Atomic Structure

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Atomic Structure
Atomic Number
• Definition: Equals the number of protons
– Identifies the element
• # protons = # electrons in a neutral atom (an
atom without a charge)
Isotopes & Mass Number
• Isotopes = atoms with the same # of protons,
but different numbers of neutrons
Ex: Isotopes
3 isotopes of hydrogen:
1.
Hydrogen-1; mass # = 1; 0 neutrons
2.
Hydrogen-2 (deuterium); mass # = 2; 1
neutron
3.
Hydrogen-3 (tritium); mass # = 3; 2
neutrons
There are two different ways to write
chemical symbols for isotopes:
1. Write the mass number after the element’s
name (gold-197)
2. Use the symbol, a superscript for mass #, and a
subscript for atomic #
Mass Number = protons + neutrons
• To calculate # neutrons
= mass # - atomic #
• Use shorthand writing of atom’s composition
• Superscript = mass #
• Subscript = atomic #
Practice 1
Practice 2
Practice 3
a) Three isotopes of sulfur are sulfur-32, sulfur33, and sulfur-34. Write the complete
symbol for each isotope, including the
atomic number and the mass number.
b) How many neutrons, protons, and electrons
are in Na+ with a mass number of 24? What
is its atomic number?
Atomic Mass
• Weighted atomic mass of all the isotopes of that
element – that is why it is a decimal, not an integer
• It reflects mass and relative abundance of the isotopes
as they occur naturally
AMU – the mass of an atom
• Amu’s are used to define the weight of a single
atom or isotope
• How measured: Mass spectrometer –
determines masses of atoms
• Carbon-12 was set to 12 atomic mass units
(amu’s)
– Used to define masses of atoms
Mass Number and Atomic Mass
Carbon-12 = 6 protons and 6 neutrons,
– mass # = 12,
– atomic mass = 12 amu, (atomic # = 6)
• 1 proton = 1 amu
• 1 neutron = 1 amu
• Mass number = Atomic Mass FOR A SINGLE ATOM
ONLY
• The defined mass of each elements is actually a
weighted average so it is NOT a whole number.
Average Atomic Mass: Uses a weighted
average to account for mass and
relative amounts
• Most elements are present as a mixture of 2+
isotopes
• Not all isotopes are present in equal amount
(percent abundance)
• Ex – hydrogen
Mass defect: the mass of any nucleus is less than the
sum of the separate masses of its protons and neutrons.
How can the actual mass be less than the mass
number? Can atoms lose mass?
•
E=mc2... Some of the mass of the protons and neutrons get
converted to energy when they come together in the
nucleus
Calculating average atomic mass
• Atomic mass = (Mass isotope A * natural
abundance of A) + (Mass isotope B * natural
abundance of B) + …
• You must know the following in order to
calculate the atomic mass of an element:
– # of stable isotopes
– mass of each isotope
– % abundance of each isotope
Chlorine: Example 1
• Use the data from the previous slide on
chlorine to determine the average atomic
mass.
Magnesium Example
Mass Number
% Abundance
24
78.99
25
10.00
26
11.01
Calculating Relative Abundance from
Average Atomic Mass
• The element Quahog has two isotopes:
Quahog-35 and Quahog-40.
– If the atomic mass of Quahog is 39.00 amu, what
is the relative abundance of each isotope?
• The sum of the relative abundance must equal
1.
• So assign X and 1-X as the relative abundance
of each and solve.
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