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Atoms: The Building Blocks of Matter
3 Guiding Principles Chemistry:
1. Law of Conservation of
Mass/Matter
2. The Law of Definite Proportions
3. The Law of Multiple Proportions
Dalton’s Atomic Theory
Problems with Dalton’s Atomic
Theory:
Matter can neither be created nor destroyed
o You can not make an atom appear or disappear
o The atoms you start with in a reaction must be
accounted for in the product
- A chemical compound contains the same elements in the
same ratio regardless of mass of sample
o A tsp of water, a glass of water, and a swimming pool
of water all have the ratio of 2H : 1 O regardless of size
of sample
- If two or more different compounds are composed of the
same elements the ratio of the elements in those compounds
will always be simple whole number ratios
o You can never have half an atom or part of an atom
the ratio of elements must always be the lowest whole
numbers.
1. All elements are made of tiny indivisible atoms. Atoms can not be
created nor destroyed
2. Atoms of the same elements are identical in size, mass, and other
properties. Every element has its own unique atom.
3. Atoms can combine to form compounds
4. Atoms can combine, separate, or be rearranged during a
chemical reaction
-
1. We have discovered that atoms can be broken down into
subatomic particles: protons, neutrons, and electrons
- We have also learned that quantum particles exist: quarks,
leptons, positrons etc.
2. Atoms of an element are not always identical. Some elements
have atoms with different masses. These atoms are called isotopes
The Subatomic Particles:
The Electron
-
Discovered in 1897 by J.J. Thomson
-
Ran a charge through the cathode-ray tube and a beam of
light appeared.
Checked to see if particles in the beam were positive or
negative
Knew that like charges repel, opposite charges attract
When he placed a positively charged plate near the tube, the
beam was attracted to it.
When he placed a negatively charged plate near the tube,
the beam was deflected away.
Cathode Ray Tube Experiment
-
o Based on the above he concluded the beam contained
negatively charged particles
o Called the negative charges Electrons
Facts about electrons
---------------------------------------THE PROTON
The Rutherford Gold-Foil
Experiment
-
Electrons are small negatively charged particles
Electrons are 10,000 times smaller than protons/neutrons
(9.109 x 10 -31 Kg = 0 amu)
- Electrons are responsible for the bonding that occurs among
atoms
- Electrons float around the nucleus in an area called the
electron cloud
- Electrons carry a full 1- charge
-------------------------------------------------------------------------------- Discovered in 1911 by Ernest Rutherford
-
3 important findings of Gold-Foil
Experiment
Facts about the Proton
Major contribution of the proton
The atomic number
Wanted to know if and where any positive particles existed in
the atom
- Placed radioactive Polonium in a lead box.
o Polonium gives off tiny positive particles called alpha
particles
- If positive alpha particles escape the box and pass straight
through the gold-foil to the X-ray film, the atom must not
contain any positive charges (like charges repel)
- If some of the positive alpha particles are deflected backward,
then the atom must contain positive charges – this theory
turned out to be true
o Most particles went straight through , but a few were
deflected
------------------------------------------------------------------1. Atoms contain positively charged particles – Protons
2. Atoms are mostly empty space in which electrons can roam
about
3. Atoms are located in a densely-packed, positively charged,
central core – called the Nucleus.
