Physical Science U4
Physical Science U4
Physical Science U4
Table of Contents 1
Essential Questions 4
Review 4
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Check Your Understanding 40
Challenge Yourself 41
Laboratory Activity 60
Performance Task 61
Self Check 63
Key Words 63
Wrap Up 64
Photo Credits 65
References 65
Answer Key 66
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G
RADES 11/12 | PHYSICAL SCIENCE
Unit 4
Structures and Shapes of
Molecules
All of the substances that you encounter every day are built from different
elements that form bonds with other elements to form compounds. Some
substances can conduct electricity, while some cannot. Some are extremely volatile,
while some are tough to melt. Have you ever wondered why these substances
have different properties? Why is it that some substances are reactive and
some are inert?
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Essential Questions
At the end of this unit, you should be able to answer the following questions.
● Which part of the atom is responsible for bonding interactions?
● What are the different types of chemical bonds and how are they formed?
● How can molecular structures be represented?
● What are the different molecular geometries?
● Why do molecules take the shape they do?
Review
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Lesson 4.1: Counting Valence Electrons
Objective
In this lesson, you should be able to:
● determine the number of valence electrons of an atom.
Warm-Up
Materials:
● valence electrons map
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Procedure:
1. Receive one element card from your teacher.
2. In numerical order, have them add one marker (bingo chips/buttons) at a
time to the map.
Guide Questions:
1. What does the marker represent?
2. How many valence electrons are there in your element?
Learn about It
Fig. 1. The periodic table of elements, where rows are labeled based on IUPAC
rules.
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The periodic table can be divided into different partitions according to their
electron configurations. These are representative elements, transition metals,
lanthanides, and actinides.
Silicon (Si), for example, has a condensed electron configuration of 1s2 2s2 2p63s2 3p2.
Since its outermost shell has an n of 3, you can count all electrons contained in
shells with n = 3 (i.e. 3s23p2 ). Thus, Si has four valence electrons. On the same hand,
Beryllium (Be) has a condensed electron configuration of 1s22s2. Its outermost shell
has an n of 2 and contains only two valence electrons.
In general, you can use the periodic table of elements in determining the number of
valence electrons of representative elements because they follow a pattern. The
number of valence electrons of representative elements is equal to their group
numbers provided by the Chemical Abstracts Service (CAS). Elements in group 1A
always have one valence electron, while elements in group 8A have eight.
Table 4.1. G
roup number systems and their respective number of valence
electrons of representative elements
CAS group IUPAC group
Number of valence electrons
number number
1A 1 1
2A 2 2
3A 13 3
4A 14 4
5A 15 5
6A 16 6
7A 17 7
8A 18 8
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The d-block
The transition metals are those in the d-block. Elements in this block have their d
subshells being filled. They are the elements belonging to groups 3 to 12 as shown
in Fig. 1. Their electrons behave differently, such as the electrons present in the
largest d- block for the atom adds to the total number of valence electrons from the
largest shell of s orbital. The valence configuration for transition metals is ns x
(n-1)dy . The total number of valence electrons is equal to the sum of x and y. For
example, titanium has a valence configuration of [Ar] 4s2 3d2. Therefore, it has 4
valence electrons. For transition metals too, the number of valence electrons is
equal to its group number.
However, when transition metals form compounds, they do not use all their
valence electrons. The number of electrons an element can gain or lose when it
forms compounds is called valence (valency). A transition element can have
multiple valencies, depending on the condition of the atom.
Table 4.2. G
roup number systems and their respective number of valence
electrons of transition metals
Number of
CAS group IUPAC group First element
valence Valency
number number of the group
electrons
3B 3 scandium 3 2-3
4B 4 titanium 4 2-4
5B 5 vanadium 5 2-5
6B 6 chromium 6 2-6
7B 7 manganese 7 2-7
8B 8 iron 8 2-3
8B 9 cobalt 9 2-3
8B 10 nickel 10 2-3
1B 11 copper 11 1-2
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2B 12 zinc 12 2
Zinc, belonging to IUPAC group 12, has 12 valence electrons and a valency of two
while copper, belonging to IUPAC group 11, has 11 valence electrons and a valency
of one or two electrons. Mercury (Hg), a transition metal having an atomic number
of 80, belongs to IUPAC group 12, has 12 valence electrons and a valency of two.
Iron, cobalt, and nickel belong to the same group and contain the same number of
valence electrons.
Example 1
How many valence electrons does Ca has?
Solution
Step 1 Identify what is required to find in the problem.
You are asked to determine the number of valence electrons in Ca.
Step 4 Determine the number of valence electrons from its group number.
Since Ca belongs to group 2A, it contains two valence electrons.
L
et us Practice
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How many valence electrons does N have?
Example 2
How many valence electrons does Mn have?
Solution
Step 1 Identify what is required to find in the problem.
