Unit Three:
Molecules & Compounds
0
Unit Three: Molecules and Compounds
Table of Contents:
Pg. 1
Table of Contents
Pg. 2-3
Bonding (Intramolecular Forces)
Covalent and Ionic Bonding Assignment on Pg. 36
Pg. 4-5
Lab Two: Investigation into Polymers
Pg. 6-9
Ionic Compounds
Determining Number of Atoms & Criss Cross Method Assign on Pg. 37
Ionic Bonding Lewis Dot Assign Pg. 38
Determining Charges Assign Pg. 39
Pg. 10-11
Covalent Compounds
Covalent Bonding Lewis Dot Assign Pg. 40
Pg. 12-16
VSEPR
VSEPR Assign Pg. 41
Pg. 16-18
Intermolecular Forces
Ionic, Polar Covalent or Non-polar Covalent Assign Pg. 42
Intermolecular Forces Assign Pg. 43
Pg. 19-20
Nomenclature for Ionic Compounds
Naming Ionic Compounds Assign Pg. 44
Pg. 21-26
Lab Three: Formula Writing and Chemical Names
Pg. 27
Nomenclature for Covalent Compounds
Naming Covalent Compounds Assign Pg. 45
Writing Formulas From Names Assign Pg. 46
Pg. 27-28
Nomenclature for Acids
Naming Acids Assign Pg. 47
Pg. 29-32
Nomenclature for Organic Compounds
Naming Organic Compounds Assign Pg. 48
Structure of Organic Compounds Assign Pg. 49
Pg. 33-34
Solids (Independent Study)
Types of Solids Assign Pg. 50
1
Bonding
Recall:

Recall the structure of the atom:
o An atom is composed of 3 different particles: protons, neutrons and electrons

The main particle involved in bonding is the electron.

Recall the octet rule: Remember all atoms like to have 8 valence electrons in order to be stable.
Collision Theory:

Bonding between atoms requires things:
o Energy: Kinetic energy is required for bonding to occur. This kinetic energy is usually
obtained through collision (ie. the atoms or molecules bump into one another).
o Orientation: The way the atoms are placed during the collision matters. Not one
orientation works for all atoms/molecules (ie. The atoms or molecules need to bump
into one another in the right spot).
Bonding:

Bonding is when 2 or more atoms share or transfer electrons to try to fill or empty their valence
shells.

An atom with 6 valence electrons would want to gain or share 2 electrons in order to
Intramolecular Forces:

“Intra” is latin for “within”.

Intramolecular forces are forces within a molecule or compound holding it together.

There are 2 types of intramolecular forces: ionic bonding and covalent bonding.
2
Ionic Bonding:
o Occurs between a metal and a non-metal
o Occurs when electrons are transferred from one atom to another (from metal to the
non-metal).
o Ionic bonding creates ions

The metal loses an electron/ electrons and becomes a cation (so it is positive).

The non-metal gains an electron/ electrons and becomes an anion (so it’s
negative).
o Formula units are formed when ionic bonding occurs.
Covalent Bonding:
o Occurs between 2 non-metals.
o Occurs when electrons are shared between atoms.
o Molecules form when covalent bonding occurs.
o Covalent bonds can occur as single, double, triple or quadruple bonds.

Single bonds occur when one electron pair is shared between 2 atoms

Double bond occur when 2 electron pairs are shared between two atoms. It is
stronger than a single bond (but not 2 times stronger)

Triple bond occur when three electron pairs are shared between 2 atoms. Tripple
bonds are very strong bonds. One of the strongest bonds known is the triple
bond between nitrogen. This is why many explosives contain nitrogen. Because
when those compounds break apart and nitrogen gas is formed (N2), a large
amount of energy is released.

Quadruple bond are man-made and do not occur naturally. They can only be
observed under high vacuum conditions and are extremely unstable. This would
occur when 4 electron pairs are shared between 2 atoms.
Examples:
To determine whether a compound is ionic or covalent we look at the atoms that make up the
compound. Remember that every element symbol starts with a capital letter, so to determine how
many different elements there are in a compound, you just look at the number of capital letters.
To determine whether the compound is ionic or covalent, look at whether the elements are metals or
non metals:
NH3
NiBr3
MgCl2
Al(OH)3
H2S
(see covalent bonding and ionic bonding assignment)
3
Lab Two
Investigation into Polymers
Background information:
Polymers are long chains of atoms bonded together covalently. These long chains are made up of
repeating structural units (same order of atoms over and over). Since polymers are long chains, they
can get twisted and tangled together. If a substance can be twisted, pulled, or compressed and it
resumes back to its original form, it is usually a type of polymer called an elastomer. Elastomers have
elastic properties. That is they can by pushed or pulled and they still go back to their original form. An
example of an elastomer is a rubber band or a car tire.
The covalent bonds along the chain are strong, but the bonds between chains are normally weak.
However, additives such as borax allow the formation of strong "cross-links" between chains. As the
number of cross-links increases, the material becomes more rigid and strong.
The liquid latex, or glue, which we will use in this experiment, contains small globules of hydrocarbons
(hydrogens and carbons bonded together) suspended in water. These globules are polymers and are
called polyvinyl acetate. The polymer putty is formed by joining the globules using a cross linker
sodium tetraborate (found in borax). Sodium borate acts as a cross-linker to the original hydrocarbon
polymers. The cross-linker can be thought of as rungs joining two sides of a ladder. The putty is held
together by very weak intermolecular bonds that provide flexibility around the bond and rotation
about the chain of the cross-linked polymer.
4
Purpose: The objective of this experiment is to cross-link a polymer and observe the changes in the
physical properties as a result of this cross-linking.
Materials:
Plastic bag
Water
White liquid glue
Food colouring.
Borax
General Safety Guidelines:

Since borax solid (a bleaching agent) and solution will burn the eyes, goggles should be worn.

