Lewis Structures & the VSEPR Model A Directed Learning Activity for Hartnell College Chemistry 1
Funded by the Title V – STEM Grant #P031S090007 through Hartnell College
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Student Learning Objectives This activity will help you to: 1. Use a Periodic Table to help determine
valence electrons & draw Lewis structures.
2. Examine the relationship between atomic electronegativities & bond polarity
3. Use VSEPR Theory and bond electronegativities to determine molecular shape & molecular polarity
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Getting Started
This set of Power Point slides will lead you through a series of short lessons and quizzes on the topics covered by this Directed Learning Activity tutorial.
Move through the slideshow at your own pace.
There are several hyperlinks you can click on to take you to additional information, take quizzes, get answers to quizzes, and to skip to other lessons.
You can end this slide show at any time by hitting the “ESC” key on your computer keyboard.
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Table of Topics What You Should Already Know Drawing Lewis Structures Determining Molecular Shapes Discovering Molecular Polarity
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What You Should Already Know General layout of the Periodic Table Electronic configurations Valence electrons Types of bonding Electronegativity trends
If you are unsure of these topics, please refer to your lecture text for more information. Next
Drawing Lewis Structures Looking at the Electrons
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Electrons and Bonding Valence electrons are the highest energy atomic electrons that participate in bonding between atoms. They are considered the electrons that are the farthest from the atomic nucleus. The Lewis structure is used to give a two-dimensional representation of covalent molecules and ions. Lewis structures are based on using dots and lines to represent bonding and non-bonding valence electrons. Lines are used to represent bonding pairs of electrons. Dots are used to signify electrons. Lewis structures may also be used to show valence electrons around individual atoms or ions.
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Electrons and Bonding cont’d For the representative elements (those in Groups IA through VIIIA or 1A through 8A of the periodic table) the number of valence electrons is typically the Group number. The periodic table in your text probably looks very similar to the one on the next page. The number of valence electrons for the transition metals (those in Groups labeled with a B) are more variable and will not be covered in detail in this tutorial.
Periodic Table
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Lewis Structure Procedures for “A” Elements: Single Atoms 1. Find the element on the periodic table. 2. Look at the Group designation at the top of
the column. It must end with the letter “A” to use this method.
3. The Group number gives the normal number of valence electrons for the neutral atom.
4. Write the symbol for the element. Draw a small dot to represent each valence electron.
a. Depending on the number of valence electrons, the dots should be placed around the top, bottom and sides of the symbol.
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Single Atoms cont’d b. Start by not pairing any electrons. After
four electrons, you begin pairing electrons. c. The maximum number of electrons around
the atom will be 8. Here are a few examples –
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Lewis Dot Structure Examples
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Single atoms cont’d 5. For single atom ions
a. For cations (positive ions), subtract one electron from the total for the molecule for each positive charge
b. For anions (negative ions), add one electron to the total for each negative charge
c. The Octet Rule states that the preferred number of valence electrons for the representative element atoms is eight (the Noble Gas configuration), provided the electronic configuration of the atom can accommodate eight valence electrons.
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Examples of Anions & Cations
The Lewis dot symbol for the F – ion showing 8 valence electrons
The Lewis dot symbol for the H+ ion showing no valence electrons – one of the exceptions to the Octet Rule
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Covalent Molecules & Polyatomic Ions 1. Calculate the total number of valence
electrons a. Use the Group number for the number of
valence electrons for each atom in a neutral molecule and add these together for every atom in the molecule or ion
b. For cations, subtract one electron from the total for each positive charge
c. For anions, add one electron to the total for each negative charge
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Covalent Molecules & Polyatomic Ions cont’d 2. Write the skeleton structure of the molecule
or ion a. This may have to be done by trial and error b. For simple molecules
• The central atom is typically surrounded by atoms of greater electronegativity
• H may surround a more electronegative element because it can only have one bond
• Try to draw symmetrical molecules as a first guess
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Covalent Molecules & Polyatomic Ions cont’d 3. Connect the atoms with a pair of dots or a line to
represent a covalent bond 4. Distribute electrons around the atoms surrounding
the central atom(s), obeying the Octet Rule for atoms with the capacity for eight electrons. The next slide contains some of the most common bonding patterns for nonmetals
5. Distribute remaining electrons as pairs to central atom(s)
6. If there are fewer than eight electrons on the central atom(s) form multiple bonds to satisfy the Octet Rule
7. The symbol for a polyatomic ion is placed in brackets with the charge shown outside
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Most Common Bonding Patterns for Nonmetals
Element # bonds # lone pairs
H 1 0 C 4 0
N, P 3 1 O, S, Se 2 2
F, Cl, Br, I 1 3
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Examples H is an exception to the Octet Rule
CH4 (methane) C satisfies the Octet Rule
Carbon monoxide requires the use of multiple bonds to satisfy the Octet Rule
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Quiz Questions Draw the Lewis structures for – BF3
NH3
H2O NO3
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Check answers
Quiz Answers
B is one of the exceptions to the Octet Rule
Note that NO3- has three possible forms, where the
double bond is in three different places. This property is referred to as “resonance”. Also, the symbol for a polyatomic ion is placed in brackets with the charge outside the brackets.
