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Atom
3-1 Atom
450 BC Democritus: Atom (Atomos)– smallest particle of an element that retains the chemical
identity of that element.
1799 AD Joseph Proust: Law of Constant Composition – Compound always contains same
elements in same proportion by mass.
1803 AD John Dalton: Atomic Theory of Matter
- Each element is composed of extremely small particles called Atoms.
- All atoms of given element are identical, but differ from any other
element.
- Atoms are neither created nor destroyed in any chemical reaction.
- A given compound always has the same relative numbers and kinds of
atoms.
3-2 Atomic Structure – Originally thought of as small and hard like marbles.
1839 Michael Faraday: Atoms contain particles that have electrical charge
1896 JJ Thomson determined the charge of an electron
1909 US Millikan measures the mass of an electron
1896 Becquerel discovers radioactivity.
Atom now thought of as like a chocolate chip cookie – round with subatomic particles randomly
throughout the atom.
1909 Rutheford proposes nucleus in center of atom:
electrons
Nucleus
Modern Theory of Atom
3-3 The Structure of the Atom - Atomic Number and Ions
Atoms are composed of:
Nucleus
Electron Cloud
Proton; p+
Neutron; no
Electron; e-1
1.673x10-24 g
1.675x10-24g
9.109x10-28 g
Mass
Volume
Diameter of atom approximately 0.1 – 0.5 nanometers
Atomic Number
1910 Henry Mosely found that atoms of each element contain a unique positive charge in their
nucleus
- The number of protons is represented by the atomic number:
7
N
Nitrogen
14.01
-
Since an element has a neutral charge, the atomic number must also represent the
number of electrons in a neutral element
- Elements have the ability to lose or gain electrons in order to bond with other
elements to form molecules.
- An element with an imbalance of protons to electrons is called an ion.
- It is then said to have an Ionic Charge
Ex) When Nitrogen is an ion, it has an ionic charge of -3 which means that it has 10
total electrons
number of protons
7
- number of electrons 10
charge of ion
-3
We would write the Nitrogen Ion as: N-3
Isotopes
As stated above, all atoms of the same element have the same number of protons. They do, vary
in the number of neutrons.
-
-
These atoms are referred to as Isotopes.
Most elements in the first two rows of the periodic chart have at least two isotopes of
which one is the most common.
Since the mass of an atom is based on the mass of the nucleus, we can determine the
number of the neutrons an atom has from the following:
Number of Neutrons = Atomic Mass- Atomic Number
If you look at boron, you might wonder, “how there can be a fractional neutron?”
-
There cannot be – The atomic Mass is an average of all of the masses of all of the
isotopes of that element found on Earth.
In fact, Hydrogen has 3 isotopes: Hydrogen, Deuterium, and Tritium..
We would write them:
o Hydrogen - 1
o Hydrogen - 2
o Hydrogen – 3
-
Notice that the number behind the isotope is the Atomic Mass, so, these numbers
come from using the above formula:
Number of Neutrons = Atomic Mass- Atomic Number
Hydrogen – 1 no = 1 – 1 = 0
Hydrogen – 2 no = 2 – 1 = 1
Hydrogen – 3 no = 3 – 1 = 2
-
To determine the number of Protons, p+; Electrons, e-1; and Neutrons, no each of these
have, we get:
p+1
e-1
no
Hydrogen – 1
1
1
0
Hydrogen – 2
1
1
1
Hydrogen – 3
1
1
2
We can also write these isotopes as follows:
1
1
H  Hydrogen - 1
2
1
H  Hydrogen - 2
3
1
H  Hydrogen - 3
where the top number is the atomic mass of the isotope and the bottom number
represents the atomic number.
Mass of Atom
We can determine the mass of an atom by using the number of protons and neutrons an atom has.
- # protons (mass of proton) + # neutrons (mass of neutrons) = mass of atom
- The mass of an atom uses two different mass scales.
- One is the metric unit grams
o p+ = 1.673x10-24 grams
o no = 1.675x10-24 grams
o e- = 9.102x10-28 grams
ex) the mass in grams of one atom of chlorine can be determine as follows:
# protons = atomic number = 17
# neutrons = atomic mass – atomic number = 35 – 17 = 18
(17)(1.673x10-24) + (18)(1.675x10-24) = 5.86x10-23 grams
the other mass scale is called Atomic Mass Units (AMU’s)
o this scale is based on the mass of 1 proton ≈ 1 amu
o therefore the mass of 1 neutrons ≈ 1 amu
o the mass of 1 electrons would be ≈ 0 amu’s
ex) the mass in amu’s of one atom of chlorine would then be:
(17 protons)(1 amu) + (18 neutrons)(1 amu) = 35 amu’s
-
3-4 Changes in the Nucleus
Nuclear Stability:
- for elements with atomic numbers between 1 and 20, the number of neutrons to
protons are about the same.
- Those elements between 20 and 83 require an increasingly larger ratio of neutrons to
protons.
- Elements beyond 83 naturally encounter nuclear decay.
- However, not only does an isotope of an element need to have fewer neutrons than
protons to be radioactive, it also can encounter nuclear decay when there is a large
excess of neutrons as compared to the number of protons.
Types of radioactive decay
- We can identify 5 types of radioactive particles which can be emitted from a nucleus.
- These are:
o Alpha particles, , which are very heavy and dangerous particles, 24 He , also
called a Helium atom
o Beta particles, , which are medium in both size and hazard, 10 e or 10 e also
called an electron or positron
o Neutron, very hazardous, 01 n
o
o
Proton, also hazardous, 11 H
Gamma particle very light high energy particle, no charge or