Hands should always be washed after kneading the polymers and finishing the experiment.
Procedure:
1. Obtain a plastic bag.
2. Obtain your sample of white glue in the plastic bag. Be sure to make your observations.
3. Place one drop of food colouring in the plastic bag. Close the back and mix (Observations)- Do
not use more than 2 drops of food colouring or that will ruin your experiment.
4. Obtain 4-5mL of 4% Borax solution. Slowly add this to your glue mixture while continually
stirring. (Observations, Observations, Observations!)
5. Make observations about what you have created. The Polymer is safe to touch, but not to eat or
taste because it is toxic if ingested. Get creative with your observations (Dropping it, stretching
slowly/quickly, compare your polymers with others in the class, etc.)
6. Place your polymer in the plastic zipper seal bag and seal it with the least amount of air
possible.
7. Clean up. Your polymer should go in the plastic bag and the bag should be sealed whenever you
are not observing.
8. Take your polymer home and place it in the fridge for about 15 minutes. Make observations
about your polymer when the temperature is decreased.
Analysis Questions:
1. What would be the effect (your thoughts) of adding more sodium borate solution?
2. Where does all of the water go? Remember that you stated with a solution made up of 96%
water.
3. How do the physical properties of the glue change as a result of adding the sodium borate?
4. Why does a car tire appear to be flat in the summer even though the gas inside is hotter than in
the winter.
5
Ionic Compounds
Chemical Formula
How to determine the charge of an ion:

If an atom gains electrons it will become negatively charged (an anion) because it is gaining
negative charges. If an atom loses electrons it is losing negative charges so it will be positive (a
cation).

For each of the atoms listed below, determine how many valence electrons they have, how
many valence electrons they would need to gain or lose to fill their valence shell, and what the
charge would be on the ion with a full valence shell.
Element
# of Valence Electrons
# of e- to gain or lose
Charge
Sulfur
Hydrogen
Sodium
Oxygen
How to Determine Chemical Formula for Ionic Compounds:

In order to determine the chemical formula for a compound we need to look at the charges that
the atoms would have with a full valence shell. For the transition metals we will always be
given its charge.
# of Atoms:

First we need to be able to read how many atoms are in a molecule:
o O=1atom
o O2= 2 atoms
o Na2O= 3 atoms
o 3Na2O=9 atoms (6 sodium, 3 oxygen)
o Mg(OH)2=5 atoms
o (NH4)2SO4= 16 atoms
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Chemical Formula for Binary Compounds:

To figure out how many of each atom goes into a molecule we use the criss-cross method:

Sr and Cl
Look on the common ion sheet to find the charges: Sr2+ Cl1Drop the Charge (leave the number): Sr2 Cl1
Criss Cross the numbers: Sr1Cl2
Ones do not need to be written: SrCl2

Ba and O
Ba+2 O-2
Ba2O2
BaO

Cu and F
Cu+2 F-1
Cu1F2
CuF2
Chemical Formula for Ionic Compounds that Contain Polyatomic Ions
For polyatomic ions, we look on the common ion sheet to find their charge. We use this charge in our
criss cross method.

Ammonium NH4+1 and Ferrocyanide Fe(CN)6 4Polyatomic ions need to have brackets around them (if there is more than one)
(NH4)4 Fe(CN)6

Magnesium and Carbonate

Cobalt (III) and Sulfate
(See Determining Number of Atoms
and Criss Cross Assignment)
7
Lewis Dot (Ionic)
How to determine the Lewis dot structure for ionic compounds
Example: Determine the lewis dot structure for the ionic compound that will form between calcium
and chlorine.
1. Determine the number of atoms of each element there will be (Criss Cross).
2. Draw out the lewis dot structure for each of the atoms. Circle the electrons that are going to
leave and create the cation and draw arrows to the anion they are going to.
3. Draw out the lewis dot for each of the ions created and put them in brackets with the charge in
the upper right corner outside of the brackets.
Examples:
Aluminum and Oxygen
(See Ionic Bonding Lewis Dot assignment)
8
Determining the Charge (Transition metals)
Determining charge when given the chemical formula:
Work backwards:

What is the charge on iron?
FeCl3
o We know chlorine has a charge of -1 and there are 3 chlorine ions, so copper must have
a charge of +3 in order for the compound to have a neutral charge.

What is the charge on copper? Cu(NO3)2
o

We know NO3 has a charge of -1 (from our common ion sheet), so copper must have a
charge of +2. This is written as copper (II) – For the transition metals.
What is the charge on Iron in Fe2O3?
o Oxygen has a charge of -2 normally. There are 3 oxygens so the oxygens have a total
charge of -6 in total. This means that the irons must have a total charge or +6. There are
2 iron atoms and the molecule is neutral so they must each have a charge of +3 in order
to equal +6.

What is the charge on iron in FeCO3?
o First we need to find the polyatomic ion (it won’t always be in brackets). A compound
contains a polyatomic ion if it is ionic and there are more than 2 elements. The
polyatomic ion will be the first few elements if it is a cation and the last few elements if
it is an anion. Most polyatomic ions are anions, so usually the polyatomic ion will be at
the end.
o Carbonate has a charge of 2- so iron has a charge of 2+.
Example:
PbCl2
As(OH)5
FePO4
(See Determining Charges Assignment)
9
Covalent Molecules
Number of Bonds
How to determine the number of covalent bonds that form:
Determining the number of covalent bonds that form is similar to determining charge. We look at how
many electrons an atom needs to fills its octet (this is the number of bonds that will form). When
atoms bond covalently they share their valence electrons.