You are asked to determine the number of valence electrons in Mn.
Step 4 Determine the number of valence electrons from its group number.
Since Mn belongs to group 7B, it contains seven valence electrons.
L
et us Practice
How many valence electrons does Zn have?
Example 3
How many valence electrons does NO2 have?
Solution
Step 1 Identify what is required to find in the problem.
You are asked to determine the number of valence electrons in N
O2.
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The compound, NO2, is made of 1 nitrogen and 2 oxygens.
L
et us Practice
Count the number of valence electrons of PCl3.
Key Points
● Valence electrons are electrons located in the outermost shell of the atom
of a given element.
● The periodic table can be divided into representative elements, transition
metals, lanthanides, a nd actinides.
● The valence electrons of representative elements are those present in the
electron shell having the highest principal quantum number, n.
● The valence electrons of the main transition and inner transition elements
vary depending on the atom.
● The valence electrons of a compound are simply the sum of the valence
electrons of each constituting atoms or elements.
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Web Links
For further information, you can check the following web links:
Challenge Yourself
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Answer the following questions briefly and clearly.
1. What are valence electrons and what important role do they play?
2. What are the columns in the periodic table called and when do
representative elements fall under the same column?
3. What is the implication of two or more representative elements being under
the same column?
4. How do you find the number of valence electrons of a representative
element?
5. How do you determine the number of valence electrons of a transition
metal?
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Lesson 4.2: Lewis Structures of Molecules
Objective
In this lesson, you should be able to:
● draw the Lewis dot structures of atoms and molecules.
To better understand how reactions occur, the structure of each compound must
be drawn first. The type of bond formed between atoms depends on their nature
and properties. A method for drawing molecular structures have been devised in
order to incorporate the individual properties of the elements into the molecule.
Have you ever wondered why the representation of chemical structures is the
way it is?
Warm-Up
Marble Madness
Before you fully understand the concept behind Lewis structures, you first have to
know how chemical bonds are formed between atoms. In this activity, you will
demonstrate covalent bonding through a game.
Materials:
● small cups
● marbles
Procedure:
1. The class will be divided into 4 groups: A, B, C, and D.
2. Each student will have 1 cup and the following number of
marbles presented on the table.
3. Listen carefully to your teacher as he will shout out the
conditions. (Example: Group A and Group C will share 4
marbles.)
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4. Do the sharing process and group yourselves as necessary to satisfy the
given conditions.
5. Raise your hand when you’re done and have allowed your teacher to check if
you are correct.
Guide Questions:
1. What do the marbles represent?
2. From the activity, how would you describe covalent bonding?
Learn about It
Lewis Structures
Gilbert Newton Lewis was an American physical chemist
who proposed the idea of covalent bonding and the
importance of the electron pairs in bonding. He devised
the use of Lewis electron-dot symbols, simply known as
Lewis symbols.
For example, a neutral nitrogen atom, with the chemical symbol N, has five valence
electrons. These electrons are placed one at a time around the element’s chemical
symbol, eventually occupying its four sides.
Fig. 2. The first four valence electrons must occupy all sides of the chemical symbol
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before pairing up.
Each of the next electrons (i.e. dots) is then paired to one of the first four. This
applies to elements with five or more valence electrons. All sides of the chemical
symbol are equivalent. This means that whichever side having two electrons rather
than one electron is not important.
Octet Rule
When atoms combine to form compounds, they tend to gain, lose or share
electrons to achieve eight electrons in the valence shell. This represents a stable
structure that is commonly found in noble gases except for helium. The pattern is
called the octet rule. For a compound between two nonmetals, atoms share
electrons to have an equivalent of eight electrons in their valence shell.
Covalent bonds
A covalent bond formation can simply be represented with these Lewis symbols as
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well. For instance, two Cl atoms can share an electron with each other to form one
covalent bond. Because of this newly formed bond, each chlorine atom would now
complete its octet. Remember that each Cl atom has seven valence electrons and
would only need one more to complete the needed eight electrons.
The bond in the chlorine molecule is composed of a shared pair of electrons. This
electron pair is formed when two electrons from two separate orbitals of two
atoms pair up, occupying a space between the two atoms. These kinds of structures
are called Lewis structures or L
ewis electron-dot structures.
Fig. 4. Each chlorine atom shares one electron to form the chlorine molecule.
The pair of electrons that forms the bond (bonding electrons) is called a bonding
pair while the pair of electrons that do not take part in bonding (nonbonding
electrons) is called a lone pair. Note that lines could be drawn between atoms that
are bonded instead of using dots.
If you count the number of electrons around one Cl atom, you will see that both of
them contains eight electrons. The Cl atoms in this molecule are stable because
they follow the octet rule. Six of these electrons are nonbonding and two are
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shared in a single bond.
