Module 3 Lab Experiment-Chemistry

Molecular Modeling and
Lewis Structures

Version 42-0080-00-02

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in
which to complete the exercise.

Experiment Summary:

In this experiment, you will draw Lewis structures
for a series of molecules and then create the VSEPR
model for the molecule using the modeling kit.

EXPERIMENT

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Define allotropes, valence electrons, and lone pairs.

● Describe the duet rule and octet rule.

● Define and create Lewis structures.

● Describe the valence shell electron pair repulsion (VSEPR) model.

● Draw Lewis structures for molecules.

● Create VSEPR models of molecules with molecular modeling kits.

● Identify the number of valence electrons of elements using the periodic table.

● Diagram resonance structures.

● Classify the VSEPR model of a molecule as: linear, trigonal planar, tetrahedral, trigonal
bipyramidal, or octahedral.

● Assemble molecules with molecular modeling kits, preparing single bonds, double bonds, and
resonance structures.

Time Allocation: 3.5 hours

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Experiment Molecular Modeling and Lewis Structures

Materials
Student Supplied Materials

Quantity Item Description
1 Camera, digital or smartphone
1 Pen
1 Sheet of white paper

HOL Supplied Materials

Quantity Item Description
1 Modeling kit (Molecular Modeling and Lewis Structures):

6 – Single bonds
4 – Double bonds
18 – Lone pairs
5 – White (1 hole)
1 – Pink (2 holes)
1 – Gray (3 holes)
2 – Red (4 holes)
2 – Black (4 holes)
6 – Green (4 holes)
1 – Blue (4 holes)
4 – Yellow (4 holes)
1 – White (5 holes)
1 – Yellow (6 holes)

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software, such as Microsoft Word® or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

Experiment Molecular Modeling and Lewis Structures

Background
Structure and bonding

The structure and molecular bonding of molecules is an important factor in the course of chemical
reactions. Both diamond and graphite are variants of the carbon atom, and while both substances
are exclusively composed of carbon, it is the difference in the structure and molecular bonding
of the carbon atoms that result in the two extraordinarily different minerals. Different structural
modifications of an element are referred to as allotropes. See Figure 1.

Figure 1. Allotropes of the element carbon. A. Crystalline structure of diamond. B. Crystalline
structure of graphite. © Michael Ströck

Many properties and characteristics are involved in chemical bonding and molecular structure
and function including bond strength, polarity, and atomic orbitals. The focus of this lab is the
localized electron model. This model assumes that a molecule is bonded through the sharing of
valence electron pairs. The localized electron model also highlights valence electron arrangement,
Lewis structures, and molecular shape.

Valence electrons and Lewis structures

Valence electrons are the electrons of an atom located in the outermost shell of an atom. The
number of valence electrons that an atom (element) contains can be found in the periodic table.
Elements in the same group (vertical column) of the periodic table contain the same number of
valence electrons. See Figure 2. For example, all elements in group 7 (VIIA), including fluorine,
chlorine, bromine, and iodide, have 7 valence electrons. All of the elements in group 6 (VIA), such
as oxygen and sulfur, have 6 valence electrons.

Experiment Molecular Modeling and Lewis Structures

Figure 2. Periodic Table of Elements. The number of valence electrons of an atom is determined
by the group number, shown directly above each group in roman numerals. Click to Download

Printable Version.

For example, all elements in group 6 (VIA); oxygen, sulfur, chlorine, bromine, iodide, and astatine,
contain 6 valence electrons. Likewise, all elements in group 4 (IVA); carbon, silicon, germanium,
tin, and lead, contain 4 valence electrons. See Figure 3.

Figure 3. Valence electrons examples.

Molecules are surrounded by an electron cloud: the electrons belong to the entire molecule,
rather than the individual atoms. It is useful, however, to model atoms and electrons in an
organized manner to better understand the structure of a molecule.

Experiment Molecular Modeling and Lewis Structures

A Lewis structure shows how valence electrons are arranged among atoms in a molecule. In
arranging valence electrons, the duet and octet rules are very important. The duet rule applies to
molecules containing hydrogen, as hydrogen is most stable when sharing two valence electrons.
The octet rule is based upon the observation that atoms (other than hydrogen) are most stable
when surrounded by eight valence electrons. These 8 valence electrons can either be shared
(bonds) or not shared (lone pairs).

Consider the bonding of hydrogen and fluorine. Hydrogen (Group IA) has 1 valence electron and
fluorine (Group VIIA) has 7 valence electrons. See Figure 4. The atoms form the molecule hydrogen
fluoride (HF), which has a total of 8 valence electrons in its electron cloud. The hydrogen follows
the duet rule, and the fluorine follows the octet rule. Two electrons are shared between them. Six
electrons surrounding fluorine are not shared, and are considered lone pairs. There are a total of
three lone pairs around fluorine, which are shown as three pairs of dots.

Figure 4. Bonding of hydrogen fluoride (HF). Hydrogen has two valence electrons and obeys
the duet rule; fluorine has eight valence electrons and obeys the octet rule.

The Lewis structure of the molecule can be drawn so that a single dash represents the shared
electrons, as shown in Figure 5.

Figure 5. A. Hydrogen and fluorine share one pair of electrons (shown in the circle); there are
also three lone pairs. B. The shared pair may be represented as a single dash that signifies a

bond, while the lone pairs are still drawn as dots.

Experiment Molecular Modeling and Lewis Structures

Creating Lewis Structures

Lewis structures are created in just a few simple steps. Follow along with the procedures listed
below and practice creating the Lewis structure for carbon tetrachloride.

Example 1: Carbon tetrachloride (CCl4) will be used:

Step 1) Calculate the total number of valence electrons in the molecule.

Carbon is located in group 4 of the periodic table; thus, carbon has 4 valence electrons. Chlorine,
located in group 7, has 7 valence electrons and there are 4 chlorine atoms total. The total number
of valence electrons in the molecule is 32.

Step 2) Arrange atoms and create single bonds.

The first atom listed in the molecular formula is often the central atom in the Lewis structure.
(A more precise description is that the least electronegative atom is usually the central atom.) A
wrong choice in the central atom will usually result in the inability to create the Lewis structure.

A single bond is composed of two valence electrons and is noted as either (● ●) or (—). In this
example, 8 electrons are used to create one single bond between each of the C-Cl atoms.

The C atom is placed in the center and surrounded by the Cl atoms. Single lines are drawn between
the atoms representing shared pairs of electrons, which may be thought of as single bonds.

When actors lose an academy
award to another actor they often say,
“It was an honor just to be nominated.”

While not an actor, Gilbert Newton Lewis, an
American chemist for whom the Lewis structure

model was named, can certainly relate to the honor
of a nomination. While Dr. Lewis was nominated for a
Nobel Prize 35 times, he never won. Although Lewis
was denied winning the Nobel Prize multiple times,
this did not reduce the impact he made on science;

rather he has the distinct honor of having been
nominated for the award more than any

other scientist (thus far!).

Experiment Molecular Modeling and Lewis Structures

Step 3) Calculate the number of remaining valence electrons; then distribute the electrons with
the goal of satisfying the duet rule or octet rule.