Carbon has 4 valence electrons. It needs 4 electrons to fill its valence shell. This means it will
form 4 bonds (if they are single, or it could form 2 double bonds, or 1 triple bond and one
single bond).

Oxygen has 6 valence electrons and needs 2 to fill its valence shell. This means it will form 2
bonds (if they are single, or 1 double bond).

Nitrogen has 5 valence electrons and needs 3 electrons to fill its valence shell. This means it
will form 3 bonds (if they are single).
Lewis Dot (Covalent)
Draw Lewis dot structure for molecules (covalent):
Fluorine bonding with itself (F2):
1. Determine how many atoms of each element there will be (if it’s not given).
2. Draw the lewis dot for each atom. (In the space below)
3. Look at the lone electrons. These are the electrons involved in bonding. Circle one lone electron
from one atom and one lone electron from another. This will form a bond. If there are three or
more atoms, make sure every atom is involved in the bonding.
4. Redraw the molecule so that the atoms are beside one another and the shared electrons are
between them.
5. If asked to draw a structural diagram, draw lines to replace bonded pairs of electrons.
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Examples:
Oxygen bonding with itself:
Chlorine bonding with hydrogen:
CF4:
Nitrogen bonding with itself:
(See Covalent Bonding Lewis Dot assignment)
11
VSEPR
What do the molecules really look like?
Next we will look at what a molecule looks like in 3 dimensional space. We use molecular geometry
and VSEPR Theory. VSEPR stands for Valence Shell Electron Pair Repulsion. The VSEPR theory
determines the shape of a molecule by looking at the electrons surrounding the central atom and
whether they are shared pairs of lone pairs (bonding or non-bonding pairs).
Determine Molecular geometry using VSEPR:
1. Determine the lewis dot formula
2. Determine the total number of electron pairs around the central atom
3. Use the table provided to determine the electron pair geometry
4. Use the table provided to determine the shape
5. Use the diagram chart to draw a 3D diagram
VSEPR Chart
Total pairs
of e- around
central atom
Number of
bonded pairs
Electron Pair geometry
Molecular geometry
2
2
Linear
Linear
3
2
Trigonal Planar
Bent
3
3
Trigonal Planar
Trigonal Planar
4
2
Tetrahedral
Bent
4
3
Tetrahedral
Trigonal Pyramidal
4
4
Tetrahedral
Tetrahedral
5
2
Trigonal Bipyramidal
Linear
5
3
Trigonal Bipyramidal
T-Shaped
5
4
Trigonal Bipyramidal
Seesaw
5
5
Trigonal Bipyramidal
Trigonal Bipyramidal
6
2
Octahedral
Linear
6
3
Octahedral
T-Shaped
6
4
Octahedral
Square Planar
6
5
Octahedral
Square Pyramidal
6
6
Octahedral
Octahedral
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VSEPR Geometry 3D diagrams:
E = Central Atom
X =Bonded Atom
  = Unbonded pair of electrons
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Examples:
OF2
Oxygen has 4 pairs of electrons around it, 2 are bonded 2 are lone, so it has a geometry of tetrahedral
bent.
The 3D drawing for tetrahedral bent is:
So you just need to fill in the atoms where they belong:
PCl3
PCl3 has 4 pairs of electrons around it, 3 of the pairs are bonded. This means that PCl3 has a geometry
of Tetrahedral, Trygonal pyramidal.
14
Sometimes in a double or triple bond you have to assume that the atom has only one shared pair of
electrons in order to get the proper geometry.
CS2
Even though this actually has 4 shared pairs of electrons and all are bonding electrons, the shape is
linear. In this molecule we assume that Carbon has 2 pairs of electrons and both are bonded.
Normally assume that atoms want to have 8 valence electrons, however sometimes atoms like to have
more. In situations like this we say that the atom breaks the octet rule. In this class you will always be
told if an element breaks the octet rule.
XeF4
Xe has 8 valence electron and Fluorine has 7,
This molecule has 6 pairs of electrons around it, and 4 of them are involved in bonds so it’s geometry
is Octahedral, Square planar.
3D drawing:
(See VSEPR assignment)
15
Balloon Molecules Assignment
Read through the whole activity before you begin, as there is some after class participation that is
required.
Create the molecules assigned to your group using the balloons provided. There are 2 different
coloured balloons. Assign one colour to represent the bonded electron pairs and the other to
represent the lone electron pairs. Make sure you specify which is which.
Create your molecules so that it is stable enough to get knocked over and still maintain its shape. You
will be provided with tape and balloons. After you have created your molecule with the balloons, label
it with masking tape and your names. You then need to make a 3 dimensional drawing of each of your
molecules using the actual atoms (not drawings of the balloons). Use the dashed and solid lines to
represent the atoms that go into the page and the atoms that come out of the page. Each member of
your group should submit a drawing of both molecules. On this paper you need to write down which
balloon colors represent bonded electrons and which represent lone pairs. Don`t forget your names!
When you have finished building and drawing your molecules, you need to find a space on the counter
and place all of your molecule representations together. You can then work on the “VSEPR”
assignment.
After today`s class, I would like you to go to the virtual classroom and post a comment saying which
molecule your group had and what you think the geometry of that molecule is. Since each group has 2
molecules, each person should post a different molecule. For the group of 3, 2 of you will have to post
the same molecule.
Intermolecular Forces

Inter is latin for “between”

Occurs when bonds or forces occur between molecules (covalent).

