carbon-12 is the standard atom and has
been assigned a mass value of exactly 12
atomic mass units (amu)
More about Relative Atomic
Mass
All other atomic masses are determined by
comparing it with the mass of the carbon12 atom
See the chart on page 80 for some typical
atomic masses
The masses could also be expressed by
adding up the mass of the e-s, p+s, and n0s
Average Atomic Mass
Average atomic mass: the weighted average
of the atomic masses of the naturally
occurring isotopes of an element
The average atomic mass depends on both
the mass and the relative abundance of
each of the element’s isotopes
Calculating Average Atomic
Mass
Copper consists of 2 naturally occurring
isotopes
A sample of copper contains 69.15% of
copper-63 (62.929601 amu) and 30.85% of
copper-64 (64.927794 amu)
Change the percentages to decimals
Continued
Solve like this:
(0.6915)(62.929601 amu ) +
(0.3085)(64.927794 amu)= 63.55 amu
The average atomic mass of copper is
63.55 amu
Always round to 2 decimal places
The Mole
The SI unit for the amount of a substance
Mole: the amount of a substance that
contains as many particles as there are
atoms in exactly 12 g of carbon-12
It is a counting unit
Avogadro’s Number
Avogadro’s number: the number of
particles in exactly one mole of a pure
substance; 6.022 x 1023 particles
Molar Mass
Molar mass: the mass of one mole of a
pure substance
Usually written in units of g/mol
The molar mass of an element is
numerically equal to the atomic mass of
the element in atomic mass units (found on
the periodic table)
More about Molar Mass
Molar mass is usually rounded to two
decimal places
What is the molar mass of carbon?
12.01 g/mol
What is the molar mass of chlorine?
35.45 g/mol
Gram/Mole Conversions
The molar mass can be used as a
conversion factor
see the board and the book for sample
problems
Conversions with Avogadro’s
Number
Since 6.022 x 1023 is the number of
particles in a mole, that can be used as a
conversion factor
See the board and your book for sample
problems