Since each single bond (represented by a line) contains 2 shared electrons, 8 electrons have
already been added to the Lewis structure. As determined in step 1, the molecule contains 32
valence electrons total. Thus, there are 24 more electrons (32-8=24) to add to the Lewis structure.

C already satisfies the octet rule.

Placing the remaining 24 valence electrons around the chlorine atoms satisfies the octet rule for
each Cl. The lone pairs are represented by pairs of dots.

Step 4) Check your work: ensure the duet and octet rules are satisfied and count the total
number of valence electrons.

Review the molecule to ensure that all atoms are surrounded by 8 valence electrons, satisfying
the octet rule. Ensure that hydrogen atoms (when present) satisfy the duet rule.

Count the number of electrons represented in your diagram. Ensure that the number of elec-
trons you count matches the number calculated in step 1.

Example 2: Oxygen gas (O2)

Step 1) Calculate the total number of valence electrons in the molecule.

Oxygen atoms have 6 valence electrons, and there are a total of 2 atoms. There are 12 valence
electrons total.

Experiment Molecular Modeling and Lewis Structures

Step 2) Arrange atoms and create single bonds.

The two oxygen atoms should be placed in a line and linked with a single bond that represents
shared electrons.

Step 3) Calculate the number of remaining valence electrons; then distribute the electrons with
the goal of satisfying the duet rule or octet rule.

Two electrons (held in 1 single bond) have already been added to the Lewis structure.
This molecule has 12 valence electrons total, as determined in step 1. Ten more electrons
must be added, and this can be accomplished through a trial-and-error process.

Creating a double bond between the atoms and then adding lone pairs satisfies the octet rule for
all atoms and also generates the correct number of valence electrons. (Double bonds represent
4 shared electrons.)

Consider the outcome if the Lewis structure was drawn with a single bond, and the octet rule was
fulfilled. The total number of electrons would be 14. In step 1, we determined that the molecule
has 12 valence electrons, so we know that the Lewis structure below is incorrect, even though
the octet rule is fulfilled.

Consider the outcome if the Lewis structure was drawn with the correct number of total valence
electrons and a single bond. The Lewis structure below is incorrect because, although 12 electrons
are present, the octet rule is not fulfilled.

Note: Multiple bonds are used ONLY when there are not enough lone pairs present for each atom
to fulfill the octet rule. In the incorrect Lewis structure above, we know a multiple bond is needed
because the correct total valence electrons are present but the octet rule is not fulfilled.

Experiment Molecular Modeling and Lewis Structures

Step 4) Check your work: ensure the duet and octet rules are satisfied and count the total
number of valence electrons.

A review of the correct molecule shows that all atoms are surrounded by 8 valence electrons,
satisfying the octet rule. The number of electrons drawn in the Lewis structure (in the lone pairs
and the double bond) totals 12, which matches the calculation in step 1.

Example 3: Cyanide (CN-)

Step 1) Calculate the total number of valence electrons in the molecule.

As indicated in the molecular formula, this molecule has a negative charge. Therefore, an additional
electron must be accounted for. Carbon contributes 4 valence electrons, nitrogen contributes 5
valence electrons, and the negative charge in the chemical formula indicates 1 additional valence
electron.

There are 10 valence electrons total in CN-.

Note: In this example, a “–“charge was shown. If a “+” charge were present, then valence electron(s)
would be subtracted rather than added.

Step 2) Arrange the atoms and create single bonds.

The two atoms are placed in a line and linked with a single bond.

Step 3) Calculate the number of remaining valence electrons; then distribute the electrons with
the goal of satisfying the duet rule or octet rule.

Two electrons were added to the Lewis structure through the addition of the single bond; 8 more
electrons were required. Through trial-and-error, it can be determined that the two atoms share
6 electrons in a triple bond. The C and N atoms each have one set of lone pairs, satisfying the
octet rule.

Experiment Molecular Modeling and Lewis Structures

Consider the outcome if the Lewis structure was drawn with a single bond, and the octet rule was
fulfilled. The total number of electrons would be 14. In step 1, we determined that the molecule
has 10 valence electrons, so we know that the Lewis structure below is incorrect, even though
the octet rule is fulfilled.

Consider the outcome if the Lewis structure was drawn with the correct number of total valence
electrons and a single bond. The Lewis structures below are both incorrect because, although 10
electrons are present in both structures, the octet rule is not fulfilled.

Consider the outcome if the Lewis structure was drawn with a double bond and the octet rule
was fulfilled. The Lewis structure below is incorrect because, yet again, the number of valence
electrons is incorrect. The total number of electrons in the structure below is 12, and it was
determined that the molecule has 10 valence electrons.

The correct Lewis structure has a triple bond; carbon has one lone pair, and nitrogen has one lone
pair. The octet rule is fulfilled, and the valence electrons total 10, as calculated in step 1.

However, the Lewis structure is not complete. The charge of the molecule is denoted by placing
the structure in brackets and writing the charge in the upper right-hand corner.

Experiment Molecular Modeling and Lewis Structures

Resonance Structures

Resonance occurs when more than one valid Lewis structure exists. In this scenario, the various
molecular structures may be referred to resonance structures.

Consider the Lewis structure for the polyatomic ion CO3
-2. Following the steps for drawing Lewis

structures, we find that CO3
-2 has a total of 24 electrons (4+(3)6+2=24). Through trial-and-error, it

is determined that one double bond, two single bonds, and 8 lone pairs exist.

A Lewis structure for CO3
-2 can be drawn as follows:

In the diagram above, the oxygen atoms are evenly distributed around the central carbon atom.
However, notice that there is more than one option for the placement of the double bond. The
double bond could be placed between the C atom and any one of the O atoms. Thus, the Lewis
structure may be represented in three drawings.

The CO3
-2 molecule has resonance, which occurs when more than one valid Lewis structure

exists. In reality, the electron structure of CO3
-2 is a combination of all three resonance structures.

Resonance is represented by double-headed arrows as follows. Notice that the -2 charge is
denoted for each resonance structure.

Exceptions to the Octet Rule

Oxygen, fluorine, nitrogen, and carbon always obey the octet rule. However, it is often said in jest
that the only rule in science is, “There is an exception to every rule.” Indeed, the octet rule applies
to most atoms in molecules, but there are exceptions. Some elements, such as Boron, tend to have
fewer than 8 valence electrons. Likewise, other elements can have more than 8 valence electrons.
Only the elements located in or below period 3 (row 3) of the periodic table tend to exceed an octet.

Experiment Molecular Modeling and Lewis Structures

Consider the Lewis structure for iodine tetrachloride (ICl4
-). The sum of the valence electrons is

36 (7+(4)7+1=36). Drawing the molecule and placing lone pairs around the Cl atoms results in a
representation of only 32 electrons.

Where do the remaining 4 electrons go? Since iodine is located below period 3 of the periodic
table, the remaining lone pairs may be placed around the central iodine atom. (You may learn
later in your course that this occurs because elements like iodine can exceed the octet by using
their empty valence d orbitals.) In the case of ICl4

-, iodide breaks the octet rule, while the Cl atoms
obey the octet rule.

As shown above, the total number of valence electrons (36) are represented in the Lewis structure.
The placement of the Cl atoms and the lone pairs are adjusted so that the central iodide atom is
evenly surrounded.