Intermolecular forces hold a molecule in a certain spot.

The three types of intermolecular forces we will look at are polar bonds, hydrogen bonds and
van der waals forces.
Polarity

Non-polar molecules have equal, or approximately equal pull on the electrons. This occurs
when the atoms in the molecule are the same or have a similar electronegativity.
Electronegativity generally increases as you move from left to right in a period and decreases
as you move from the top to the bottom of a family.

Polar molecules have a slightly positive or slightly negative charge due to one atom having a
stronger pull on the electrons. This occurs when there is a large difference in electronegativity
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so one atom pulls stronger on the electrons than the other atom. The more electronegative
atom will have a partial negative charge because it has a stronger pull on the electrons, and the
less electronegative atom will have a partial positive charge because the electrons are being
pulled away.
o Example, Chlorine has an electronegativity of 3.16 and hydrogen has an
electronegativity of 2.0. Since chlorine has a stronger pull on the electrons they will be
closer to the chlorine atom and the chlorine atom will have a slightly negative charge.
This means the hydrogen will have a slightly positive charge.
o To show the direction of polarity we draw a dipole. A dipole is drawn to show where the
electrons in the molecule are more likely to be found. It is drawn like this:
o Another way to show the slight negative charge is to use the symbol that stands for
slightly.

The difference in the electronegativity can tell us what type of bond is formed.
o If the difference is 0-0.4 then the bond is non-polar covalent
o If the difference is 0.4- 1.7 then the bong is polar covalent
o If the difference in electronegativity is greater than 1.7, then the bond is ionic. If the
difference in electronegativity is very very large then it is usually an ionic bond because
one atom pulls so hard on the electrons that it pulls them completely away from the
other atom.

Remember that there is no distinct line of separation so the rules above are more like
guidelines. Remember an ionic bond will occur between a metal and a non-metal and covalent
is between 2 non metals.
H-I
(See Polar vs. Non-Polar Assign)
17
Hydrogen bonds:

Hydrogen bonds are weak bonds formed by the attractions of slightly negative atoms to the
slightly positive hydrogen when bonded covalently.

Hydrogen bonds occur between hydrogen, and Nitrogen, Oxygen or Fluorine.

Hydrogen bonds are also called H-bonds

Hydrogen bonds are weaker than covalent bonds but stronger than van der waals.
Examples: H2O
Van der Waals:

Van der Waals forces are a type of intermolecular force that deals with the dipoles of molecules.

The Van der Waals forces are those where a slightly negative atom from one molecule is
attracted to a slightly positive atom from another molecule.

Van der Waals forces are weaker than both the covalent bond and the hydrogen bond

(Technically H-Bonds and London Disperions forces are both types of Van Der Waals Forces)
Example: CO2
London Dispersion Force:

The weakest intermolecular force

Electrons in a molecule are constantly moving.

In a non-polar molecule, on average the electrons are distributed equally over the molecule, but
occasionally one side or the other will gain a small excess of electron density. When this occurs,
the molecule has a temporary dipole- one side of the molecule has a slight + charge, the other a
slight - charge.

All molecules have dispersion forces.
Example: Cl2
(see Intermolecular Forces Assign)
18
Nomenclature for Inorganic Compounds
IUPAC Naming System (Nomenclature)
IUPAC = International Union of Pure and Applied Chemistry
Naming Ionic Compounds:
When naming ionic inorganic compounds we always write the metal first (or the cation) and then the
non-metal (or the anion).
For compounds that do not contain polyatomic ions:
We first write the metal, then the non metal with “ide” added on.
Ex. Li2S=
AlCl3=
MgO=
For compounds that do contain polyatomic ions:
We write the metal then the polyatomic ion (do not add the “ide” ending)
Ex. K3PO4=
(NH4)2CO3=
Na2SiO3=
MgSO4 =
Transition Metals
When the metal involved in the ionic bond has more than one oxidations state/possible charge we
need to identify which ion it is. This usually occurs with the transition metals. Any ion that appears on
your common ion sheet more than once is one of these types of metals. There are 2 ways to do this.

The first way (the way that is the newer and preferred method) is to use roman numerals to
identify which ion it is.
o Ex. Fe(CN)3= First we need to find what charge iron has (Since cyanide has -1 charge and
there are 3 of them, iron must have a 3+ charge). We would then use the roman numeral
III to identify this ion.
o The rules are the same as above but now we add the roman numeral in the middle.
o The compound Fe(CN)3 would be named:

