Download Prerequisite Knowledge for Chemistry

Review of the Chemistry You Learned in Physical Science
Important vocabulary terms are underlined the first time they appear in this review
Matter

All matter is made of atoms. Atoms are not the smallest form of matter, but atoms are
the smallest particles of matter that retains a chemical identity (still can be considered
an element and display many of the properties of that elements).
Matter including liquids are made of smaller units made of bonded atoms.
These small units can be called molecules or more generally compounds.
These smaller units which are made of individual atoms can be broken down
into single atoms. When groups of atoms are broken many or all of the
properties of that group / molecule are lost.

Elements are any amount of one type of atom. Elements are a form of pure substance.

Compounds are made of two or more elements joined by chemical bonds. Compounds are a
form of a pure substance.

Mixtures are different elements or compounds NOT chemically bonded.

There are three states of matter: Solid, Liquid and Gas. To compare the characteristics
of each state of matter to each other think about a single substance (like water).
Shape
Volume
Energy per Unit Mass
Degree of Order
Solid
Retains shape
Maintains volume
Lowest
Highest
Liquid
Takes shape of
container
Maintains volume
Moderate
Low
Gas
Takes shape of container
Takes volume of container
Highest
Lowest
Atomic Structure

Atoms are made of three subatomic particles: Electrons, Protons and Neutrons. Below
are table describing the mass, charge and location of each particle.
Particle
Proton
Neutron
Electron
Symbol
p
n
e-
Relative Charge
+1
0
-1
Relative Mass (AMU)
1
1
0 (actually 0.000545)
Location
Nucleus
Nucleus
Energy Levels

Almost all the mass of an atom is contained within the nucleus. Electrons practically have no
mass 0.000545 AMU (atomic mass units) compared to the 1 AMU of protons and neutrons.
By adding the protons and neutrons you can calculate the mass of the atom. For instance, an
atom with 9 protons and 10 neutrons in its nucleus would have a mass of 19 AMUs.

Atoms have a core occupied by neutrons and protons we call the nucleus. The nucleus is a
very small part of the atom by volume, but the nucleus contains nearly all the mass of the
atom. Positioned around the nucleus are electrons which are in energy levels. The lowest
energy level is closest to the nucleus. The highest energy level (the 7th) is farthest away.
The electron energy levels occupy almost all of the volume of the atom, but contain
essentially no mass (remember electrons nearly 0 mass compared to protons & neutrons).
Model of Chlorine Atom

All atoms of one element will have the same number of protons. If you change the number of
protons then the element changes. For instance, all carbon atoms have 6 protons. If a
proton is added to a carbon atom, the atom would become nitrogen. All atoms with 7 protons
are nitrogen.

If an atom has the same number of protons (positively charged subatomic particles) and
electrons (negatively charged subatomic particles) then the atom is neutral (no charge).

Atoms can lose or gain electrons and remain the same element.

Ions are atoms with an unequal number of protons and electrons. When there are more
protons than electrons the atom is positive and we call the atom a “cation”. When there are
more electrons than protons the atom is negative and we call the atom an “anion”.
A neutral atom of sodium would have 11
protons and 11 electrons. If a sodium atom
were to lose a negatively charged electron as
shown in the diagram the resulting atom
would have a positive charge because it
would have one more proton than electron.

Positives and negatives attract. Therefore, protons and electrons attract. Also, positive
atoms (cations) and negative atoms (anions) attract.

Like charges repel. Protons repel other protons. Electrons repel electrons.
Cations repel other cations. Anions repel other anions.

Isotopes are atoms of the same element with varying number of neutrons & mass. For
instance, there are three isotopes of hydrogen. The most common isotope does not have
a neutron. There are less common isotopes of hydrogen with one neutron & two neutrons.
Periodic Table

Elements are all given abbreviations consisting of one or two letters (e.g. Carbon is “C”).
These abbreviations are called element symbols. Not all element symbols are obvious (e.g.
Sodium is “Na”). The first letter of an element symbol is always capitalized and the
second (if there is one) is always lower case.