Valence shell electron pair repulsion (VSEPR)

While Lewis structure models describe atoms in the two-dimensional sense, the valence shell
electron pair repulsion (VSEPR) model describes the three-dimensional arrangement of the
molecule.

The VSEPR model arranges atoms in a manner that minimizes electron pair repulsion, maintaining
the most stable form of the molecule. Consider the two-dimensional Lewis structure for methane
(CH4) shown in Figure 6. The three-dimensional VSEPR model has bond angles of 109.5 degrees,
placing the hydrogen atoms as far from one another as possible

Figure 6. Lewis structure and VSEPR model of methane (CH4)

Experiment Molecular Modeling and Lewis Structures

The molecular geometry (or molecular shape) of CH4 is classified as a tetrahedral arrangement. In
fact, anytime one atom is surrounded by four other atoms (and no lone pairs), the geometry will
be tetrahedral and the molecule will exist in its most stable form.

A selection of molecular geometries is shown in Figure 7. Central atoms are represented by grey
spheres, and surrounding atoms are represented by red spheres. The black bars represent bonds,
which could be single or double bonds. Lone pairs are represented by green clouds.

Figure 7. Selection of molecular geometries showing central atoms (grey Spheres), surrounding
atoms (red spheres), single/double bonds (black bars), and electron pairs (green clouds).

Experiment Molecular Modeling and Lewis Structures

Experiment Molecular Modeling and Lewis Structures

Lone pairs are not considered part of the molecular geometry, but the presence of lone pairs
around the central atom affects the arrangement of the other atoms and dictates the shape of
the entire molecule. For example, compare the trigonal planar and trigonal pyramidal structures
in Figure 7. Both structures have one central atom (grey) and three surrounding atoms (red).
Notice that the trigonal planar arrangement has no lone pairs and it is flat. The trigonal pyramidal
arrangement does have a lone pair and the structure is not flat- it is “lifted” like a pyramid. In
your course, you may also learn about electron-pair geometry, which does take into account the
shapes of the electron configuration; however, the focus of this lab experience will be molecular
geometries around the central atom.

Review all of the geometries in Figure 7. Try to predict which geometries will allow the central
atom to have more than eight valence electrons.

Note: Please note that there are additional shapes that are not shown in Figure 7. For example,
a central atom surrounded by one lone pair and four atoms is an “irregular tetrahedron” (not
shown). Variations on the geometries listed also exist. A bent molecule can include a central atom
surrounded by two atoms and only one lone pair as opposed to the two lone pairs shown in Figure 7.
The figure includes all of the geometries you need to be successful in the Experimentation. For more
geometries, consult a textbook or a reliable internet source.

Molecular Modeling Kits

Molecular modeling kits are used to visualize the three-dimensional structure of a molecule.
Kits typically contain balls that represent atoms, sticks that represent bonds, and paddles that
represent lone pairs. See Figure 8.

Figure 8. Molecular modeling kit.

Experiment Molecular Modeling and Lewis Structures

The Lewis structure and molecular model for chloromethane (CH3Cl) are shown in Figure 9. The
Lewis structure includes a carbon atom surrounded by three hydrogen atoms and one chlorine
atom. The chlorine atom has 3 lone pairs. Creating the molecule with a modeling kit shows the
three-dimensional placement of the atoms and lone pairs. The model includes one carbon atom
(black) surrounded by three hydrogen atoms (small, white), and one chlorine atom (dark green)
surrounded by three lone pairs (pink paddles).

Figure 9. Methyl chloride (CH3Cl): Lewis structure and molecular model.

To identify the molecular geometry, the central atom must be identified. Then the geometry
can be determined using Figure 7 as a guide. For example, the central atom in CH3Cl is carbon.
Using Figure 7 as a guide, it may be determined that the geometry around the carbon atom
is tetrahedral. Remember, the lone pairs around the chlorine atom are not included when
determining the molecular geometry.

Experiment Molecular Modeling and Lewis Structures

Geometry around Multiple Central Atoms

When a molecule has more than one interior atom, more than one geometry may be used to
describe the molecule. Consider the Lewis structure and molecular model of ethylene (C2H4) in
Figure 10. Note that the two carbons are connected by a double bond. Each carbon is considered
a central atom.

Figure 10. Ethylene C2H4: Lewis structure and molecular model.

Focus on only one interior atom at a time. Using Figure 7 as a guide, the geometry around only
the leftmost carbon atom of C2H4 is trigonal planar. Taking only the rightmost carbon into account,
the geometry is again trigonal planar. For C2H4 it may be concluded that the geometry is “trigonal
planar around each carbon atom.”

Experiment Molecular Modeling and Lewis Structures

Exercise 1: Lewis Structures and Molecular Modeling
In this exercise, the student will draw Lewis structures for a series of molecules and then create
the VSEPR model for the molecule using the modeling kit.

Procedure

1. Gather the modeling kit, a white sheet of paper, a pen, and the digital camera.

2. Review the key in Table 1. Check that the contents of your modeling kit match the key.

Table 1. Element, color, and holes.

Element/bond/etc. Color Holes # Included
Hydrogen White 1 5
Magnesium Pink 2 1
Aluminum Gray 3 1
Carbon Black 4 2
Oxygen Red 4 2
Chlorine and Florine Green 4 6
Lead and Nitrogen Blue 4 1
Sulfur and Iodine Yellow 4 4
Tin White 5 1
Sulfur and Iodine Yellow 6 1
Single Bond Gray Short and rigid 6
Double Bond Gray Long and flexible 4
Lone Pair (2 valence electrons) Pink Paddle shaped 18

Part 1: Practice Describing Molecular Structures

3. Practice describing the molecular structure of CHO2
– (See Lab Report Assistant for Answers).

Number of Valence Electrons:

● How many valence electrons does CHO2
– have? Use the periodic table to calculate the

total.

Lewis Structures:

● What is the Lewis structure for CHO2
– ? Practice drawing the Lewis structure on a separate

sheet of paper.

Hint: CHO2
– has resonance structures, and there are two forms of the drawn molecule. Draw both

structures.

VSEPR Models:

● Use your molecular modeling kit to create a CHO2
– molecule. Although the molecule has

two Lewis structures, you only need to build one molecule.

Experiment Molecular Modeling and Lewis Structures

Note: Consult Table 1 to determine which pieces represent the C, H, and O atoms. To create a double
bond, use TWO of the long, flexible gray connectors. To create a single bond, use one of the short,
inflexible connectors. Pink paddles represent lone pairs. The completed molecule should have no
“open” or unfilled holes.

Atoms:

● What is the central atom? If there is more than one interior atom, list each.

● How many bonds and electron pairs surround the central atom(s)?

Geometry:

● Identify the molecular geometry of the molecule. Refer to Figure 11 as needed.

Figure 11. Molecular geometries.

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Experiment Molecular Modeling and Lewis Structures

4. Practice the questions above until you feel comfortable describing the CHO2
– molecule.

Note: If you would like additional practice, describe the molecules in the background section.

Part 2: Describing Molecular Structures

5. Using Table 1 as a guide, label and organize the atoms of the molecular modeling kit. See
Figure 12.

Figure 12. Organized molecular modeling kit.

Note: You might not use all of the pieces included in the modeling kit during this exercise.