Multiple bonds
Sometimes, atoms connected with single bonds do not meet the eight-rule
requirement. One way to reinforce the octet rule in these atoms is to form multiple
bonds such as double bonds and triple bonds.
In the case of double bonds, each of the double covalent bonds is represented by a
double line. Consider the CO2 molecule as an example. The Lewis structure of CO2
is shown in Fig. 6. A carbon atom contains four electrons in its valence shell. It
would therefore need to share four electrons resulting in four-electron bonding
pairs. On the other hand, an oxygen atom has six electrons in its valence shell, so it
needs to share two electrons. Therefore, two atoms of oxygen can combine with
one atom of carbon to form CO2. This bonding indicates the presence of two double
covalent bonds.
1. Atoms possessing less than eight electrons such as hydrogen, boron, and
beryllium follow reduced electron requirements. Hydrogen and helium
follow the duet rule and needs two electrons, beryllium follows the quartet
rule and needs four electrons, and boron follows the sextet rule and needs
six electrons.
2. Atoms possessing an odd number of electrons do not meet the octet
requirements. Most compounds involving nitrogen (and other Group V
elements) can possess only seven electrons. These compounds are also
known as free radicals. An example is nitrogen monoxide, NO, with a total
valence electron of 11. In this case, nitrogen will have one unpaired electron
and will not follow the octet rule.
3. Atoms possessing more than eight electrons exhibit hypervalence where
the elements go even higher than ten. Due to their empty d orbitals,
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elements in Period 3 or higher such as. P, Br, I, Xe, and Sb can accommodate
electrons exceeding eight. An example of a hypervalent compound is PF5
which uses hybridized sp3d orbitals from empty d orbitals and filled s and p
orbitals. The five valence electrons of phosphorus will be distributed evenly
to these hybridized orbitals which allow the atom to bond five times.
For chlorine, it has the electron configuration 1s2 2s2 2p5 . It belongs to
Group 7A. Thus, it has 7 valence electrons. At its valence shell (n = 2), it
has 7 electrons with an unpaired electron in its orbital, which then can
pair with one more electron from other atoms.
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Step 3 Draw the skeleton structure of the molecule.
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two electrons in the skeleton structure and needs six more electrons.
These remaining electrons, then, will be distributed on each atom.
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule. In
this compound, all atoms follow the octet rule.
Worked Examples
Example 1
Draw the Lewis structure of NH3.
Solution
Step 1 Determine the number of valence electrons in each constituent
element of the compound.
Since N belongs to Group 5A, it contains five valence electrons. Since H
belongs to Group 1A, it contains one valence electron.
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N 5 1 5×1=5
H 1 3 1×3=3
Total number of valence electrons in NH3 8
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule.
There is no need to add multiple bonds because there were enough
electrons to be distributed.
L
et us Practice
Draw the Lewis structure of CHCl3.
Example 2
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Draw the Lewis structure of BF3.
Solution
Step 1 Determine the number of valence electrons in each constituent
element of the compound.
Since B belongs to Group 3A, it contains three valence electrons. Since F
belongs to Group 7A, it contains seven valence electrons.
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Lewis structure of BF3
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule.
There is no need to add multiple bonds because there were enough
electrons to be distributed. Boron also follows sextet, a form of the
reduced octet.
L
et us Practice
Draw the Lewis structure of SF6.
Example 3
Draw the Lewis structure of NO2.
Solution
Step 1 Determine the number of valence electrons in each constituent
element of the compound.
Since N belongs to Group 5A, it contains five valence electrons. Since O
belongs to Group 6A, it contains six valence electrons.
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Skeleton structure of NO2
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule.
If you noticed, the central atom only has 5 electrons. Nitrogen has to have
at least 7. Since there are not enough electrons, putting a double bond
would help.
L
et us Practice
Draw the Lewis structure of CS.
Key Points
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○ Electron-deficient atoms which can only achieve duet and sextet;
○ electron-rich atoms which can exhibit hypervalence; and
○ odd-electron atoms that form r adicals.
● Sometimes, the octet rule can be reinforced by forming multiple bonds.
Web Links
For further information, you can check the following web links:
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7. O2F2 9. SO3
8. SCl6 10. N2O5
Challenge Yourself
Answer the following questions briefly and clearly.
1. Why is hydrogen an exception to the octet rule?
2. Why are elements in Period 3 or higher exceptions to the octet rule?
3. Why must valence electrons be accounted for in drawing Lewis structures?
4. When are multiple bonds necessary in drawing Lewis structures?
5. Why is the least electronegative atom usually the central atom?
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Lesson 4.3: Predicting Shapes through
VSEPR Theory
Objective
In this lesson, you should be able to:
● determine the structures and shapes of molecules using the
VSEPR model.
