Covalent Bonding Note: Students and classrooms with iPads should - - PDF document
Slide 1 / 186 Slide 2 / 186 Covalent Bonding Note: Students and classrooms with iPads should download the free "Lewis Dots" App and can use that on all the slides where Lewis Dot drawings are to be done. Slide 3 / 186 Table of
Note: Students and classrooms with iPads should download the free "Lewis Dots" App and can use that on all the slides where Lewis Dot drawings are to be done.
· Properties of Ionic and Covalent Materials
Click on the topic to go to that section
· Naming Binary Molecular Compounds · VSEPR Theory · Covalent versus Ionic Bonds · Resonance Structures · Molecular Geometry · Lewis Structures · Polarity
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Examine these two forms of the same compound, ibuprofen.
This form of ibuprofen has virtually no anti-inflammatory effect. This form of ibuprofen is about 100x more effective at alleviating pain than the other form. Even though they consist of the exact same number and kinds of atoms, these two molecules have very different chemical properties.
In this unit, we will explore what causes molecules to have various shapes. Later, we will then examine how molecular geometry affects different chemical properties.
Take a look around you. The chemistry of everything you see, hear, feel, touch and taste is a result of not only what it's made of but also how it's put together.(Remember this for next year in biology!)
Ionic - The electrostatic attraction between ions Covalent - The sharing of electrons between atoms Metallic - Each metal atom bonds to
electrons (covered in a later unit) Chemical bonds hold atoms together to create chemical
How ionic or covalent a bond is depends on the difference in
are "shared" and the bond is considered covalent, the greater the difference, the more likely electrons have been transferred and the atoms are ionized resulting in an ionic bond. Li Be B C N O F
Electronegativity 1.0 1.6 2.0 2.5 3.0 3.5 4.0 Bond Li-F Be-F B-F C-F N-F O-O F-F Electronegativity 3 2.4 2.0 1.5 1 0.5 0
Increasing Covalent Character
We can make a few simplifications... Ionic Bonding Ionic bonds occur when the difference in electronegativity between two atoms is more than 1.7. Na ---- F electronegativity = 3 Covalent Bonding If the difference of electronegativity is less than 1.7, neither atom takes electrons from the other; they share electrons. This type of bonding typically takes place between two non-metals or between two metals. H ---- Cl electronegativity = 1.1
In the case of ionic bonding, a 3-D lattice of ions is the result . . . not individual molecules. The chemical formula for an ionic compound is just the ratio of each type of ion in the lattice, not a particular number of ions in a molecule. In contrast, covalent bonding can result in individual molecules or 3-D lattices depending on the elements
shapes of these molecules help determine the physical and chemical properties of everything around us!
click here for an animation about ionic and covalent bonding
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Since the attractions between the ions span a short distance, these forces are quite strong resulting in high melting points and boiling points! Na+ -- Cl- it takes a lot of energy to break an ionic lattice! Compound Melting Point (C) NaCl 801 MgO 2852
Since ionic compounds consist of ions, when these ions are free to move, the substance can conduct electricity. To move, they must be in the liquid or molten state.
NaCl (s) Molten NaCl(l) Lattice is strong, no conductivity Lattice is broken, ions are free to move and conduct
Metallic compounds are held together by non-directional covalent bonds in which some electrons are shared but are loosely held and free to roam. The covalent bonds between the metal atoms are strong! This gives rise to high melting and boiling points! Metallic Lattice strong metallic covalent bonds Metal Melting Point Cu 1085 C Fe 1585 C
In order to obtain pure metals, the ancients had to melt the metal (metallic substance) out of the rock (an ionic compound). Copper has a lower melting point so it could be obtained in furnaces at lower temperatures. Furnaces hot enough to extract iron would come later.
Why do you think the bronze age (copper mixed with tin) came before the iron age?
Since the electrons in metals are free to roam somewhat, metals are good conductors of electricity! Silver is the most conductive metal and is roughly 5-10 times more conductive than steel (mostly iron).
Copper is often used in electrical cable rather than silver even though it is roughly 10% less conductive than silver. Why?
Copper currently trades for roughly 3 dollars an ounce while silver trades for about 30 dollars a month. It's about the money!!!!
5 Which of the following would NOT conduct electricity in the solid state?
A Al B Al2O3 C NaCl D
Both A and B
E
Both B and C
5 Which of the following would NOT conduct electricity in the solid state?
A Al B Al2O3 C NaCl D
Both A and B
E
Both B and C
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Like ionic and metallic substances, covalent network solids are giant molecules arranged in 3-D crystalline shapes. Here, the atoms involved tend to semi-metals like Silicon or Germanium or elemental carbon. Since the bonds are covalent, they are quite strong! This gives rise to high melting and boiling points!
