Download 04 Mass Spectrometer and Isotopes

Isotopes & Mass
Spectrometry
Unit 10
Thank You Scientists!
 We
now have a modern model of the
atom!!!
Atom
Neutral
Nucleus
Outer Area
Center
Protons
Positive
Charge
Neutron
No Charge
Electrons
Negative
Charge
Atomic Mass Unit (AMU)
g/mol
 Atomic
number = # of protons
 The # of protons defines the element.
 The amount of electrons and neutrons in
an element may change, BUT if the
number of protons changes, it becomes
a new element.
 In a neutral atom, proton = electrons
 Atomic mass = protons + neutrons
What information do these
numbers give us about
aluminum?
X = symbol of the element.
A = atomic or molar mass of the
element.
Z = atomic number or the number
of protons.
 Elements
are defined by the number of
protons in the nucleus:
 Hydrogen has
proton.
 Boron has
protons.
 Thallium has
protons.
 What else contributes to the mass of the of
the element?
 Neutrons! Neutrons contain pretty much the
same amount of mass as protons; electrons
contain very little mass (small enough that
we can ignore their mass)
How
would we determine the
number of neutrons?
How would we determine the
number of electrons for a neutral
atom of this element?
What would indicate that the atom is
not neutral?
2+
How
many electrons does this ion
have?
Isotopes
 Isotopes
are atoms of the same element
(same # protons) with a different # of
neutrons. Therefore they have different
atomic masses.
 Because isotopes have different atomic
masses, this is usually indicated in the
name. Example Hydrogen-2 This
isotope has a mass of 2 g/mol.
 The atomic mass of an element is a
weighted average based on relative or
percent abundance of all of the
isotopes that exist for that element.
Example of Calculating Atomic Mass
 You
are given a sample containing 98% carbon12 and 2% carbon-13. What is the atomic mass of
the element?
 98% of the atomic mass will come from carbon-12.
12 x 0.98 = 11.76 g/mol
Carbon-12 will contribute 11.76 g/mol
 2% of the atomic mass will come from carbon-13.
13 x 0.02 = 0.26 g/mol
Carbon-13 will contribute 0.26 g/mol
 The atomic mass of the sample is found by adding
the contribution of each isotope’s weighted mass.
11.76 g/mol + 0.26 g/mol = 12.02 g/mol
How do we measure the relative
abundance of isotope?
Imagine: You are pushing a shopping cart filled to the
max with groceries (heavy ones). You are going to
be pretty late for your party if you don’t hurry.
 A little girl jumps out at the end of the aisle and you
have to turn sharply to miss her

 Would
the situation have
been different if your buggy
was not quite so full?
 How easily you can change
the direction of your
movement is affected by
how much mass you have.
 Isotopes have different
masses! Maybe we can use
this to determine the mass of
different isotopes!
So how does it work with
atoms:
We
use an instrument called a
Mass Spectrometer
It measures the atomic mass
and the relative abundance of
each isotope in a sample.
Step 1: Ionization
 First
we need to charge our atoms.
 What do we call a charged atom?
Ions with +1
Charge
Step 2: Acceleration
 The
ions are then accelerated so that
they are traveling at the same speed.
Step 3: Deflection
A
magnetic force is applied to the ions
traveling through the tube.
 The smaller the mass of the atom, the more the
atom is deflected.
+
Ions with +1
Charge
-
Step 4: Detection
What is the molar mass of the isotope
27 g/mole
represented by spectrum A?
What are the name and atomic symbol
of element A?
What are the symbols, including superscripts
and subscripts for the isotopes in spectrum
B?
Based on the experimentally obtained values of
atomic mass and percent abundance, calculate
the average molar mass of this element. Show
your work.
(0.787 x 24 ) + (0.103 x 25) + (0.112 x 26) =
24.4 grams/mole