6. For each molecule listed below, you will: 1- calculate the number of valence electrons, 2- draw
the Lewis structure(s), 3- upload an image of the VSEPR model, 4- list the number of bonds
and lone pairs surrounding the central atom(s), and 5- identify the structure geometry; or
identify multiple geometries if there is more than one central atom.

Experiment Molecular Modeling and Lewis Structures

7. Use the periodic table in Figure 2 of the Background to calculate the number of valence
electrons in CCl4. Click here to download a printable copy of the periodic table. Record in Data
Table 1 of your Lab Report Assistant.

8. Draw the Lewis structure for CCl4 and insert into Data Table 1. Ensure that the features of
the molecule are large enough and clear enough for your instructor to grade. Refer to the
background for step-be-step instructions on drawing Lewis structures.

Note: To draw the Lewis structure, either use a computer drawing program, such as Power Point, or
other computerized drawing tool and cut and paste into Data Table 1, or draw by hand, scan, resize,
and then insert into Data Table 1, or draw by hand and take a photograph, resize and then insert
into Data Table 1. Refer to the appendix entitled, “How to Label an Image” for guidance Refer to
the appendix entitled, “Resizing an Image” for guidance. If resonance structures exist for any of the
molecules, make sure to draw all of the structures.

9. Use the molecular modeling kit to build CCl4. Ensure that the appropriate atoms are used, as
described in Table 1. The completed molecule should have no “open” or unfilled holes.

10. Place the molecular model of CCl4 on a white sheet of paper with your name and the date
written on it, take a photograph of the model, as shown in Figure 13. Resize and insert the
image into Data Table 1.

Figure 13. NCH3CH3CH3: molecular model with student name and date.

11. Determine the number of bonds and lone pairs that surround the central atom. Record the
data in Data Table 2 of your Lab Report Assistant.

Note: If there is more than one central atom, identify each atom and list the number of electron
clouds that surround each.

Experiment Molecular Modeling and Lewis Structures

12. Use Figure 11 to determine the name of the geometric structure for ClC4. Structure names
include: linear, bent, trigonal planar, trigonal pyramidal, tetrahedral, square pyramidal,
and octahedral. Record the name Data Table 2.

13. Repeat steps 5-12 for the remaining molecules.

Hints:

● Some molecules have more than one central atom: multiple geometries will need to be
described.

● Some molecules have a charge: draw brackets and indicate the charge in the Lewis
structure.

● Some molecules have resonance: draw all of the forms for the structure.

● Most, but not all, of the molecules fulfill the duet and octet rules.

14. When you are finished uploading photos and data into your Lab Report Assistant, save and
zip your file to send to your …

Naming Ionic and
Molecular Compounds
Hands-On Labs, Inc.
Version 42-0315-00-01

Review the safety materials and wear goggles when
working with chemicals. Read the entire exercise
before you begin. Take time to organize the materials
you will need and set aside a safe work space in which
to complete the exercise.

Experiment Summary:

You will apply the rules for naming ionic and molecular
compounds to write the names of compounds when
given the chemical formula. You will also write the
formula for ionic and molecular compounds when
given the name.

EXPERIMENT

Learning Objectives
Upon completion of this laboratory, you will be able to:

● Describe how the periodic table arranges elements by their chemical properties.

● Discuss the IUPAC naming system.

● Define molecular compound, ionic compound, polyatomic ion, oxidation state, and diatomic
element.

● Identify the prefixes and suffixes used to name polyatomic ions and list the name, formula,
and charge of common polyatomic ions.

● List the rules for converting formulas to names, and names to formulas for ionic and molecular
compounds.

● Explain the difference between binary and oxoacids, and identify the rules for naming each.

● Generate a colored periodic table to distinguish between the groups of elements, and create
a list of common polyatomic ions and strong acids to aid in naming chemical compounds.

● Write the names for ionic compounds, molecular compounds, polyatomic ions, and acids by
interpreting their formulas.

● Write the chemical formula for ionic compounds, molecular compounds, polyatomic ions, and
acids by interpreting their compound names.

Time Allocation: 3 hours

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Experiment Naming Ionic and Molecular Compounds

Materials
Student Supplied Materials

Quantity Item Description
1 Box of colored pencils or highlighters
1 Computer printer
1 Digital camera or smartphone
1 Package of note cards
1 Pen or pencil
1 Sheet of paper

Note: To fully and accurately complete all lab exercises, you will need access to:

1. A computer to upload digital camera images.

2. Basic photo editing software such as Microsoft® Word or PowerPoint®, to add labels, leader
lines, or text to digital photos.

3. Subject-specific textbook or appropriate reference resources from lecture content or other
suggested resources.

Note: The packaging and/or materials in this LabPaq kit may differ slightly from that which is listed
above. For an exact listing of materials, refer to the Contents List included in your LabPaq kit.

Experiment Naming Ionic and Molecular Compounds

Background
The Periodic Table

The periodic table is a reference for the arrangement of chemical elements. See Figure 1. The
periodic table not only organizes the elements by their atomic number and electron configurations,
it also organizes the elements by their chemical properties. The periodic table is the most important
tool to have on hand when studying chemistry. A periodic table is located on the inside cover of
almost every chemistry textbook.

Figure 1. Periodic Table of Elements. Click to Download Printable Version.

Communication skills are important in any field, and the language of chemistry has a vocabulary
of its own. In the medical field, it is essential to communicate clearly and effectively. For example,
it is important to be correct and unambiguous when transferring responsibility for a patient to
another person, or asking the doctor or dentist for required medication. Reports may be written
for the doctor or for communicating with a patient about their condition, and learning how to
write chemical names and formulas is the first step in pharmacology.

For example, the ionic compound potassium nitrite (KNO2) is used to treat chest pain, whereas the
compound potassium nitrate (KNO3) is used to treat asthma and is also found in toothpastes for
sensitive teeth. A pharmacologist or doctor would need to ensure the proper indication of these
compounds to properly and effectively treat a patient’s symptoms.

Experiment Naming Ionic and Molecular Compounds

The IUPAC Naming System

The International Union of Pure and Applied Chemistry (IUPAC) naming system provides a systematic
method of naming compounds around the globe. The IUPAC naming system was designed with the
fundamental principle that “each different compound should have a different name.” Each IUPAC
name for organic compounds consists of several parts: 3D relationship, numbered substituents,
number of carbons in the longest chain, and the ending suffix that describes functional groups.
Figure 2 shows the molecular structure for vitamin C, also called ascorbic acid. The formal, IUPAC
name for the chemical is (5R)-5-[(1S)-1,2-dihydroxyethyl]-3,4-dihydroxyfuran-2(5H)-one.

Figure 2. The unique IUPAC name for ascorbic acid (vitamin C) includes the numbered
constituents, 3D relationship of atoms, the number of carbons, and functional groups.

©Macrovector

The IUPAC naming system assures safety and consistency when using chemicals. It would be difficult
to replicate experiments if scientists used different names for the same compound. Safety would
also be a concern if there was no consistent system for naming because of the various hazards
associated with mixing chemicals.

Experiment Naming Ionic and Molecular Compounds

Naming Molecular Compounds

Molecular compounds consist of nonmetallic elements that share electrons through covalent
bonds. The type of molecular compounds we will focus on naming in this lesson are binary
molecular compounds. Binary molecular compounds consist of two nonmetal elements. For
example, a water molecule (H2O) is a binary molecular compound consisting of the nonmetals
hydrogen and oxygen.