Molecules take different shapes and forms depending on their nature. Like Lewis
structures, the geometry of the molecules depends largely on the valence electrons
of the atoms comprising it. In fact, given a Lewis structure, one can predict the
shape of the molecule based on the number of electron domains. A different
number of domains yield different geometry. In what ways can you find out
which shapes molecules take?
Warm-Up
Molecular Balloons
Molecules come in various geometric structures depending on the number of
electron domains around their central atom. There are five shapes from which all
the molecular geometries are derived, which you will build in this activity.
Materials:
● balloons
Procedure:
A. Building a Linear Geometry Model
1. Inflate two balloons of the same size.
2. Build the model representing linear geometry by tying these balloons
together. You can use the image below as your reference.
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Linear Geometry
Tetrahedral Geometry
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D. Building a Trigonal Bipyramidal Geometry Model
1. Inflate two balloons of the same size and tie them together.
2. Inflate three balloons of the same size and tie them together.
3. To build the final model representing trigonal bipyramidal geometry, twist
both sets through the center and merge them together. You can use the
image below as your reference.
Octahedral Geometry
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Guide Questions:
1. What does each balloon represent?
2. If a full rotation is 360o, what is the angle between the two balloons in the
linear geometry? How about In the trigonal planar geometry?
3. There are instances when the balloons would just configure themselves to
maximize their space and relieve spatial stress. Explain how this analogy
related to that of orbitals.
Learn about It
Remember that electrons in the valence shell are usually present in pairs with
opposite spins. There could be a single, two, or three pairs of electrons. They can
also be present as nonbonding pairs or even a single unpaired electron.
The ED model assumes that the shape of a molecule can be predicted by arranging
the electrons in a geometry that keeps them separated as far as possible.
Therefore, two pairs of electrons in the atoms of a molecule can be separated by an
angle of 180°, three bonding domains can be separated by an angle of 120°, and
four bonding domains forming a tetrahedron are separated by an angle of 109.5°.
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For instance, apply the ED model to formaldehyde, H2CO. The Lewis structure of the
molecule shows that four pairs of electrons are present in the valence shell of the
central atom. The two pairs of electrons form a double bond, and there are two
single bonds present. Since the four pairs of electrons are located in the three
bonding domains, the ED model predicts that the geometry of H2CO has a trigonal
planar geometry.
These angles predicted by VSEPR resulted from the fact that repulsion between
lone pairs differs from that between bonding pairs. In terms of repulsion strength,
the theory states that
lone pair - lone pair repulsion > lone pair - bonding pair repulsion > bonding pair -
bonding pair repulsion
For instance, there are four domains in the valence shell of the central atom in H2O.
The valence electrons should be distributed towards the corners of a tetrahedron.
However, the main goal of using the ED model is to predict the geometry of the
molecule, rather than knowing the distribution of electrons.
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closer to one another as the lone pairs repel each other stronger. In fact, the bond
angle between atoms in a water molecule is approximately equivalent to 104.5o.
Molecular Geometry
If all the electrons around the central atom are bonded, then the molecular
geometry is the same as the ED geometry. However, when there is at least one lone
pair around the central atom, the molecular geometry is just a derivative of the ED
geometry. Table 4.1 shows the molecular geometry of any molecule with a given
number of bonding and nonbonding domains.
Table 4.3. T
he different molecular shapes for different electron domains.
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The actual geometry of the molecule can be determined by following the steps
below. Consider the molecule boron trifluoride (BF3).
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Boron is an element known to follow a reduced octet. Instead of
having eight electrons, it can only accommodate a maximum of six
electrons. Fluorine and other halogens usually occupy terminal
positions and follow the octet rule. The Lewis structure of BF3 is shown
below.
Step 2 Count the total number of electron domains of the central atom.
The central atom B in the molecule BF3 has three bonds and has no
lone pairs. Hence, the total number of electron domains in the
molecule is 3.
You can use Table 4.3 as your reference. Molecules with three electron
domains are classified under the electron domain geometry of trigonal
planar.
Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
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the same if no lone pairs are initially present in the central atom of the
molecule.
Worked Examples
Example 1
Determine the molecular group geometry and the bond angles between atoms of
PCl5.
Solution
Step 1 Draw the Lewis structure.
Phosphorus is the least electronegative atom and it occupies the central
position. It is also known to exhibit hypervalency. The Lewis structure of
PCl5 is shown below.
Step 2 Count the total number of electron domains of the central atom.
The central atom P has five bonding pairs and has no lone pairs. The total
number of electron domains of the central atom is five.
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Step 3 Identify the electron domain geometry of the molecule.
Using Table 4.3, molecules with five electron domains are classified as
trigonal bipyramidal.
Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Based on Table 4.3, the molecular group geometry of a trigonal
bipyramidal electron domain geometry with no lone pairs is also trigonal
bipyramidal. Hence, the molecular group geometry of PCl5 is trigonal
planar and the bond angles between atoms are 120o and 90o.