-
The proton is a positively charged particle- carries a full 1+
charge
- The proton lives in the nucleus of the atom
- The proton has a mass of 1.673 x 10-27 kg = 1amu
o Amu stands for: atomic mass unit
o 1 amu = 1.67 x 10 -27 kg
- A proton and an electron have opposite charges which can
cancel each other out and make the overall charge neutral
-----------------------------------------------------------The number of protons in an atom are unique to each element
- the number of protons in the atom
o serves as the atoms fingerprint
-------------------------------------------------The Neutron
o identifies one element from another
o elements are organized on the periodic table according
to their atomic number
o atomic # = # of protons
------------------------------------------------------------------------------------- discovered by James Chadwick in 1932
- Has approximately the same mass as a proton
o 1.675 x 10 -27 kg = 1 amu
- Found in the nucleus
- Adds weight/heft to the nucleus
- Has no charge (composed of equal amounts of positive and
negative quantum particles) = neutral
- Partially responsible for radiation
- The number of neutrons in an element may be different 
creates isotopes
The Nucleus:
The mass of atom is located in the nucleus
The nucleus has a net positive charge
Contains protons and neutrons
Held together by atomic forces
Densely- packed core: contains all the mass of the atom, but
takes up very little space (if atom is the size of a football
stadium, the nucleus would be a marble sitting in the middle of
the field)
2 Masses associated with an element
- Mass number
- Average atomic mass
-
How is the mass of an atom
determined?
Mass Number
-
Formula
-
Solving for neutrons or protons:
o
Average Atomic Mass
Let’s review what isotopes are.
-
Determined by the total number of protons and neutrons in an
atom
Mass number = # protons + # neutrons
Explanation: I know that…
o 1 proton weighs 1 amu
o 1 neutron weighs 1 amu
o 1 electron has no real mass (negligible) so we do
not count it
o If I have 4 protons they must weigh 4 amus and if I
have 6 neutrons they must weigh 6 amus.
Therefore the mass of the atom is the total of 4 amu
+ 6 amu = 10 amu
If I know the mass of an atom is 39 amus and I know that its
atomic number is 19, I can solve for the number of neutrons by
plugging the given values into the formula.
 Mass # = #protons + # neutrons

39 =
19
+ X
 Solve for X
 X = 39 - 19 = 20. There are 20 neutrons in this
atom.
Isotopes: Atoms of the same element that have different masses.
o Isotopes of the same element must always have
the same number of protons
o Isotopes can have different numbers of neutrons
and different mass numbers
(30 represents the mass of isotope E, the 15
represents isotope E’s atomic number
3017E
3515E
3512E
Example 1: 3015E
A.
B.
C.
D.
o
Example of an Isotope Question
Which two letters represent
isotopes of the same element?
Answer:
-----------------------------------------------Let’s review the common math
formula for finding the average of
a list of numbers:
3015E
The correct answer is A and C. A and C both have an atomic
number 15, which signifies that they are the same element. However,
both A and C have different masses  they are isotopes of that
element.
------------------------------------------------------------------------------------------------Most of us learned that to find the average of a given series of
numbers, we must add the numbers together and divide by the total
number of numbers.
For Example: Find the average of 3, 5, 4, and 2
3 + 5 + 4 + 2 = 14 divide by the number of numbers (4)
14/ 4 = 3.5 therefore 3.5 is the average
But why do we divide by 4? We divide by 4 because each number
in our series makes up an equal percentage in our average. In other
words each number makes up 25% of the total.
Let’s find the average in a different way. This time instead of dividing
by 4 at the end, let’s multiply by the percentage weight of each
number.
(4 x 25/100) + (3 X 25/100) + (5 X 25/100) + (2 x 25/100) = 3.5
The number is the same.
What do you do if the numbers
being added together are not
represented equally?
Formula for average atomic mass:
Example 1:
What is the average atomic mass
of carbon if C-12 makes up ~ 98%
of all carbon atoms and C-14
makes up only ~2%?
In nature, the majority of the isotopes of an element do not exist in
equal amounts. For example Carbon -12 makes up ~ 98% of all
carbon atoms, while Carbon -14 only makes up ~2% of the total
carbon atoms. If I were to find the average of both isotopes in the
traditional way, the average mass would be 13. This, however, is not
correct; the majority of carbon atoms have a mass of 12, so why
should Carbon-14 weigh so heavily on the average. To deal with this
inequality we use the average atomic mass formula.
(mass of isotope A x % abundance) + (mass of isotope B x % abundance) + etc.
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Solving Example 1
(12 amu x 98) + (14 amu x 2 ) = 11.76 + .28 = 12.04
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