Next lesson Go back to review lesson
Determining Molecular Shape
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VSEPR Model – Predicting Molecular Shapes Valence Shell Electron Pair Repulsion
Bonding pairs of electrons on a central atom are
oriented as far away from each other as possible in the space around the central atom to minimize the repulsion between areas of negative charge (the valence electrons)
Nonbonding pairs of valence electrons take more space than bonding pairs and slightly reduce the spacing between the bonding pairs
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Using the VSEPR Model Draw the Lewis structure of the molecule
or polyatomic ion using the method in the previous lesson
Look at the Lewis structure to see how many atoms (bonding pairs) and lone pairs (unshared pairs) of electrons are bonded to the central atom(s)
Use the table on the next slide to determine the molecular geometry (there is probably a similar chart in your lecture text)
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Common Bonding Patterns
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Examples
CH4 (methane) 4 atoms around central C atom – tetrahedral geometry
NH3 (ammonia) 3 atoms and one lone pair around central N atom – trigonal pyramidal geometry
H2O (water) 2 atoms and two lone pairs around central O atom – bent geometry
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Quiz Questions Determine the molecular geometry of the central atom in the following species: 1. PO4
3-
2. SOF2
3. NO2
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Check answers
Quiz Answers 1. PO4
3- tetrahedral 32 valence electrons 4 atoms around P 2. SOF2 trigonal pyramidal 26 valence electrons around S 3 atoms + 1 nonbonding pair around S
3. NO2- bent
18 valence electrons 2 atoms + 1 nonbonding pair around N
Next lesson Review lesson
Discovering Molecular Polarity
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Bond Polarity Remember that there are two general types of bonds: covalent and ionic. Covalent bonds imply that there is at least one shared pair of electrons in the bond between two atoms. Covalent bonds can be classified as polar or nonpolar. In polar covalent bonds, there is a difference in electronegativity between the two atoms – the more electronegative atom has the partial negative charge and the less electronegative atom has the partial positive charge. Ionic bonds imply that the two species involved are ions.
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More on Bonding To predict whether a bond is going to be nonpolar covalent, polar covalent, or ionic, you can look at the difference in electronegativity (ΔEN) between the atoms and other clues. If ΔEN is less than 0.4, the bond is most likely
nonpolar covalent. If ΔEN is between 0.4 and 1.7, the bond is
most likely polar covalent. If ΔEN is greater than 1.7, the bond is most
likely ionic. Non-metal to non-metal bonding is always
covalent. Metal to non-metal bonding is usually ionic.
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Symmetry Just because a molecular compound has polar bonds does not necessarily mean that the molecule is polar overall. How can that be? We have to consider if there is symmetry in the molecules that make up the compound. Let’s consider CO2. By now you know that it is a linear molecule. If we look at the ΔEN in the C-O bond, ΔEN = 3.44 - 2.55 = 0.89.
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This would make each of the bonds in CO2 polar covalent. But if we look at the shape of the molecule, it is linear. The two oxygen atoms are partially negatively charged, while the central carbon atom is partially positively charged. So, the molecule itself cannot be polar – there is no way for one end of the molecule to be negative relative to the other end, because both ends are negative. So while the bonds are polar, the molecule itself has a symmetrical charge distribution and is not polar.
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How about CF4? If we analyze the bonds, we will find that each of the four bonds is polar. However, when we look at the overall shape of the molecule, it is tetrahedral. Again, this is a molecule with symmetry and it has a symmetrical charge distribution. So although each of the individual bonds is polar, the overall polarity of the molecule is nonpolar.
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Quiz Questions
Which of the following molecular compounds is polar? CH4 H2O HBr CHCl3
Check answers
Quiz Answers CH4 H2O HBr CHCl3 CH4 is tetrahedral and has nonpolar bonds – therefore nonpolar molecule H2O is a bent molecule and has two polar bonds and two lone pairs of electrons, making it asymmetrical in charge distribution – therefore polar HBr is a linear molecule and has a polar bond, making it asymmetrical in charge distribution – therefore polar CHCl3 is tetrahedral and has three polar bonds and one nonpolar bond making it asymmetrical in charge distribution – therefore polar Next Review lesson
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