To name a binary molecular compound, the first element is given its elemental name, and the
second element is given its root (i.e. carb-, hydr-, ox-, fluor-) with the suffix “-ide.” For example:

● HF = hydrogen fluoride

Greek prefixes are used for molecular compounds to account for the number of each element in
the compound. See Table 1 for the Greek prefixes that represent the numbers 1 – 10.

Table 1. Greek prefixes.
Number of Atoms Prefix

1 Mono-
2 Di-
3 Tri-
4 Tetra-
5 Penta-
6 Hexa-
7 Hepta-
8 Octa-
9 Nona-

10 Deca-

For example:

● CO2 = carbon dioxide

● N2O3 = dinitrogen trioxide

The prefix “mono-” is never used for the first element, and only used for the second element
if ambiguity exists in the naming. A few examples for using the prefix “mono-” on the second
element are:

● carbon monoxide

● dinitrogen monoxide

● nitrogen monoxide

Note: If the final vowel in a prefix is “a” or “o” it is dropped before the vowel in a stem name, for ease
of pronunciation.

Experiment Naming Ionic and Molecular Compounds

Converting a Molecular Compound Formula to a Name

Example 1: Writing a Nonmetal + Nonmetal Name: P2S5
1. Read the formula and look at the subscripts.

P2S5

Note that P and S are nonmetals, which can be determined by referencing a periodic table like the
one in Figure 1.

2. Write the name of the first element with the correct Greek prefix.

P2 = diphosphorus

3. Write the root name of the second element with the suffix “-ide.”

S = sulfur = sulfide

4. Write the correct Greek prefix of the second element.

S5 = pentasulfide

5. Write the name of the molecular compound.

P2S5 = diphosphorus pentasulfide

Example 2: Writing a Nonmetal + Nonmetal Name: CO

1. Read the formula and look at the subscripts.

CO

2. Write the name of the first element with the correct Greek prefix. If the 1st element has the
prefix “mono”, it is dropped.

C = carbon (NOT monocarbon)

3. Write the root name of the second element with the suffix “-ide.”

O = oxygen = oxide

4. Write the correct Greek prefix of the second element. If the final vowel in a prefix is “a” or “o”
it is dropped before the vowel in a stem name, for ease of pronunciation.

O = monoxide (NOT monooxide)

5. Write the name of the molecular compound.

CO = carbon monoxide

Note: There is no charge indicated in the above formulas which indicates that they are molecular
compounds and not ionic compounds.

Experiment Naming Ionic and Molecular Compounds

Converting a Name to a Molecular Compound Formula

Example 3: Writing a Nonmetal + Nonmetal Formula: Carbon Tetrafluoride

1. Read the name of the compound.
carbon tetrafluoride

2. Write the first chemical symbol based on the first name written in the compound.
carbon = C

3. Include the number of atoms based on the prefix (if any) included in the first name.
1 carbon atom = C

Note: Since the word “carbon” in carbon tetrafluoride has no prefix, it can be assumed that there is
only one carbon atom in the molecule.

4. Write the second chemical symbol based on the second name written in the compound.
fluoride = fluorine = F

5. Include the number of atoms based on the prefix (if any) included in the second name.
tetrafluoride = 4 fluorine atoms = F4

6. Write the formula of the molecular compound.
carbon tetrafluoride = CF4

Note: You will not be asked to “balance” the formula of molecular compounds because there are no
ions to balance. Ionic charges and balancing equations will be introduced later.

Example 4: Writing a Nonmetal + Nonmetal Formula: Diboron Trioxide

1. Read the name of the compound.
diboron trioxide

2. Write the first chemical symbol based on the first name written in the compound.
boron = B

3. Include the number of atoms based on the prefix (if any) included in the first name.
diboron = 2 boron atoms = B2

4. Write the second chemical symbol based on the second name written in the compound.
oxide = oxygen = O

5. Include the number of atoms based on the prefix (if any) included in the second name.

trioxide = 3 oxygen atoms = O3

6. Write the formula of the molecular compound.
diboron trioxide = B2O3

Experiment Naming Ionic and Molecular Compounds

Naming Ionic Compounds

Ionic compounds are chemical compounds containing both a cation (positive ion) and an anion
(negative ion) held together by electrostatic forces, also known as ionic bonds. Ionic compounds
are different from molecular compounds in that they usually contain a metal or an ammonium ion
(NH4

+) and molecular compounds are composed of nonmetals. A binary ionic compound contains
two elements, one metal and one nonmetal.

To name a binary ionic compound, the cation elemental name is listed first, followed by the root
of the anion ending in “-ide.” For example:

● NaCl = sodium chloride

Greek prefixes are not used in naming the number of atoms of each element for ionic compounds.
For example:

● Li3N = lithium nitride, NOT trilithium nitride.

Ionic compounds are written as neutral compounds, meaning the overall charge of the compound
must equal zero. In order to determine the overall charge, the oxidation state of the metal and
nonmetal must be known in order to ensure the correct number of atoms of each element are
present in the compound. The oxidation state represents the number of electrons that an atom can
gain, lose, or share when bonded with an atom of another element. The oxidation state for each
element can be found on the periodic table in Figure 1. We will discuss how to use the oxidation
state to calculate the number of atoms of a compound in further detail later.

Converting an Ionic Compound Formula to a Name

1. Use the periodic table to determine if a metal is present in the compound.

Note: Be aware that many elements have similar names and symbols.

a. If a metal is present, the compound is likely an ionic compound.

2. Determine if the metal ion (cation) has a “fixed charge.” A fixed charge means that there is
only one possible oxidation state, as shown in Figure 3. Cations with a fixed charge include.

a. All elements in Group IA and Group IIA

b. Al3+

c. Transition metals with a fixed charge: Ag1+, Zn2+, and Cd2+

Experiment Naming Ionic and Molecular Compounds

Figure 3. The fixed charge oxidation states, or charges, of the element potassium (K) in Group IA
and the element beryllium (Be) in Group IIA.

3. Proceed to step 6 if the metal does not have a fixed charge, otherwise continue to step 4.

4. Name the cation first.

a. Record the entire name of the cation.

5. Name the anion second.

a. Combine the root name of the nonmetal anion with the suffix “-ide.” The conversion
from a formula to a name is complete once the anion has been named.

6. Determine the variable charges for the metal, as shown in Figure 4. Variable charge means
that there is more than one possible oxidation state. Elements with a variable charge include:

a. All transition metals in Groups IIIA through IIB, EXCEPT: Ag1+, Zn2+, and Cd2+

b. All basic metals EXCEPT: Al3+

Figure 4. The variable charge oxidation states, or charges, that iron (Fe) and gold (Au) can carry.

Experiment Naming Ionic and Molecular Compounds

7. Name the cation first.

a. Include the entire name of the cation.

b. Balance the ionic charges of the cation and anion. Identify the number of ions required
to generate a neutral compound.

c. List the Roman numeral in parentheses based on the ionic charge. Do not put a space
between the cation name and the parentheses. For example, iron(III) chloride.

8. Name the anion second.

a. Include the root name of the anion and the suffix “-ide.”

