Chemistry Notes: Chapter 3
Law of Conservation of Mass—Mass is neither created nor destroyed during ordinary chemical reactions or
changes.
Law of Definite Proportions—Any chemical compound contains the same elements in exactly the same
proportions by mass regardless of the size of the sample or source of the compound.
Water is always 11% Hydrogen and 89% oxygen by mass.
Law of Multiple Proportions—If two or more different compounds are composed of the same two elements, the
ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of
whole numbers. (read page 66)
Dalton’s Atomic Theory: p. 66
Explains why all chemical reactions involve the combining, separating and rearrangement of atoms. The atoms
are not subdivided, created, or destroyed.
Not all aspects of this theory are correct. Atoms can be subdivided into smaller particles. However, we still use
today the major components of Dalton’s atomic theory.
Section 3-2
Atom—smallest particle of an element that retains the chemical properties of that element.
Atoms contain:
1. Nucleus—small region located near the center.
a. Contains the protons, which have a positive charge.
b. Also contains the neutrons, which have a neutral charge.
2. Electron cloud—large area surrounding the nucleus that contains negatively charged particles called
electrons.
a. Electrons were discovered with cathode rays (read p. 70)
b. Electrons have a large charge for their tiny mass.
c. Mass of an electron is 9.109 x 10-31 kg.
3. Because atoms are electrically neutral, they contain equal numbers of positive and negative charges.
4. Most of the mass of the atom is located in the nucleus.
5. Rutherford discovered the”+” charged center of the atom. He also discovered that this center was
very small compared to the size of the entire atom.
6. Protons and neutrons have almost the same mass but both are much more massive than an electron.
7. Different elements differ in the number of protons in the nucleus. Therefore, the number of protons
determines the atom’s identity.
8. When two protons are extremely close to one another there is a strong force of attraction (gravity)
between particles. This force along with neutron-neutron forces holds an atom together. They are called
Nuclear forces.
9. Because atoms are so small, their radii are expressed in picometers (10-12) pm.
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Atomic number—number of protons in the nucleus of an atom of that element. Atomic number identifies the
element.
All atoms of the same element contain the same number of protons. However, they can contain varying
amounts of neutrons. Because they contain different amounts of neutrons, their atomic masses are different.
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Isotopes are atoms of the same element that have different atomic masses because of different numbers of
neutrons.
Mass number—total number of protons and neutrons in the nucleus of an isotope. Hydrogen has three isotopes:
Protium has no neutrons and a mass # of 1.
Deuterium has 1 neutron and a mass # of 2
Tritium has 2 neutrons and a mass # of 3.
All three have 1 proton.
Specifying their mass numbers identifies isotopes. There are two methods for specifying isotopes:
1. Mass number is written with a hyphen after the name of the element.
(Uranium-235)
2. Shows the composition of a nucleus as the isotopes “nuclear symbol”.
235
92U
Nuclide—general term for any isotope of any element.
Knowing the atomic number and the mass number of a particular element allows one to determine the number
of protons, neutrons, and electrons of that element.
Atomic number = #of protons = # of electrons
Mass Number- atomic number = number of neutrons.
Because the mass of atoms is so small, we use relative atomic masses for calculations. To use relative atomic
masses, one element had to be designated as the standard and all other atoms are compared to it.
Carbon-12 was designated as the standard. Carbon-12 was assigned a mass of exactly 12 atomic mass units
(amu’s). Therefore, 1 amu is 1/12 of a carbon-12 atom. All other atoms are compared to this. Since Mg is
almost twice the mass of Carbon-12, its mass is 23.985042 amu’s.
Protons and neutrons are both almost equal to 1 amu. Electrons are much, much smaller.
Average atomic mass—weighted average of the atomic masses of the naturally occurring isotopes of an
element. (read page 80)
Mole (mol) amount of a substance that contains as many particles as there are atoms in exactly 12 g of carbon12. There are exactly 6.0221367 x 1023 carbon atoms in a mol of Carbon-12.
A mole is a counting unit similar to a dozen or a gross.
Avogadro’s Number—number of particles in exactly one mole of a pure substance
(6.022 x 1023)
Molar mass—mass of one mole of a pure substance. It is written in units of g/mol. It is equal to the atomic
mass only it is written in grams instead of amu’s.
Gram/Mole conversions—see page 82.
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