Atomic Number is the number of protons contained by an atom of a given element.
Atomic number is almost always included in a periodic table. Atomic number is always a
whole number as you cannot have a fraction of a proton.

The periodic table is ordered from left to right and down by increasing atomic number.

Atomic Mass is the “weight” of an atom. Atomic mass is also most often included in
periodic tables, but does not have to be a whole number.

Columns of elements in the periodic table are called families. Families of elements have
similar properties (e.g. state of matter, reactivity, tendency to form a positive or
negative ion, etc.).
Chemical Formulas and Chemical Equations

Compounds are represented by chemical formulas consisting of element symbols and
subscripts. For instance, water’s chemical formula is H2O. The “H” stands for hydrogen
and the “O” stands for oxygen. The subscripts come after the element to which they
refer. Subscripts state the number of atoms of that element in the compound.
Therefore the “2” in H2O says that there are two atoms of hydrogen in water. If there
is not a subscript shown (as for oxygen), then the assumption is a “1”. So, there is only
one atom of oxygen. Again H2O means 2 atoms hydrogen and 1 atom of oxygen.

Coefficients can be placed in from of a chemical formula to indicate a some multiple of
that formula. For instance, 3 H2O is equal to 3 molecules of H2O.
Note:

3 H2O = H2O + H2O + H2O
Chemical equations are representations of chemical reactions written out in symbol form.
For instance, the reaction of Octane (C8H18, also known as gasoline) and Oxygen in a car’s
engine can be written out as:
2 C8H18 (g) + 25 O2 (g)  16 CO2 (g) + 18 H2O (g)

The chemicals before the arrow in a chemical equation are called reactants and those
after the arrow are called products. The reactants in the octane reaction above are
C8H18 and O2 and the products are CO2 and H2O.

Arrows in a chemical reaction can be thought of as an equal sign and also as an indicator
of change (transformation).
Chemical Reactions

During chemical reactions, the bonds between atoms are broken and reformed. A
chemical reaction has happened if the formulas change. Using the octane reaction again,
we notice that the formulas C8H18 and O2 do not appear in the product side of the
reaction equation. The product formulas are different (CO2 and H2O) and so we may say
that a chemical reaction has occurred.

Chemical reactions follow the Law of Conservation of Matter which says that atoms are
not created or destroyed during a chemical reaction. This means that there must be the
same number of each type of atom on either side of the equation’s arrow. To make sure
we follow the Law Conservation of Matter, coefficients in a chemical equation can be
changed. Consider the following reaction of sodium (Na) and chlorine molecules (Cl2) to
become sodium chloride (table salt). There are no coefficients shown so we may assume
that they are all “1”.
Na + Cl2  NaCl
(or
1 Na + 1 Cl2  1 NaCl but we rarely write the “1”s)
The equation above contains the correct formulas, but it does not follow the Conservation
of Matter because there are two atoms of Chlorine on the reactant side and only one on
the product side. We CANNOT change the formula of sodium chloride to NaCl2 because
no such compound exists. Instead, we place coefficients out front of the compounds so
we have the correct number of atoms.
2 Na + Cl2

2 NaCl
Notice that when a “2” is placed in front of the sodium chloride to balance the chlorine
atoms, another would be needed in front of the sodium on the reactant side to balance
sodium atoms.

Chemical reactions overall either release energy (exothermic) or absorb energy
(endothermic). The burning of octane (gasoline) is a good example of an exothermic
reaction while an instant cold pack is a good example of an endothermic reaction.
Heat and Kinetic Theory

Heat is defined as the motion (translation, vibration and rotation) of atoms and
molecules.

Changing the states of matter (solid, liquid, gas) of a substance requires energy to be
absorbed or released. Moving from a solid to a liquid and eventually a gas requires an
absorption of energy (this change would be endothermic). Reversing the direction going
from a gas to liquid to solid requires that energy is released by the substance (this
change would be exothermic). Remembering that solids have the least amount of energy
per mass and gases have the most energy per mass might help you remember whether
energy needs to be absorbed or released.