Example 5: Writing a Metal + Nonmetal Name: Li2O

1. Read the formula and look at the subscripts.

Li2O

Note that Li is a metal and O is a nonmetal, which can be determined by referencing a periodic table
like the one in Figure 1.

2. Determine if the metal ion has a fixed charge or variable charge.

Li = Group IA = fixed charge = 1+

3. Record the entire name of the cation.

Li = lithium

4. Write the root name of the nonmetal ion with the suffix “-ide.”

O = oxygen = oxide

5. Write the name of the ionic compound.

lithium oxide

Example 6: Writing a Metal + Nonmetal Name: Fe2O3
1. Read the formula and look at the subscripts.

Fe2O3
2. Determine if the metal ion has a fixed charge or variable charge.

Fe = Group VIIIB = variable charge = 2+ or 3+

3. Record the entire name of the cation.

Fe = iron

www.HOLscience.com 11 ©Hands-On Labs, Inc.

Experiment Naming Ionic and Molecular Compounds

Experiment Naming Ionic and Molecular Compounds

4. Determine the Roman numeral by balancing the charges to create a neutral compound.

a. There are 2 atoms of iron and 3 atoms of oxygen.

b. We know oxygen always carries a 2- charge so all 3 atoms of oxygen have a 2- charge.

c. We need to determine if each iron atom has a 2+ or a 3+ charge by a simple algebraic
calculation where the charge for each iron atom equals x and the equation is set to 0:

Tip: To balance Fe2O3, first consider the oxygen atoms. Oxygen always has a charge of 2-, therefore
O3 has a total charge of 6-. This means that Fe2 must have a total charge of 6+. Each F atom will
have a charge of 3+.

d. After solving for x, we find that each iron atom carries a 3+ charge. Thus, the Roman
numeral used in the compound name is “III.”

iron(III)

5. Write the root name of the nonmetal ion with the suffix “-ide.”

O = oxygen = oxide

6. Write the name of the ionic compound.

iron(III) oxide

Converting a Name to an Ionic Compound Formula

When writing formulas for ionic compounds, the sum of the positive and negative charges MUST
equal zero to obtain a neutral compound.

Example 7: Writing a Metal + Nonmetal Formula: Calcium Fluoride

1. Read the name of the compound.

calcium fluoride

Experiment Naming Ionic and Molecular Compounds

2. Determine if the metal ion has a fixed charge or variable charge.

calcium = Group IIA = fixed charge = 2+

Note: If the name does not include a Roman numeral, then the metal ion has a fixed charge.

3. Write the first chemical symbol based on the cation in the compound, including the charge.

calcium = Ca2+

4. Write the second chemical symbol based on the anion in the compound, including the charge.

fluoride = fluorine = F-

5. Balance the charges to determine the correct number of atoms of each element and write the
formula for a neutral compound.

Tip: To balance CaF2, first consider the Ca has a fixed charge of 2+ and F carries a 1- charge. For the
compound to have a net charge of 0, there must be two F atoms and one Ca atom.

6. Write the formula of the molecular compound.

calcium fluoride = CaF2

Example 8: Writing a Metal + Nonmetal Formula: Gold(III) Chloride

1. Read the name of the compound.

gold(III) chloride

2. Determine if the metal ion has a fixed charge or variable charge.

gold = Group IB = variable charge = 1+ or 3+

3. Write the first chemical symbol based on the cation in the compound, including the charge.
The name tells us it is gold(III) meaning it carries a 3+ charge:

gold(III) = Au3+

4. Write the second chemical symbol based on the anion in the compound, including the charge.

chloride = chlorine = Cl-

Experiment Naming Ionic and Molecular Compounds

5. Balance the charges to determine the correct number of atoms of each element to write the
formula for a neutral compound.

6. Write the formula of the molecular compound.

gold(III) chloride = AuCl3

Polyatomic Ions

Polyatomic ions are a group of two or more covalently bonded atoms that function as a single
ion. For example, oxygen carries a 2- charge and hydrogen carries a 1+ charge. When combined
they form a covalent bond, producing a single anion with a 1- charge called hydroxide (OH-). The
1- charge of the hydroxide ion is a product of the 2- charge of the O with the 1+ charge of the
H (-2 + 1 = -1). Polyatomic ions are the fundamental unit in the majority of ionic compounds;
therefore, knowing polyatomic ion names, formulas, and charges is important. Common ions and
their charges are listed in Table 2.

Table 2. Common polyatomic ions.

Name Formula Charge

Ammonium NH4
+ 1+

Hydroxide OH- 1-

Cyanide CN- 1-

Nitrite** NO2
– 1-

Nitrate** NO3
– 1-

Sulfite** SO3
2- 2-

Sulfate** SO4
2- 2-

Hydrogen sulfite** HSO3
– 1-

Hydrogen sulfate** HSO4
– 1-

Experiment Naming Ionic and Molecular Compounds

Name Formula Charge

Carbonate CO3
2- 2-

Hydrogen carbonate HCO3
– 1-

Phosphate PO4
3- 3-

Hydrogen phosphate HPO4
2- 2-

Dihydrogen phosphate H2PO4
– 1-

Hypochlorite ClO- 1-

Chlorite** ClO2
– 1-

Chlorate** ClO3
– 1-

Perchlorate ClO4
– 1-

Peroxide O2
2- 2-

Chromate*** CrO4
2- 2-

Dichromate*** Cr2O7
2- 2-

Permanganate*** MnO4
– 1-

**Note the very subtle differences in names and subscripts due to the oxidation states.

*** Note that permanganate, chromate, and dichromate each have a metal and a
nonmetal.

Note: There are additional tables of polyatomic ions available online and in textbooks. This table
includes only the most common ions encountered in general chemistry classes.

Here are a few helpful hints for naming polyatomic ions:

a. Suffixes: the name of the ion usually ends in “-ite” or “-ate.” A low oxidation state will
have an ion ending in “-ite,” versus a higher oxidation state that ends in “-ate.” The
oxidation state is dependent upon a calculation of the charges of the polyatomic ion.
Figure 5 shows the calculations of the oxidation state for the sulfite ion (SO3

2-) and the
sulfate ion (SO4

2-). A simple algebraic calculation is done to determine the oxidation
state for each polyatomic ion by solving for “x” since we know that oxygen (O) carries
a 2- charge:

Experiment Naming Ionic and Molecular Compounds

Figure 5. Algebraic calculations for the oxidation state of the polyatomic ions sulfite and sulfate.

b. Prefixes: “hypo-” indicates the very lowest oxidation state and “per-” indicates the very
highest oxidation state. Figure 6 shows the calculations of the oxidation state for the
hypochlorite ion (ClO-) and the perchlorate ion (ClO4

-). A simple algebraic calculation
is done to determine the oxidation state for each polyatomic ion by solving for “x”
since we know that oxygen (O) carries a 2- charge:

Figure 6. Algebraic calculations for the oxidation state of the polyatomic ions hypochlorite and
perchlorate.

Experiment Naming Ionic and Molecular Compounds

c. Few polyatomic ions have positive charges. These ions have names ending in “-onium.”
For example, ammonium (NH4

+) and hydronium (H3O
+).

d. There are a few exceptions to these rules. The following polyatomic ions were once
thought to be monatomic ions so they end in “-ide:” hydroxide (OH-), cyanide (CN-),
and peroxide (O2

2-).