The second way to identify the ion is to use the -ous or -ic endings.
o The ion with the lower oxidation state (or lower numerical charge ignoring the + or -)
gets the –ous ending, and the ion with the higher oxidation state gets the –ic ending.
19
o So in the example above, Iron 3+ is the iron ion with the higher oxidation state (the
other ion is iron (II) ) so it would get the –ic ending. Notice that on your table of
common ions, Iron (III) also has the name ferric. So Fe(CN)3 would be named ferric
cyanide.
(See Naming Ionic Compounds Assign)
Writing Compounds from the names:
Just work backwards (and Use Criss Cross):
Ex. Magnesium Chlorite
Ex. Lead (II) Iodide
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Name:
Lab Partner:
Lab Three
Teacher:
Date:
Formula Writing and Chemical Names
Introduction:
A chemical formula is a combination symbols and numerical subscripts that represents the
compostition of a compound. The symbols indicate which elements are present and the numerical
subscripts indicate the relative proportion of each element in the compound. These proportions can
be predicted using the oxidation numbers (charges) of the elements. When atoms acquire a charge
they are called ions.
It is important that all scientists use the same system for writing chemical formulas. This helps
to ensure clear and consistent transmission of information. Therefore, the following rules should be
used for writing chemical formulas:
1. In a neutral compounds the sum of the charges of the elements (ions) must equal zero. One
positive(+) charge will neutralize one negative(-) charge.
2. Elements (ions) with a positive charges are written first
3. When the relative proportion of the polyatomic ion in a compound is greater than one, the
symbol for that ion must be enclosed in parenthesis and followed by a numerical subscript
indicating its relative proportion, as in the ternary compound Aluminum Sulfate whose
formula would be Al2(SO4)3.
Purpose:
Students will observe precipitate formation and write chemical formulas and chemical names for the
precipitate.
Equipment:
Materials:
Pipettes
Solutions of
Glass Plate
NO3-, CO32-, PO43-, Zn2+,
SO42-, Cu2+, Mg2+, K
Procedure:
1. Read through the entire lab procedure, put on your safety glasses.
2. Insert the work page under your glass plate and place on top of work table
3. Make you hypothesis before each reaction occurs.
4. Make your observations of each solution before any reaction occurs.
5. Combine two drops of cation solution with two drops of anion solution in the appropriate
grid square. Be careful not to let the dropper touch the drops of the other solutions.
21
6. Observe the reaction (if any) and record you observations on the corresponding square of
the data table. Also record your more detailed observations in the spaces below.
7. Repeat steps above until you have combined all sixteen possible reactions.
8. Make observations of reaction’s using the corresponding number (in the chart).
Data Analysis:
Observations Before:
NO3-:
CO32-:
PO43-:
SO42-:
Zn2+:
Cu2+:
Mg2+:
K+:
22
Hypothesis:
NO3-
CO32-
PO43-
SO42-
Zn2+
1
2
3
4
Cu2+
5
6
7
8
Mg2+
9
10
11
12
K+
13
14
15
16
Data: Descriptive observations
1.
2.
3.
4.
23
5.
6.
7.
8.
9.
10.
11.
24
12.
13.
14.
15.
16.
Analysis:
1. What is a chemical formula?
2. What information does a subscript in a chemical formula provide?
3. What is a formula unit?
4. When do you need to use a parenthesis in writing a chemical formula?
5. When do you need to use a roman numeral in the name of a compound?
6.
Write out the chemical formula and chemical name for each solution that underwent a chemical
change.
Conclusion:
State which ions experience chemical reactions when combined.
25
Work Sheet
Cl-
S2-
CO32-
OH-
Ag+
1
2
3
4
Co2+
5
6
7
8
Fe3+
9
10
11
12
Cu2+
13
14
15
16
26
Naming Covalent Compounds:
Binary compounds: These are compounds that contain only 2 elements. When naming binary
compounds we use prefixes to indicate the amount of each element present.
1
2
3
4
5
6
7
8
9
10
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
We rarely use mono, because normally we just assume that there is one if no prefix is present. Mono is
only used for oxygen.
Ex.
CO2
P4O10
SO3
N2O4
(See Naming Covalent Compounds Assign)
Write Chemical Formula from Name:
Just write what the formula says:
Dinitrogen trioxide
Diphosphorus pentoxide
(See Writing Formulas From Names Assign)
Naming Acids:
Naming acids that are binary:

These would be covalent molecules that contain a hydrogen and another molecule (Usually a
halogen or sulpfur or phosphorus)
H__
Hydro______ic Acid
o Ex. HI
Hydroiodic acid
o Ex. Hydrochloric acid:
HCl
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Naming Acids that are polyatomic:

These are covalent molecules containing hydrogen and a polyatomic ion.

First we need to point out a pattern that you may or may not have already recognized:
o There are many sets of polyatomic ions that are similar to each other where the only
difference is the number of oxygen atoms.

The ions with more oxygens get an “ate” ending and the ion with less oxygens
gets an “ite” ending.

Ex. Phosphate PO43- and Phosphite PO33Nitrate: NO3- Nitrite: NO2The Oxygen containing polyatomic ions are named as follows:
Least Oxygens: hypo_____ite
______ite
______ate
Most Oxygens: per_______ate
o These are the polyatomic ions that are involved in acids

When naming acids containing polyatomic ions we name them first by writing the name of the
ion.
o We change the ending. If the ion ends in -ate we put an –“ic acid” on it. If the ion ends in
ite we put an –“ous acid” ending on it.
o Ex. H3PO3 Phosphite is the polyatomic ion in this acid, so the name is phosphorous acid.
o Ex. HClO3: ClO3 is chlorate. Since it has an ate ending we name it Chloric acid.
o Ex. Iodic acid. This has an ic ending so the ion must have an –ate ending. Iodate (IO3-).
Then we add however many hydrogens we need to make the molecule neutral. So this
molecule needs one hydrogen

HIO3 (We write the hydrogen first, then the polyatomic ion)
o Ex. Chlorous acid. This has an ous ending so we are looking for an ion with an ite (with
chlorine and oxygen). Chlorite (ClO2-). We add however many hydrogens we need.