L
et us Practice
Determine the molecular group geometry and the bond angles between atoms of
SF6.
Example 2
Determine the molecular group geometry and the bond angles between atoms of
NH3.
Solution
Step 1 Draw the Lewis structure.
Oxygen occupies the central position since hydrogen atoms always occupy
terminal positions. The Lewis structure of NH3 is shown below.
Step 2 Count the total number of electron domains of the central atom.
The central atom N has three bonding pairs and has one lone pair. The
total number of electron domains of the central atom is four.
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Step 3 Identify the electron domain geometry of the molecule.
Using Table 4.3, molecules with four electron domains are classified as
tetrahedral.
Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Based on Table 4.3, the molecular group geometry of a tetrahedral
electron domain geometry with a lone pair is trigonal pyramidal. Hence,
the molecular group geometry of NH3 is trigonal pyramidal and the bond
angles between atoms are less than 120o.
L
et us Practice
Determine the molecular group geometry and the bond angles between atoms of
NI3.
Example 3
Determine the molecular group geometry and the bond angles between atoms of
H2O.
Solution
Step 1 Draw the Lewis structure.
Oxygen occupies the central position since hydrogen atoms always occupy
terminal positions. The Lewis structure of H2O is shown below.
Step 2 Count the total number of electron domains of the central atom.
The central atom O has two bonding pairs and has two lone pairs. The
total number of electron domains of the central atom is four.
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Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Based on Table 4.3, the molecular group geometry of a tetrahedral
electron domain geometry with two lone pairs is bent. Hence, the
molecular group geometry of H2O is bent and the bond angles between
atoms are less than 109.5o.
L
et us Practice
Determine the molecular geometry and bond angles between atoms of XeI4.
Key Points
Web Links
For further information, you can check the following web links:
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● Play a game of Jeopardy with your friends through this
template.
Jeopardy Labs. 2012. ‘Molecular Geometry and Bonding Theories’
https://jeopardylabs.com/play/molecular-geometry-and-bonding-theories-jeopardy
6.
1. 7.
8.
2.
3.
9.
4.
10.
5.
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B. Determine the molecular geometry of the following molecules.
1. PH3 6. BrF5
2. ClF3 7. OCl2
3. PCl3 8. SCl6
4. SeCl2 9. SO3
5. XeO4 10.CHCl3
Challenge Yourself
Answer the following questions briefly and clearly.
1. What factors affect the geometry of a molecule?
2. What keeps the outer atoms in a trigonal planar molecule 120o from each
other?
3. Explain why the bond angles between H and O atoms in water is
approximately equal to 104.5o.
4. How is a T-shaped geometry formed?
5. How is a square pyramidal geometry formed?
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Lesson 4.4: Practice Examples
Objective
In this lesson, you should be able to:
● improve current knowledge and skills in counting valence
electrons, drawing Lewis structures, and determining the
electron domain and molecular geometry of covalent
compounds.
At this point, you have already learned everything you need to know in counting
valence electrons, drawing Lewis structures, and determining the electron domain
and molecular geometry of covalent compounds. However, it takes a lot of practice
to completely master this unit. In this lesson, you will be reviewing the techniques
and apply them to several practice exercises.
Warm-Up
Procedure:
1. Given the following molecules, count the total number of valence electrons
by adding the number of valence electrons of each constituting atoms.
2. Then, draw the Lewis structure of each molecule. Make sure to satisfy octet
unless the elements are exemptions to the rule.
3. Determine the electron domain geometry by counting the total number of
electron domains present in each molecule. Recall that both bonding and
nonbonding electron pairs are classified as electron domains. Use Table 4.3
as your reference.
4. Determine the molecular geometry using Table 4.3 as your reference. Lone
pairs determine the type of molecular geometry derived from the electron
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domain geometry.
5. Determine the bond angles between atoms. Specify which atoms are
separated by these angles.
SO2
BSF
Guide Questions:
1. What criteria do you use in choosing a central atom if all atoms are of the
same number?
2. When will ED geometry be similar to molecular geometry?
Learn about It
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Step 3 Determine the group number of the element.
Use the periodic table to determine the group number of the element. The
group number will help you determine the number of valence electrons in
the element in question.
Step 4 Determine the number of valence electrons from its group number.
Use Table 4.1 and Table 4.2 as your reference. Alternatively, for
representative elements, you can figure out the number of valence
electrons directly from their CAS group number designation.
Example 1
How many valence electrons does P have?
Solution
Step 1 Identify what is required to find in the problem.
You are asked to determine the number of valence electrons in P.
Step 4 Determine the number of valence electrons from its group number.
Since P belongs to group 5A, it contains five valence electrons.
L
et us Practice
How many valence electrons does Rb have?
Example 2
How many valence electrons does Cr have?
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Solution
Step 1 Identify what is required to find in the problem.
You are asked to determine the number of valence electrons in Cr.
