The Structure of Matter
Physical Science
Chapter 6
Review
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Compound: atoms of two or more elements
that are chemically combined
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Most of the matter around us is a compound or a
mixture of compounds
Compounds have properties unlike those of their
elements
During a chemical change, a new substance is
produced.
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Chemical bonds
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Forces that hold together the atoms in a
compound
When atoms gain, lose, or share electrons
they are forming chemical bonds.
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Chemical Structure
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The way the atoms are bonded in a compound
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Chemical formulas
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Used as shorthand for writing compounds.
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NaCl is sodium chloride
Subscript – means written below
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Tells us how many atoms of an element are in a
compound
If there is no subscript, then there is one.
Example: H2O has 2 atoms of hydrogen and one
atom of oxygen
The ratio of hydrogen atoms to oxygen atoms is 2
to 1
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Chemical structure representations
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Chemical formula – show how many of each
type of atom there is
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Water: H2O
Methane CH4
Structural formula – shows how atoms are
arranged.
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Chemical structure representations
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Space filling model – shows relative volumes
of the electron clouds.
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Ball-and-stick model – shows bond angles
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Effects of chemical bonds
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Compounds with strong chemical bonds
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Are rigid and difficult to break
Have high melting and boiling points
Compounds made of molecules
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Have strong bonds within each molecule
Have weak attractions between molecules
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Molecules are easy to separate
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Lower melting and boiling points
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Attractions between molecules
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Some molecules have stronger attractions
between them
Example: water
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Has hydrogen bonding between molecules
Why water has a relatively high boiling point for
a molecular compound
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Discuss
1.
Classify the following
as mixtures or
compounds
a.
b.
c.
d.
2.
Air
CO
SnF2
Pure water
3.
Draw a ball-and stick
model of a boron
trifluoride, BF3,
molecule. A boron
atom is attached to
three fluorine atoms.
Each bond angle is
120 degrees and each
bond is the same
length.
Predict which
molecules have a
greater attraction for
each other: C3H8O
molecules in liquid
rubbing alcohol or
CH4 molecules in
methane gas.
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Ionic compounds
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Ionic compound – a compound made up of
two or more ions
Form networks of ions, not individual units.
Ionic bond – the force that holds the ions in an
ionic compound together.
Ionic compounds have a net charge of zero,
so the compound is electrically neutral.
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Examples
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NaCl
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MgF2
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Ionic compounds
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Smallest unit is a formula unit.
Generally have high melting points and high
boiling points.
Are usually crystalline solids at room
temperature.
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Covalent compounds
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Covalent compounds are composed of
molecules that are created when atoms share
electrons
Covalent bonds – the bonds between atoms in
a molecule.
Molecules are also neutral.
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Examples
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HCl
Cl2
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N2

O2
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Covalent compounds
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Smallest unit is a molecule.
Generally have low melting points and boiling
points.
Are usually liquid or gaseous at room
temperature, but not always.
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Discuss
1.
Determine if the
2.
following compounds are
likely to have ionic or
covalent bonds.
a.
b.
c.
d.
Magnesium oxide, MgO
Strontium Chloride, SrCl2
Ozone, O3
Methanol, CH3OH
Identify which two of the
following substances will
conduct electric current,
and explain why.
a.
b.
c.
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Aluminum foil
Sugar, C12H22O11
dissolved in water
Potassium hydroxide,
KOH, dissolved in water
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Polar molecules
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Atoms in molecules don’t always share their
electrons equally.
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Examples

HCl

H2O
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Polar molecule
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Has a positive end and a negative end.
Example: stream of water
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Nonpolar molecules
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Do not have negative and positive ends.
Example: CO2
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Nonpolar vs. polar
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Metallic Bonds
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Occur between metal atoms
Atoms are closely packed together
Electron clouds overlap
Electrons move freely between atoms
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Polyatomic ions in compounds
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Poly means many
Polyatomic ions have more than one atom in
them.
See Figure 10 on page 190
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Compounds with polyatomic ions
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They form compounds just like monatomic
(one atom) ions do.
Examples

LiOH, lithium hyrdoxide

Mg(NO3)2, magnesium nitrate
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Discuss
1.
2.
3.
4.
5.
Compare bonds
Compare bonds
What is the difference between polar molecules and
nonpolar molecules?
What are polyatomic ions?
Identify which of the bonds in calcium hydroxide,
Ca(OH)2 are ionic and which are covalent.
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Naming Ionic Compounds
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List the (positive) cation first
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Name is usually the same as the element
List the (negative) anion second
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Change ending to –ide
See figure 2 on page 192
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Examples
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CaF2
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Li2O
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Calcium fluoride
Lithium oxide
K2S

Potassium sulfide
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Writing formulas for ionic
compounds
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The charge on the compound must add up to
zero.
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Add subscripts as needed
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Examples
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Cesium Oxide
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Beryllium chloride
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Cs2O
BeCl2
Calcium Phosphide

Ca3P2
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Transition Metals
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Groups 3 – 12
Can have more than one charge when forming
compounds
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Copper and oxygen can make CuO or Cu2O3
To name them, we need to specify the charge
of the cation using a roman numeral
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CuO is copper (II) oxide
Cu2O3 is copper (III) oxide
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Examples
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Titanium (III) nitride
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Fe2O3
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TiN
Iron (III) oxide
Iron (II) oxide

FeO
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Naming Covalent Compounds
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Different rules
Use numerical prefixes (figure 5 on page 194)
If there is only one atom of the first element, the
prefix mono- is omitted.
Change the ending of the second element to -ide
If the element starts with a vowel, drop the a or o
at the end of the prefix
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Example: tetroxide, not tetraoxide
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Examples
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PF5
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N2O5
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Oxygen difluoride
Phosphorus trichloride
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Dinitrogen pentoxide
OF2
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Phosphorus pentafluoride
PCl3
Dinitrogen pentoxide

N2O5
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Empirical Formulas
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Shows the smallest whole-number ratio of
atoms that are in a compound
Ionic compounds
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Almost always the same as the chemical formula
Covalent compounds
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Not always the same
Example: glucose chemical formula is C6H12O6,
empirical formula is CH2O
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Molecular formula
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Shows the actual numbers of atoms of each
type in one molecule
The same as the chemical formula
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Determining empirical formula
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Convert the mass of each element to moles.
Find the molar ratio, which gives the
empirical formula
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The ratio must be whole numbers, because the
subscripts in the formula must be whole numbers.
If the ratio isn’t whole numbers, multiply it by a
whole number to get rid of the fractions.
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Example: if the ratio is 1.5:1, multiply by 2 to get 3:2
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Examples
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A sample of an unknown compound has 36.04
g of carbon and 6.04 g of hydrogen. What is
the compound’s empirical formula?
A sample of a compound contains 3.6 g of
boron and 1.0 g of hydrogen. What is the
compound’s empirical formula?
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You try
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A certain compound sample contains 207.2 g of
lead and 32.00 g of oxygen. What is its
empirical formula?
A compound is analyzed and found to contain
36.70g potassium, 33.27g chlorine, and 30.03g
oxygen. What is the empirical formula of the
compound?
Find the empirical formula of a compound that
contains 53.70g iron and 46.30g sulfur.
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