HClO2
Note: On exams and assignments where naming is required, it is important that you are able to
identify acids and name them appropriately. For example you name Hydrochloric acid hydrochloric
acid, not hydrogen chloride.
(See Naming Acids Assignment)
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Nomenclature for Organic Compounds
Types of Hydrocarbons:
IUPAC Naming System
Organic compounds are those which contain carbon (excluding oxides and carbonates).
Things to know for naming:

Alkanes are organic compounds that contain only single bonds,

Alkenes are carbon compounds that contain one or more double bonds

Alkynes are organic compounds that contain one of more triple bonds.
All Alkanes follow the CnH2n+2 rule.
All Alkenes containing one double bond between carbons follow the CnH2n rule.
These rules will help you to determine whether a molecule is an alkane, or an alkene.
Ex. CH4 is an Alkane because if n=1, then the molecule would be C1H2(1)+2, Which is CH4.
Is C2H6 and alkane or alkene?
Is C2H2 an alkane or an alkene?
Is C3H6 an alkane or alkene?
Naming
Prefixes
When naming organic compounds we must first find the longest chain of carbons in the molecule.
Each number of carbons in the chains corresponds to a different prefix.
Number of
Carbons
Prefix
Number of
Carbons
Prefix
1
Meth
6
Hex
2
Eth
7
Hept
3
Prop
8
Oct
4
But
9
Non
5
Pent
10
Dec
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Suffix
If the molecule contains all single bonds the suffix is “ane”. If the molecule contains a double bond the
suffix is “ene”. If the molecule contains a triple bond the suffix is “yne”.
Ex 1
Alkenes and alkynes are named using a number to indicate where the double or triple bond is located.
Always number the carbons so that you get the lowest possible number.
Examples:
Ex. 2
Ex. 3
1-butene or but-1-ene
Ex 4.
1-propene
Ex. 4
1-butene
ethyne
If the Compound contains 2 or more double bonds you would need to use di, tri etc.
Ex. 1,2 butadiene or buta-1,2-diene
Alkyl groups
If you remove a hydrogen atom from one of these carbon chains and add a branch, you get an alkyl
group. These are also named based on the number of carbons but the suffix is “yl”
For example methyl –CH3. These alkyl groups can then bond with another compound in place of a
hydrogen.
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To name these we first identify the longest carbon chain. Next number the carbons. To write the name
we first write the number, then the name of the alkyl group, then the name of the longest carbon chain.
# - alkyl group longest carbon chain
Ex 5.
When naming compounds with more than one alkyl group, you must name them alphabetically. You
also number the carbons so that the first group in the name has the lowest number.
If you have two of the same alkyl group you name them as follows:
#, # - dialkyl group longest carbon chain
If you had three of the same alkyl groups you would use 3 numbers to indicate where they are located
and the prefix tri on the name of the alkyl group.
Ex 6.
Ex 7.
Ex 8. 2-methylpropane
Ex 9. 3, 3- diethyl hexane
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Line Forms:
= Propane
= 3-methyl pentane
= Propene
Draw the molecules from examples 1-9 in line form
(See Organic Compounds Assignments)
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Solids- Independent Study
http://www.chm.davidson.edu/vce/crystals/networksolids.html
Crystalline Solids
Crystalline solids fall into one of four categories.
Type of
Solid
Interaction
Properties
Examples
Ionic
Ionic
High Melting Point, Brittle, Hard, Often
dissolve in water.
NaCl, MgO
Molecular
Hydrogen Bonding,
Dipole-Dipole,
London dispersion
Low Melting Point, Nonconducting, can be
crystal
H2, CO2
Metallic
Metallic Bonding
Variable Hardness and Melting Point
(depending upon strength of metallic
bonding), Conducting, usually not crystals
Fe, Mg
Covalent Bonding
High Melting Point, Hard, Nonconducting,
tend not to dissolve in water.
Some forms of
C,
SiO2
Network
All four categories involve packing discrete molecules or atoms into a lattice or repeating array,
though network solids are a special case. The categories are distinguished by the nature of the
interactions holding the discrete molecules or atoms together.
In ionic and molecular solids, there are no chemical bonds between the molecules, atoms, or ions. The
solid consists of discrete chemical species held together by intermolecular forces that are
electrostatic or coulombic in nature. This behavior is most obvious for an ionic solid such as NaCl,
where the positively charged Na+ ions are attracted to the negatively charged Cl- ions. Even in the
absence of ions, however, electrostatic forces are operational. For polar molecules such as CH2Cl2, the
positively charged region of one molecular is attracted to the negatively charged region of another
molecule (dipole-dipole interactions). For a nonpolar molecule such as CO2, which has no
permanent dipole moment, the random motion of electrons gives rise to temporary polarity (a
temporary dipole moment). Electrostatic attractions between two temporarily polarized molecules
are called London Dispersion Forces.
Hydrogen bonding is a term describing an attractive interaction between a hydrogen atom from a
molecule or a molecular fragment X–H in which X is more electronegative than H, and an atom or a
group of atoms in the same or a different molecule, in which there is evidence of bond formation. Dots
are employed to indicate the presence of a hydrogen bond: X–H···Y. The attractive interaction in a
hydrogen bond typically has a strong electrostatic contribution, but dispersion forces and weak
covalent bonding are also present.
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In metallic solids and network solids, however, chemical bonds hold the individual chemical subunits
together. The crystal is essential a single, macroscopic molecule with continuous chemical bonding
throughout the entire structure.
In metallic solids, the valence electrons are no longer exclusively associated with a single atom.
Instead these electrons exist in molecular orbitals that are delocalized over many atoms, producing an
electronic band structure. The metallic crystal essentially consists of a set of metal cations in a sea of