Step 4 Determine the number of valence electrons from its group number.
Since Cr belongs to group 6B, it contains 6 valence electrons.
L
et us Practice
How many valence electrons does Ti have?
Example 3
How many valence electrons does NH3 have?
Solution
Step 1 Identify what is required to find in the problem.
You are asked to determine the number of valence electrons in N
H3.
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Step 5 Find the answer.
L
et us Practice
Count the number of valence electrons of BF3.
A periodic table that has CAS group numbers is very helpful. Recall that
representative elements belong to group A (1A to 8A) and transition elements
belong to group B (1B to 8B). Determining the number of valence electrons of
representative metals is also different from that of the valence electrons of
transition metals.
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Step 3 Draw the skeleton structure of the molecule.
The skeleton structure is a preliminary structure where all atoms are
connected together with single bonds. As a rule, the less electronegative
atom is placed in the center and to be bonded by more electronegative
atoms. However, when the bond is between hydrogen and an
electronegative atom, the more electronegative atom is at the center of the
structure.
Worked Examples
Example 4
Draw the Lewis structure of CS.
Solution
Step 1 Determine the number of valence electrons in each constituent
element of the compound.
Since C and S belongs to Group 4A, each contains four valence electrons.
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Step 3 Draw the skeleton structure of the molecule.
Skeleton structure of CS
Lewis structure of CS
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule.
The tentative Lewis structure does not fulfill the octet rule for carbon and
sulfur. Hence, triple bonds need to be formed. The correct Lewis structure
is drawn below.
L
et us Practice
Draw the Lewis structure of sulfur dioxide.
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Example 5
Draw the Lewis structure of PF5.
Solution
Step 1 Determine the number of valence electrons in each constituent
element of the compound.
Since P belongs to Group 5A, it contains five valence electrons. Since F
belongs to Group 7A, it contains seven valence electrons.
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Lewis structure of PF5
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule.
There is no need to add multiple bonds because there were enough
electrons to be distributed. Boron also follows sextet, a form of a reduced
octet.
L
et us Practice
Draw the Lewis structure of BH3.
Example 6
Draw the Lewis structure of NO.
Solution
Step 1 Determine the number of valence electrons in each constituent
element of the compound.
Since N belongs to Group 5A, it contains five valence electrons. Since O
belongs to Group 6A, it contains six valence electrons.
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Valence electrons per Total number of valence electrons per
Number of atoms
atom atom
N 5 1 5×1=5
O 6 1 6×1=6
Total number of valence electrons in NO2 11
Skeleton structure of NO
Lewis structure of NO
Step 5 If the valence electrons seem to be deficient to satisfy the octet rule
for each element, multiple bonds can be drawn to satisfy the rule.
If you noticed, the central atom only has 5 electrons. Nitrogen has to have
at least 7. Since there are not enough electrons, putting a triple bond
would help.
Lewis structure of NO
L
et us Practice
Draw the Lewis structure of NO3-.
Always check if the atoms in your Lewis structure follow the octet rule. If not, verify
if the atoms are able to follow other rules (i.e. reduced octet, expanded octet, or
odd-electron species). Nitrogen always produces odd-electron compounds while
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smaller elements tend to follow reduced octet. Elements in period 3 and beyond
exhibit hypervalency.
Step 2 Count the total number of electron domains of the central atom.
Recall that both bonding and nonbonding electrons are classified as
electron domains. To count the total number of electron domains of a
molecule, simply get the sum of the number of bonds and the number of
lone pairs present in the central atom.
Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Table 4.3 is also useful in identifying molecular group geometries from
electron domain geometry. The molecular group geometry is a derivative
of the electron domain geometry. The two geometries are the same if no
lone pairs are initially present in the central atom of the molecule.
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Worked Examples
Example 7
Determine the molecular group geometry and the bond angles between atoms of
CCl4.
Solution
Step 1 Draw the Lewis structure.
Carbon is the least electronegative atom and it occupies the central
position. The Lewis structure of CCl4 is shown below.
Step 2 Count the total number of electron domains of the central atom.
The central atom C has four bonding pairs and has no lone pairs. The total
number of electron domains of the central atom is four.
Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Based on Table 4.3, the molecular group geometry of a tetrahedral
electron domain geometry with no lone pairs is also tetrahedral. Hence,
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the molecular group geometry of CCl4 is tetrahedral and the bond angles
between atoms are 109.5o.
L
et us Practice
Determine the molecular group geometry and the bond angles between atoms of
CHClBr.
Example 8
Determine the molecular group geometry and the bond angles between atoms of
BrF5.
Solution
Step 1 Draw the Lewis structure.
Bromine is less electronegative than fluorine. Hence, it will occupy the
central atom. It is also known to exhibit hypervalency. The Lewis structure
of BrF5 is shown below.
