What Do I Need to Know for the EOC Chemistry?
*These are really important...
Reporting Category 1: Matter
12 questions
*Know the Difference Between Physical and Chemical Changes and Properties...
 Chemical Properties: anything related to a chemical change...flammability,
toxicity, pH...
 Physical Properties: anything related to a physical change...color, texture, density,
mass, volume, solubility...
 Mixtures and solutions can be separated by PHYSICAL means because they are
PHYSICALLY combined.
 Know that physical changes DO NOT change what a substance is made of, so it’s
formula won’t change. Changing phases is physical (melting, freezing,
sublimation) because it is still the same stuff. Dissolving is physical!
 Know your phase changes: freezing, melting, boiling, sublimation, condensation
 Know that during a phase change, the temperature DOES NOT CHANGE
Know the difference between Extensive and Intensive Properties....
 Intensive properties are INDEPENDENT of amount. That is, the property doesn’t
change if the amount changes.
 Intensive properties are used to IDENTIFY substances since they don’t change.
Example: Does the density of an piece of wood change if you cut it in half?
NO—both pieces still float!
So, density is an INTENSIVE property.
Example: Does the heat capacity of a piece of metal change if you cut it in
half?
YES—each piece contains half of the heat the original piece
contained.
So, heat capacity is an EXTENSIVE property.
 If you are still confused, just memorize these EXTENSIVE properties:
Mass, volume, length, width, height, heat capacity, weight
Know the Difference Between Solids, Liquids, and Gases....
Solid, Liquid, Gas
Solid: Definite shape and volume, poorly compressible, little movement of
particles, low energy of particles
Liquid: Indefinite shape and definite volume, slightly compressible, some motion
of particles, some energy of particles
Gases: Indefinite shape and volume, very compressible, constant motion of
particles, high energy of particles
*Classify Matter as Pure Substances or Mixtures Through the Investigation of
Their Properties....
 Elements and compounds are PURE SUBSTANCES, and therefore have a formula
that NEVER changes: H2, H2O, sugar, alcohol, pure metals
 Heterogeneous mixtures and homogeneous solutions are NOT pure. They are
physically combined and can be separated into their components.
 Solutions can be aqueous (dissolved in water), gas mixtures (air), or metal
solutions (alloys).
 Know your separation methods: distillation, sifting, filtering, evaporation,
decanting, centrifugation
 Remember that pure substances always have a specific melting/freezing and
boiling point. Boiling point is dependent on the air pressure in addition to
temperature. In a heating curve diagram or cooling curve graph, the flat parts
represent phase changes.
 Intensive properties such as density or specific heat can be used to identify pure
substances.
Explain the Use Chemical and Physical Properties in the Historical Development
of the Periodic Table:
 Dimitri Mendeleev developed the first periodic table in 1860’s by organizing the
elements in order of atomic mass. He realized that there were properties that
changed periodically every so many elements. He called these “periods”.
 So, he began putting the periods on top of one another so that there were columns
that had similar chemical and physical properties.
 He predicted the properties of 3 elements that had not been discovered yet!
 His predictions were based on the fact that his table had trends in various
properties. So, he realized that he could look at the elements on all four sides of
the missing element and predict what it would be like.
*Use the Periodic Table to identify and explain the properties of chemical
families, including alkali metals, alkaline earth metals, halogens, noble gases, and
transition metals.
 Know the location AND name of each of the families listed above.
 Know the charge each family forms AND how to tell where metals and nonmetals
are found.
 Know that each family has a specific set of properties it is associated with:
Alkali metals—1 valence electron, very reactive, reactivity increases down the family,
react violently in water to form hydrogen gas, form alkaline (basic) solutions, react with
halogens to form salts or ionic compounds, form +1 ions, low ionization energy, all have
electron configuration that ends in s1. The fact that they have only one valence electron
is the reason for their reactivity—it is very easy to get rid of that one electron so that
there will be a full shell underneath.