electrons. This type of chemical bonding is called metallic bonding.
Network Solids
In network solids, conventional chemical bonds hold the chemical subunits together. The bonding
between chemical subunits, however, is identical to that within the subunits, resulting in a continuous
network of chemical bonds. Two common examples of network solids are diamond (a form of pure
carbon) and quartz (silicon dioxide). In quartz one cannot detect discrete SiO2 molecules. Instead the
solid is an extended three-dimensional network of ...-Si-O-Si-O-... bonding.
(See Types of Solids Assign)
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Name:__________________________________________
Covalent and Ionic Bonding Assignment
Classify the following compounds as ionic (metal + non-metal), covalent (non-metal + non-metal) or
both (compound containing a polyatomic ion).
1. CaCl2 _________________________________
2. CO2
_________________________________
3. H2O
_________________________________
4. BaSO4 _________________________________
5. K2O
_________________________________
6. NaF
_________________________________
7. Na2CO3_________________________________
8. CH4
_________________________________
9. SO3
_________________________________
10. LiBr
_________________________________
11. MgO
_________________________________
12. NH4Cl _________________________________
13. HCl
_________________________________
14. KI
_________________________________
15. NaOH _________________________________
16. NO2
_________________________________
17. AlPO4 _________________________________
18. FeCl3 _________________________________
19. P2O5
_________________________________
20. N2O3 _________________________________
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Name: _______________________________________
Determining Number of Atoms
Determine the number of atoms in each of the following compounds:
1.
KCl: ___________
8.
MgCl2: ___________
15. Hg2Cl2: ___________
2.
NaCl: ___________
9.
CH3COOH: ___________
16. (NH4)3PO4: ___________
3.
CaCl2: ___________
10. CuC2H3O2: ___________
17. Mg(C2H3O2)2: ___________
4.
KNO3: ___________
11. Ba(OH)2: ___________
18. Mg(NO2)2: ___________
5.
H2SO4: ___________
12. NH4Br: ___________
19. As2(SO4)5: ___________
6.
CaCO3: ___________
13. Ca3(PO4)2: ___________
20. Zn3(PO4)2: ___________
7.
C2H6: ___________
14. Al(OH)3: ___________
Criss Cross Method
Determine the chemical formula for the compounds created by the combination of the ions below.
Cl-
O2-
SO32-
PO43-
NO2-
OH-
CH3COO-
Na+
Mg2+
Co3+
NH4+
Fe3+
As5+
H+
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Name:_____________________________
Ionic Bonding Lewis Dot Assignment
Draw the lewis dot structure for the ionic compounds created from the elements below. Be sure to
draw the lewis dot in 2 steps, to show how the electrons are involved in bonding.
1. K + F
2. Mg + I
3. Be + S
4. Na + O
5. Al + Br
37
Name:_________________________________
Determining Charges Assignment:
1. What is the charge of copper in the following molecule?
CuBr2
2. What is the charge of iron in the following molecule?
FeBrO3
3. What is the charge of cobalt in the following molecule?
CoCrO4
4. What is the charge of copper in the following molecule?
Cu3(PO4)2
5. What is the charge of nickel in the following molecule?
Ni3(PO4)2
6. What is the charge of Silver in the following molecule?
Ag2HPO4
7. What is the charge of Manganese in the following molecule?
MnSO4
8. What is the charge of lead in the following molecule?
Pb(CrO4)2
9. What is the charge of chromium in the following molecule?
Cr2(C2O4)3
10. What is the charge of Arsenic in the following molecule?
As3(BO3)5
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Name:_____________________________________
Covalent Bonding Lewis Dot Assignment
Draw the lewis dot structure for the covalent molecules created from the elements below. Be sure to
draw the lewis dot in 2 steps, to show how the electrons are involved in bonding.
1. H + H (H2)
2. F + F (F2)
3. O + O (O2)
4. N + N (N2)
5. C + O (CO2)
6. H + O (H2O)
39
Name: ____________________________
VSEPR Assign
Determine the electron pair geometry and the molecular geometry for the following, then draw the 3D
structure of the molecule:
Follow Octet:
Break the Octet:
CH4
XeH2
NF3
ArCl4
PH3
NH5
SBr2
FCl3
OBr4
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Name:______________________________
Ionic, Polar Covalent or Non-polar Covalent Assignment
Determine whether the following bonds are ionic, non-polar covalent or polar covalent.
1. N2: The Nitrogen Nitrogen bond
2. HF: The hydrogen fluorine bond
3. LiBr: The lithium Bromine bond
4. AuCl3: The gold Chlorine bond
5. H2O: The hydrogen oxygen bond
6. CH3OH: The carbon hydrogen bond
7. CH3OH: The Carbon Oxygen bond
8. SO3: The Sulphur Oxygen bond
9. KI: The Potassium Iodine bond
10. Fe2O3: The iron, Oxygen bond.
41
Name:___________________________________
Intermolecular Forces Assignment
Determine whether the following molecules will have polar bonds, H-bonds, Van der waals forces or
London Dispersion forces. Be sure to include all forces (i.e. there may be more than one that applies)
1. NH3
2. I2
3. CH4
4. HF
5. O2
6. HBr
7. Br2
8. BF3
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Name:________________________________
Naming Ionic Compounds Assignment
Name the following ionic compounds using Roman Numerals where necessary:
1. BaCl2 ___________________________________
2. Ag2O ___________________________________
3. CuBr ___________________________________
4. FeO ___________________________________
5. MgS ___________________________________
6. Al2O3 ___________________________________
7. K2S ___________________________________
8. CrCl2 ___________________________________
9. Ba3P2 ___________________________________
10. BeS ___________________________________
11. MnO ___________________________________
12. FeSO4 ___________________________________
13. FeCl3 ___________________________________
14. Zn3(PO4)2 ___________________________________
15. NH4NO3 ___________________________________
16. Al(OH)3 ___________________________________
17. CuC2H3O2 ___________________________________