Step 2 Count the total number of electron domains of the central atom.
The central atom Br has five bonding pairs and has one lone pair. The total
number of electron domains of the central atom is six.
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Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Based on Table 4.3, the molecular group geometry of an octahedral
electron domain geometry with a lone pair is square pyramidal. Hence, the
molecular group geometry of BrF5 is square pyramidal and the bond
angles between atoms are less than 90o.
L
et us Practice
Determine the molecular group geometry and the bond angles between atoms of
XeOF4.
Example 9
Determine the molecular group geometry and the bond angles between atoms of
ClF3.
Solution
Step 1 Draw the Lewis structure.
Chlorine is less electronegative than fluorine. Hence, it will occupy the
central atom. It is also known to exhibit hypervalency. The Lewis structure
of ClF3 is shown below
Step 2 Count the total number of electron domains of the central atom.
The central atom Cl has three bonding pairs and has two lone pairs. The
total number of electron domains of the central atom is five.
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Step 4 Identify the molecular group geometry and the bond angles between
atoms in the molecule.
Based on Table 4.3, the molecular group geometry of a trigonal
bipyramidal electron domain geometry with two lone pairs is T-shaped.
Hence, the molecular group geometry of ClF3 is T-shaped and the bond
angles between atoms are less than 180o and 90o.
L
et us Practice
Determine the molecular geometry and bond angles between atoms of XeF3-.
Key Points
● The total number of valence electrons of a compound is the sum of all the
valence electrons of the atoms comprising it.
● Atoms follow the octet rule in covalent bonding to be stable, however, there
are some exceptions.
● The ED geometry depends solely on the number of electron domains
around the central atom, which is the sum of the number of bonding and
lone pairs.
● The molecular geometry, which is derived from the ED geometry, depends
on the number of bonding pairs a round the central atom.
● It is necessary to determine the valence electrons of the compound before
drawing the Lewis structure.
● It is necessary to draw the Lewis structure before determining the ED
geometry.
● It is necessary to determine the ED geometry before determining molecular
geometry.
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Web Links
For further information, you can check the following web links:
● Try out your skills in drawing Lewis structures with this practice
quiz.
ScienceGeek. 2014. ‘Review of Lewis Structures’
http://www.sciencegeek.net/Chemistry/Review/LewisStructures/
1. XeF2
2. SeBr6
3. NO2Cl
4. SCO
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5. CBr4
6. CHN
7. PI3
8. CS2
9. AsBr3
10. FCN
1. XeF2
2. SeBr6
3. NO2Cl
4. SCO
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5. CBr4
6. CHN
7. PI3
8. CS2
9. AsBr3
10. FCN
Challenge Yourself
Answer the following questions briefly and clearly.
1. How do you determine the number of valence electrons of a compound?
2. How do you determine the central atom?
3. How do you determine how many lone pairs must be placed at the central
atom?
4. Give three examples of elements that do not obey the octet rule.
5. What is the difference between ED geometry and molecular geometry?
Laboratory Activity
Activity 4.1
Molecular Models
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Objectives
At the end of this laboratory activity, the students should be able to:
● draw Lewis structures of the assigned molecule;
● determine its molecular geometry using the VSEPR model;
● identify the proper bond angles; and
● make a 3D model of the molecule using clay and sticks.
Procedure
1. Your teacher will assign a molecule to each pair.
2. Count the total number of valence electrons.
3. Draw the Lewis structure.
4. Identify the electron domain geometry and molecular geometry.
5. Using clay as atoms and sticks as bonds, make a 3D model of the molecule.
Use different colors for different elements. Also, consider the bond angles in
your 3D models.
Lewis structure
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Electron domain geometry
Molecular geometry
Bond angles
Guide Questions
1. What is the central atom of your molecule? Justify.
2. How many bonded pairs and lone pairs are there in your molecule?
3. How many multiple bonds are there?
Performance Task
Modeling Pollutants
There are many chemical pollutants that harm the environment. Raising awareness
about them can help a lot in reducing, if not completely eliminating them. Taking a
closer look at their structure may also help in understanding how they work.
Goal
● Your task is to make a molecular model of your chosen environment
pollutant using recycled materials.
Role
● Your job is to draw the Lewis structure of the molecule, and from there,
assess its molecular geometry and make a model.
Audience
● The target audience is your classmates and teacher.
Situation
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● The challenge involves doing research about environmental pollutants,
coming up with its molecular geometry, and using recycled materials to make
a model.
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logical manner. a way that should and easily
be rearranged for understandable by
better the audience.
understanding.
Self Check
determine the structures and shapes of molecules using the VSEPR model.
Key Words
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shell.
Valence It is the outermost energy level of an atom.
Valence electrons These are the electrons found in the valence shell.
Wrap Up
Photo Credits
Unit photo. 3D Glass Molecular Model by Purpy Pupple is licensed under public
domain via Wikimedia Commons.