Naming Polyatomic Ion Compounds

Example 9: Writing a Metal + Polyatomic Ion Name: KNO2
1. Read the formula and look at the subscripts.

KNO2
2. Determine if the metal ion has a fixed charge or variable charge.

K = Group IA = fixed charge = 1+

3. Record the entire name of the cation.

K = potassium

4. Record the entire name of the anion.

NO2 = nitrite

5. Write the name of the ionic compound.

potassium nitrite

Example 10: Writing a Metal + Polyatomic Ion Formula: Gold(I) Nitrate

1. Read the name of the compound.

gold(I) nitrate

2. Determine if the metal ion has a fixed charge or variable charge.

gold = Group IB = variable charge = 1+ or 3+

3. Write the first chemical symbol based on the cation in the compound, including the charge.
The name tells us it is gold(I) meaning it carries a 1+ charge.

gold(I) = Au+

4. Write the second chemical symbol based on the anion in the compound, including the charge.

nitrate ion = NO3

Experiment Naming Ionic and Molecular Compounds

5. Balance the charges to determine the correct number of atoms of each element to write the
formula for a neutral compound.

6. Write the formula of the molecular compound.

gold(I) nitrate = AuNO3

Naming Acids

Simple covalent compounds that contain hydrogen often dissolve in water to produce acids.
For example, HF in its gaseous state (g) is hydrogen fluoride, but HF in aqueous solution (aq) is
hydrofluoric acid. This lesson focuses on binary acids and oxoacids. Binary acids are acids in which
hydrogen bonds with a second nonmetallic element. An oxoacid is an acid containing oxygen,
hydrogen, and a third element. An oxoacid contains at least one hydrogen atom bound to the
oxygen. These acids dissociate in water by breaking the OH bond to form a H+ ion and an anion.

There are a unique set of rules for naming both binary and oxoacids. The systematic names for
binary acids include the prefix “hydro-” (indicating the water the acid is dissolved in) and the root
of the second element’s name (the non-metal). The suffix of the anion changes from “-ide” to “-ic,”
followed by the word “acid.” For example:

water + H+ + F- = HF(aq) = hydrofluoric acid

Note: Acids containing sulfur use the full name “sulfur” instead of the root of the name. For example,
H2S(aq) is called hydrosulfuric acid NOT hydrosulfic acid.

Oxoacids are named based on the nonmetal from which they are derived. The prefix “hydro-” is
not used, and the suffix “-ate” is changed to “-ic,” and the suffix “-ite” is changed to “-ous.” For
example, HNO3 (contains the nitrate ion) is named nitric acid. The prefixes used for lowest and
highest oxidation states are also used in the naming of oxoacids. For example, HClO (contains the
hypochlorite ion) is named hypochlorous acid. See Table 3 for examples of formulas and names
for strong acids that may be encountered.

Note: Since acids are covalent compounds dissolved in water, the formula must indicate the physical
state the compound is in to distinguish it from covalent compounds NOT dissolved in water. For
example, HCl(g) is hydrogen chloride whereas HCl(aq) is hydrochloric acid.

Experiment Naming Ionic and Molecular Compounds

Table 3. The formulas and systematic names for a selection of common strong acids.

Acid Formula Non-metal Present in Acid Systematic Name
HCl(aq) Chloride Hydrochloric acid
HBr(aq) Bromide Hydrobromic acid
HI(aq) Iodide Hydroiodic acid

HNO3(aq) Nitrate Nitric acid
H2SO4(aq) Sulfate Sulfuric acid

Diatomic Elements

Diatomic elements do not have full valence electron shells and cannot exist as a single atom.
For example, the reactivity of hydrogen causes lone atoms to combine into diatomic (two atom)
molecules forming hydrogen gas (H2(g)) or liquid hydrogen (H2(l)). Element names ending in “-gen”
or “-ine” are diatomic. The names and formulas for the seven diatomic elements can be found in
Table 4.

Note: These elements are diatomic ONLY when they are the only element present, NOT when they
are chemically bonded to other elements.

Table 4. The formulas and names of the seven diatomic elements.

Seven Diatomic Elements

Hydrogen H2
Nitrogen N2
Oxygen O2
Fluorine F2
Chlorine Cl2

Bromine Br2

Iodine I2

Review and Tips

In the Exercises, you will study the periodic table and familiarize yourself with polyatomic ions,
common acids, and the diatomic elements. Then, you will practice the discipline of following a
set of rules to write the names and formulas of ionic and molecular compounds. It is important
to follow all of the rules and not to skip steps. Do not go too fast, or guess at names, charges, or
formulas. If you guess incorrectly you may learn or memorize the rules incorrectly. Review the
flowchart in Figure 7 for help differentiating between molecular and ionic compounds.

Experiment Naming Ionic and Molecular Compounds

Figure 7. Flowchart for naming ionic and molecular compounds.

The following tips will help you in the coming Exercises:

1. If there is no metal present in the compound, the compound is likely a molecular compound.

2. One way to distinguish between binary ionic compounds and binary molecular compounds is
to look for a prefix. Molecular compounds use a Greek prefix to indicate the number of each
atom in the compound. Refer to Table 1.

3. When naming ionic compounds (a metal and a nonmetal or polyatomic ion), if the metal ion is
in Group IA or IIA, or one of the following elements: Ag, Zn, Cd, or Al, then the cation is named
first. The anion is named second: monoatomic ions have the suffix “-ide,” and polyatomic ions
have no suffix.

Experiment Naming Ionic and Molecular Compounds

4. If the metal is in Groups IIIA through IIB (EXCEPT for Ag, Zn, Cd, or Al) then the ionic compound
is named as in step 3, except a Roman numeral is added in parentheses after the metal name
to indicate the charge of the metal ion.

5. When writing the formula for ionic compounds, write the symbol for the cation first, followed
by the anion.

6. EVERY time a compound contains a metal, balance the charges of the compound when writing
the formula.

7. Print or write down the steps on naming ionic and molecular compounds in the Background,
and use them for every example.

8. An aqueous physical state distinguishes binary acids from simple molecular compounds.

9. Double check your work. After you have written the name of a chemical compound, cover the
name and try to write the corresponding formula, and vice versa.

10. Use the note cards that will be created in Exercise 1 for every compound in Exercise 2.

Practicing is the best way to learn how to name chemical compounds and write the corresponding
formulas.

Chemists can have a sense
of humor at times when naming

molecules. Some of the more classroom
“appropriate” include, Draculin which is a

large glycoprotein found in vampire bat saliva.
Penguinone, named from its similarity in 2D

structure to a penguin. The keto acid of morolic
acid, isolated from the mora tree, is named moronic
acid with derivatives called moronates, as in “which
moron-ate all of the pie?” Traumatic acid is a plant
hormone that causes injured cells to divide and help
repair trauma to the plant. Uranium has resulted in

the creation of numerous silly names such as the
uranium oxide anions known as urinates, uranium

nitrate which is also known as uranyl nitrate,
and U4+ known as the uranous ion.