18. PbSO3 ___________________________________
19. NaClO3 ___________________________________
20. CaC2O4 ___________________________________
21. Fe2O3 ___________________________________
22. NaHSO4 ___________________________________
23. Hg2Cl2 ___________________________________
24. Mg(NO2)2 ___________________________________
25. CuSO4 ___________________________________
26. NaHCO3 ___________________________________
27. NiBr3 ___________________________________
28. Be(NO3)2 ___________________________________
29. AuCl3 ___________________________________
30. KMnO4 ___________________________________
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Name:______________________________
Naming Covalent Compounds Assignment
Name the following compounds using the prefix method:
1. CO ___________________________________
2. CO2 ___________________________________
3. SO2 ___________________________________
4. NO2 ___________________________________
5. N2O ___________________________________
6. SO3 ___________________________________
7. CCl4 ___________________________________
8. NO ___________________________________
9. N2O5 ___________________________________
10. P2O5 ___________________________________
11. N2O4 ___________________________________
12. CS2 ___________________________________
13. OF2 ___________________________________
14. PCl3 ___________________________________
15. PBr5 ___________________________________
16. N2O3 ___________________________________
17. PCl5 ___________________________________
18. NH3 ___________________________________
19. SCl6 ___________________________________
20. SiO2 ___________________________________
44
Name:_______________________________
Writing Formulas from Names Assignment
Write the chemical formula for the compounds listed below:
1. Sodium chloride _____________________________
2. Carbon tetrachloride ________________________
3. Magnesium bromide ________________________
4. Aluminum iodide ____________________________
5. Hydrogen hydroxide _________________________
6. Dihydrogen monoxide _______________________
7. Iron (II) fluoride _____________________________
8. Carbon dioxide ______________________________
9. Sodium carbonate ___________________________
10. Ammonium sulfide __________________________
11. Magnesium sulfate ___________________________
12. Dinitrogen pentoxide ________________________
13. Phosphorous trichloride ____________________
14. Copper (I) carbonate ________________________
15. Potassium hydrogen carbonate _______________
16. Sulfur trioxide _______________________________
17. Ammonium phosphate _____________________
18. Iron (II) oxide _______________________________
19. Iron (III) oxide ______________________________
20. Carbon monoxide ___________________________
21. Calcium chloride ____________________________
22. Potassium nitrate ___________________________
23. Magnesium hydroxide _____________________
24. Copper (II) sulfate __________________________
25. Lead (IV) chromate _________________________
26. Diphosphorous pentoxide _________________
27. Potassium permagnate _____________________
28. Sodium hydrogen carbonate ________________
29. Zinc nitrate ___________________________________
30. Aluminum sulfite ______________________________
45
Name:_________________________________
Naming Acids Assignment
Name the following acids:
1. HNO3 ___________________________________
2. HCl ___________________________________
3. H2SO4 ___________________________________
4. H2SO3 ___________________________________
5. HC2H3O2 ___________________________________
6. HBr ___________________________________
7. HNO3 ___________________________________
8. H3PO4 ___________________________________
9. H2S ___________________________________
10. H2CO3 ___________________________________
Write the formulas of the following acids:
1. Sulfuric acid ___________________________________
2. Nitric acid ___________________________________
3. Hydrochloric acid ___________________________________
4. Acetic acid ___________________________________
5. Hydrofluoric acid ___________________________________
6. Phosphorous acid ___________________________________
7. Carbonic acid ___________________________________
8. Nitrous acid ___________________________________
9. Phosphoric acid ___________________________________
10. Hydrosulfuric acid ___________________________________
46
Name:___________________________________
Naming Organic Compounds Assignment
47
48
Name:___________________________________
Structure of Organic Compounds Assignment
Ethane
Ethyne
Propane
3, 3-dimethyl pentane
2-butene
2, 3-dimethyl pentane
Methane
1-butyne
49
Name:________________________
Types of Solids Assignment
Determine if the following compounds are metallic solids, ionic solids, network atomic solids, molecular
solids, or amorphous solids based on their properties. These are all actual chemical compounds.
1)
This material forms crumbly crystals and has a melting point of 16.60 Celsius. It has a low
density in solid form.
molecular_______________________________ (acetic acid)
2)
This material forms very hard colorless crystals. It does not dissolve in water and burns at high
temperatures.
network_______________________________ (diamond, C-C bond)
3)
This material forms colorless crystals that have a melting point of 6610 C. It is hard, brittle, and
dissolves well in water.
Ionic solid_______________________________ (sodium iodide)
4)
This material forms silver crystals that do not dissolve in water and have a melting point of
14140 C. This material is very hard and is not a good conductor of electricity.
Network solid_______________________________ (silicon)
5)
This material is hard and melts at a temperature of 16100 C. It dissolves only with difficulty in
very reactive acids and doesn’t conduct electricity when molten. It forms colorless crystals.
Network solid_______________________________ (quartz)
6)
This material is soft and doesn’t form crystals. It has a melting point of 6600 C. It doesn’t
dissolve in water. It is used as a structural material in the construction of airplanes and
rockets.
Metallic_______________________________ (aluminum)
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