Alkaline earth metals—2 valence electrons, reactive (but less than the alkali metals),
reactivity increases down the family, react with water to form slightly less alkaline
solutions, react with halogens to form ionic salts, form +2 ions, low ionization energy, all
have electron configurations that end in s2.
Transition metals—valence electrons can vary between 1-3, much less reactive than
Groups 1-2, form ionic salts that are usually brightly colored, at the right end of the
period some are almost nonreactive, hard, dense, form + 1 to +5 ions, and some can
have more than one oxidation number (for example, Fe+2 or Fe+3). All of their electron
configurations end in a d-orbital.
Halogens—7 valence electrons, all form -1 ions. Very reactive nonmetals that react
with the alkali and alkaline earth metals to form ionic salts (“halogen” means “saltformer”). They get less reactive down the family, with Fluorine being the most reactive.
High ionization energies and electronegativities. All have electron configurations that
end in p5. They only need one more valence electron to make an octet, which accounts
for their reactivity.
Noble Gases—8 valence electrons (a full octet), so these gases are very nonreactive
(inert), and do not form ions. All end in the electron configuration p6. Remember that
Helium has only 2 electrons, but that is a full shell for Helium since it only has one
energy level.
*Use the Periodic Table to identify and explain periodic trends, including atomic
and ionic radii, electronegativity, and ionization energy.
 Know the properties of metals and nonmetals.
Metals: lusterous, malleable, ductile, dense, high melting points, almost all solids,
+ ions, reactivity increases down a family, low ionization energies
Nonmetals: brittle, colored, low dense, low melting points, reactivity decreases
down a family, solids, liquids, or gases, high ionization energies
 Know the definitions of the trends:
Electronegativity: tendency for an element to attract electrons within a chemical
bond
Atomic radii: size of an atom from the center to the outside
Ionic radii: size of an atom AFTER it becomes an ion
Ionization energy: energy required to REMOVE an electron and form an ion
Metallic character: how “metallic-like” an atom is
 Know your trends above...
Electronegativity: Fluorine is the highest—so the closer you are to Fluorine, the
higher it is. The noble gases have ZERO electronegativity.
Atomic radii: increases all the way down to Francium, Fr, in the bottom left
corner.
Ionization Energy: Increases all the way up to Helium, which is the highest.
Metallic character: the same as Atomic radii—with Francium having the most
metallic character.
Density: Increases down a group—as does melting and boiling points.
Ionic radii: see the chart on the next page! Atoms that form + ions get smaller,
while atoms that form – ions get larger.
The left hand side of the staircase forms small, + ions that are SMALLER than the atom.
The right hand side of the staircase forms larger – ions that are LARGER than the atom.
 Understand the reasoning behind the trends...this will get you more points!
Atomic radii increases down a group because you are adding energy levels, but
decreases across a period because you are adding protons to the nucleus, but not
energy levels, so the electrons are pulled in toward the nucleus.
Ionization energy: increases as you go across a period toward the right because
the number of valence electrons is increasing, and atoms are getting closer to a
full octet. It becomes more difficult to remove an electron the closer the electrons
are to the nucleus, AND the closer to an octet the atom gets.
Electronegativity: the electronegativity of an element depends on three things:
the number of protons in the nucleus, the distance from the nucleus and the
amount of shielding by inner electrons. Electronegativity and ionization energy
decrease down a group because the outer shell electrons are getting farther away
from the nucleus. Electronegativity and ionization energy INCREASE across a
period from left to right because protons are being added to the nucleus, but no
energy levels—so electrons are being pulled in by the increasing positive nucleus.
Ionic radii: Positive ions are smaller than the atoms they came from because they
lose electrons—but not protons. The fewer electrons are pulled in by the more
positive nucleus. The opposite is true of negative ions—they have added
electrons, and there isn’t adequate protons to pull the electrons inward, so they
“stray” away from the nucleus, and the ion expands.