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References
Brown T.L. et al. 2012. Chemistry: The Central Science. Pearson Prentice Hall.
Bettelheim F.A. et al. 2015. Introduction to General, Organic and Biochemistry.
Boston: Cengage Learning.
Ebbing, Darrell and Gammon, Steven. 2016. General Chemistry. Boston: Cengage
Learning.
Reger D.L. et al. 2009. Chemistry: Principles and Practice. Boston: Cengage
Learning.
Spencer J.N. et al. 2010. Chemistry: Structure and Dynamics. New Jersey: John Wiley
& Sons.
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Answer Key
1.
2.
3.
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Lesson 4.3: Predicting Shapes Through VSEPR
Theory
Let us Practice
1. octahedral, 90o
2. trigonal pyramidal, less than 120o
3. square planar, 90o and 180o
4.
5.
6.
7. tetrahedral, 109.5o
8. square pyramidal, less than 90o
9. T-shaped, less than 180o and 90o
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GRADES 11/12 | PHYSICAL SCIENCE
Unit 4
Structures and Shapes of
Molecules
Answer Key
B.
1. 17 6. 12
2. 21 7. 26
3. 26 8. 10
4. 20 9. 24
5. 8 10. 40
Challenge Yourself
1. Valence electrons are the electrons found in the outermost energy level of an
atom. They are the electrons involved in chemical bonding.
2. The columns are called “groups” and representative elements fall under the
same group because they have the same number of valence electrons.
3. Elements belonging to the same group have mostly the same bonding
properties because they have the same number of valence electrons.
4. Using the CAS group numbering system, the number of valence electrons of
representative elements is equal to the numerical value in the group
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number. Using the IUPAC group numbering system, the number of valence
electrons is equal to the ones digit of the group to which the element
belongs.
5. It is equal to its group number.
1. 7.
2. 8.
3.
9.
4.
5.
10.
6.
B.
3.
1.
4.
5.
2.
6.
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7.
9.
10.
8.
Challenge Yourself
1. Hydrogen only has one s orbital which can hold two electrons at most.
2. Elements in Period 3 or higher can accommodate more than eight electrons
because they have an empty d orbital which can hold extra electrons.
3. Counting the valence electrons is the first step in drawing Lewis structures. It
is important because it affects whether or not there will be lone pairs or
multiple bonds.
4. Multiple bonds may be created when there are not enough electrons to
satisfy the octet of the atoms. Electrons in multiple bonds are available to
two atoms, instead of just one.
5. The least electronegative atom is more likely to share its electrons and form
bonds with more atoms as compared to electronegative atoms. This is why
putting it at the center will yield a more stable structure.
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9. 6, octahedral 10. 4, tetrahedral
B.
1. trigonal pyramidal 6. square pyramidal
2. T-shaped 7. bent
3. trigonal pyramidal 8. octahedral
4. bent 9. trigonal planar
5. tetrahedral 10. tetrahedral
Challenge Yourself
1. The number of electron domains and the types (whether they’re bonding or
nonbonding domain) are the two main factors that affect the geometry of a
molecule.
2. Repulsion between bonding electrons keeps them as far away from each
other as possible. Since there are three atoms in a plane, the farthest
distance they get from each other is by forming a 120-degree angle in
between.
3. A trigonal pyramidal geometry is derived from a tetrahedral with one lone
pair. The lone pair is found at the top axial position, pushing the bonding
pairs downwards to form a pyramid with a triangular base.
4. A T-shaped geometry is derived from a trigonal bipyramidal with three lone
pairs. These lone pairs are found at the equatorial position, leaving the
bonding pairs to form a T-shape.
5. A square pyramidal geometry is derived from an octahedral with one lone
pair. The lone pairs are found at the axial position, leaving the bonding pairs
to form a pyramid with a square base.
1. XeF2 22
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2. SeBr6 48
3. NO2Cl 24
4. SCO 16
5. CBr4 32
6. CHN 10
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7. PI3 26
8. CS2 16
9. AsBr3 26
10. FCN 16
B.
Molecule Electron Domain Molecular Geometry
Geometry
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7. PI3 Tetrahedral Trigonal pyramidal
Challenge Yourself
1. To get the number of valence electrons of a molecule, you have to add the
valence electrons from each atom comprising the molecule.
2. The central atom is the least electronegative atom and usually the fewest in
the chemical formula.
3. To determine the number of nonbonding electrons, you have to subtract the
number of bonding electrons from the total number of valence electrons.
Distribute then these non-bonding electrons to the outer atoms first.
Whatever is left must be placed at the central atom.
4. Answers may vary. Any three from the aforementioned exceptions.
5. The ED geometry depends solely on the number of electron domains around
the central atom. The molecular geometry is derived from the ED geometry
and depends on the number of bonding pairs around the central atom.
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