Name:__________________

College Chemistry I – Chem. 1806

Lab Report 3 assignment

Naming Chemical Compounds

1. Name the following elements, binary ionic compounds, and polyatomic ions using the rules that have been discussed. (5 points)

Data Table 1
Item to nameName
1SClick or tap here to enter text.
2SbClick or tap here to enter text.
3N2Click or tap here to enter text.
4PO43-Click or tap here to enter text.
5AuClick or tap here to enter text.
6RbClick or tap here to enter text.
7LiClClick or tap here to enter text.
8AlBr3Click or tap here to enter text.
9KMnO4Click or tap here to enter text.
10Cu(OH)2Click or tap here to enter text.
11FeSO4Click or tap here to enter text.
12NH4ClClick or tap here to enter text.
13ZnCO3Click or tap here to enter text.
14SnF2Click or tap here to enter text.
15MgSO4Click or tap here to enter text.
16MnO2Click or tap here to enter text.
17Ca3(PO4)2Click or tap here to enter text.
18NaOHClick or tap here to enter text.
19HClClick or tap here to enter text.
20Ba(HSO3)2Click or tap here to enter text.
21HNO3Click or tap here to enter text.
22FeSO3Click or tap here to enter text.
23Al2S3Click or tap here to enter text.
24KNO3Click or tap here to enter text.

2. Write the formula of the following compounds using the rules that have been discussed. (5 points)

Data Table 2
Compound NameFormula
1Potassium cyanideClick or tap here to enter text.
2Ammonium carbonateClick or tap here to enter text.
3Lithium selenideClick or tap here to enter text.
4Calcium hydrogen carbonateClick or tap here to enter text.
5Potassium carbonateClick or tap here to enter text.
6Strontium hydroxideClick or tap here to enter text.
7Cobalt (III) phosphateClick or tap here to enter text.
8Iron (II) sulfideClick or tap here to enter text.
9Zinc permanganateClick or tap here to enter text.
10Silver nitrateClick or tap here to enter text.
11Cadmium sulfideClick or tap here to enter text.
12Sodium sulfideClick or tap here to enter text.
13Barium iodideClick or tap here to enter text.
14Lead (II) chlorideClick or tap here to enter text.
15Aluminum hydroxideClick or tap here to enter text.
16Barium permanganateClick or tap here to enter text.
17Lithium sulfateClick or tap here to enter text.
18Sulfuric acidClick or tap here to enter text.
19Magnesium chlorideClick or tap here to enter text.
20Potassium hydroxideClick or tap here to enter text.
21Sodium sulfiteClick or tap here to enter text.
22Tin (II) fluorideClick or tap here to enter text.
23Iron (II) chlorideClick or tap here to enter text.
24Copper (II) nitrateClick or tap here to enter text.

3. Write the name of the following compounds using the rules that have been discussed. Nitric oxide will not be accepted for the compound name of NO since it has not been discussed. (5 points)

Data Table 3
FormulaCompound Name
1B2O3Click or tap here to enter text.
2NOClick or tap here to enter text.
3N2O4Click or tap here to enter text.
4S2F10Click or tap here to enter text.
5BrF3Click or tap here to enter text.
6H2OClick or tap here to enter text.
7SiCl4Click or tap here to enter text.
8H2O2Click or tap here to enter text.
9COClick or tap here to enter text.
10N2OClick or tap here to enter text.
11SiO2Click or tap here to enter text.
12CCl4Click or tap here to enter text.
13PCl3Click or tap here to enter text.
14PCl5Click or tap here to enter text.
15SF6Click or tap here to enter text.
16O2F2Click or tap here to enter text.
17P4S3Click or tap here to enter text.
18XeF4Click or tap here to enter text.

4. Write the formulas of the following compounds using the rules that have been discussed. (5 points)

Data Table 4
Compound NameFormula
1Tribromine octoxideClick or tap here to enter text.
2Tetraiodide nonoxideClick or tap here to enter text.
3Dihydrogen monosulfideClick or tap here to enter text.
4Iodine monochlorideClick or tap here to enter text.
5Nitrogen monoxideClick or tap here to enter text.
6Hydrogen monochlorideClick or tap here to enter text.
7Tricarbon dioxideClick or tap here to enter text.
8Bromine trichlorideClick or tap here to enter text.
9Nitrogen trihydrideClick or tap here to enter text.
10Dinitrogen monosulfideClick or tap here to enter text.
11Carbon tetrabromideClick or tap here to enter text.
12Sulfur trioxideClick or tap here to enter text.
13Iodine hexafluorideClick or tap here to enter text.
14Tetraphosphorus decasulfideClick or tap here to enter text.
15Dichlorine pentoxideClick or tap here to enter text.
16Diphosphorus tetraiodideClick or tap here to enter text.
17Diphosphorus pentasulfideClick or tap here to enter text.
18Carbon dioxideClick or tap here to enter text.

Lewis Structure Model

5. Complete the following table. No pictures from the internet are allowed. The VSEPR model must be made using the model kit from your lab kit.(5 points)

Data Table 1
Molecule or Ionic CompoundNumber of Valence ElectronsLewis Structure (drawn by hand)VSEPR Model (insert a picture of your model)
1CCl4Click or tap here to enter text.
2MgCl2Click or tap here to enter text.
3AlCl3Click or tap here to enter text.
4PbI4Click or tap here to enter text.
5CH4Click or tap here to enter text.
6ICl5Click or tap here to enter text.
7CH2OClick or tap here to enter text.
8NF3Click or tap here to enter text.
9H2OClick or tap here to enter text.
10CO2Click or tap here to enter text.
11[NO2]-Click or tap here to enter text.
12[NH4]+Click or tap here to enter text.
13SO2Click or tap here to enter text.
14NH3Click or tap here to enter text.
15H2SClick or tap here to enter text.
16SF6Click or tap here to enter text.
17SbCl5Click or tap here to enter text.

6. Complete the following table. (5 points)

Data Table 2
MoleculeNumber of atoms or sets of lone pairs surrounding the central atomStructure Geometry (Name)
1CCl4Click or tap here to enter text.Click or tap here to enter text.
2MgCl2Click or tap here to enter text.Click or tap here to enter text.
3AlCl3Click or tap here to enter text.Click or tap here to enter text.
4PbI4Click or tap here to enter text.Click or tap here to enter text.
5CH4Click or tap here to enter text.Click or tap here to enter text.
6ICl5Click or tap here to enter text.Click or tap here to enter text.
7CH2OClick or tap here to enter text.Click or tap here to enter text.
8NF3Click or tap here to enter text.Click or tap here to enter text.
9H2OClick or tap here to enter text.Click or tap here to enter text.
10CO2Click or tap here to enter text.Click or tap here to enter text.
11[NO2]-Click or tap here to enter text.Click or tap here to enter text.
12[NH4]+Click or tap here to enter text.Click or tap here to enter text.
13SO2Click or tap here to enter text.Click or tap here to enter text.
14NH3Click or tap here to enter text.Click or tap here to enter text.
15H2SClick or tap here to enter text.Click or tap here to enter text.
16SF6Click or tap here to enter text.Click or tap here to enter text.
17SbCl5Click or tap here to enter text.Click or tap here to enter text.

7. Which, if any, of the molecules in Data Table 1 (Lewis Structure Model) had resonance structures? How many resonance structures did each of the molecules have? (5 points)

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8. Lewis structures are drawn for molecules and not typically for ionic compounds. Explain why Lewis structures are not typically drawn for ionic compounds. (5 points)

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