Glass (75% SiO2) Diamond (pure C) Melts at 1500 C Melts at 3500 C
Since these substances have higher electronegativities, they keep good tabs on their electrons thereby preventing the electrons from
Diamond and graphite are both allotropes or different versions of carbon and vary somewhat in their conductivity.
Diamond (C) Graphite (C) non-conductive a little conductive
Diamond is notorious for being HARD! This is true for lots of covalent network crystals. Can you think of some applications where hardness is important? Body Armor B4C (boron carbide) Drill Bits polycrystalline diamond
6 Which of the following would be classified as a covalent network solid?
A NaCl B HF C CO2 D
Ge2O3
E
Fe
6 Which of the following would be classified as a covalent network solid?
A NaCl B HF C CO2 D
Ge2O3
E
Fe
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When atoms are bonded covalently, the atoms are held together by sharing electrons. This occurs between non-metals such as C,O,S,H,P,N, etc. Unlike in all of the other substances, the atoms form small individual molecules that then interact with each other and their environment. These are called molecular compounds. P O H H O = C = O Cl Cl Cl In covalent bonds, electron sharing usually occurs so that atoms attain the electron configurations of noble gases. Both atoms use the shared electrons to reach that goal. Click here to view interactive website
Since these substances contain lots of small molecules, the bonds holding these small molecules together are fundamentally different from the covalent bonds found inside the molecule.
weak inter-molecular forces between molecules
They cover a much larger distance and are quite weak giving rise to LOW melting and boiling points!
Molecular compounds contain electronegative non-metals and do not lose their electrons easily so they are non-conductive. As a result they are excellent INSULATORS!
Rubber: (C5H9)250
Ionic Metallic
Molecular
metals and non- metals metals semi-metals and pure carbon non-metals Na2O Fe
C(diamond)
CH4 High MP High MP High MP Low MP conduct as liquid conduct in all states non-conductive non-conductive Brittle Malleable Brittle Brittle
7 Which of the following would have the lowest melting point?
A N2 B C(graphite) C C(diamond) D
W
E
LiF
7 Which of the following would have the lowest melting point?
A N2 B C(graphite) C C(diamond) D
W
E
LiF
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8 Which of the following will not conduct electricity in any state?
A Cu B NaF C Fe D
CO2
E
All of these will conduct
8 Which of the following will not conduct electricity in any state?
A Cu B NaF C Fe D
CO2
E
All of these will conduct
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9 Which of the following consists of small individual molecules?
A C(diamond) B SiO2 C Cu2O D
Na
E
SO3
9 Which of the following consists of small individual molecules?
A C(diamond) B SiO2 C Cu2O D
Na
E
SO3
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10 Which of the following substances has both ionic and covalent bonding within the crystal?
A Cu B CuCO3 C LiCl D
Ba
E
BaF2
10 Which of the following substances has both ionic and covalent bonding within the crystal?
A Cu B CuCO3 C LiCl D
Ba
E
BaF2
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Use prefixes to indicate the number the atoms. All end in "ide" Examples NO2 nitrogen dioxide P2O5 diphosphorous pentoxide ( penta-oxide-->pentoxide)
Look on your reference sheets for the prefixes. The atom with the lower electronegativity is usually written first. If there is only one of the first atom, the mono- is left off. Examples CO carbon monoxide CO2 carbon dioxide
11 Chlorine monoxide is
ClO2
O2Cl
11 Chlorine monoxide is
ClO2
O2Cl
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12 Dinitrogen tetroxide is A NO2 B
N2O4
NO3-
D N4O2
12 Dinitrogen tetroxide is A NO2 B
N2O4
NO3-
D N4O2
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13 H2O is
A
Hydrogen monoxide B
Dihydrogen monoxide
C Hydrogen oxide D Hydrogen dioxide
13 H2O is
A
Hydrogen monoxide B
Dihydrogen monoxide
C Hydrogen oxide D Hydrogen dioxide
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14 SO3 is
14 SO3 is
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15 MgO is
A monomagnesium monoxide
B magnesium monoxide C monomagnesium oxide
D magnesium oxide
15 MgO is
A monomagnesium monoxide
B magnesium monoxide C monomagnesium oxide
D magnesium oxide
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16 P4O10 is
Tetraphosphorous decoxide
16 P4O10 is
Tetraphosphorous decoxide
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Lewis structures are diagrams that show valence electrons as dots. Lewis structures are also known as Lewis dot or electron dot diagrams. Note that no electrons are paired until after the fourth one.
17 How many valence electrons does nitrogen have?
17 How many valence electrons does nitrogen have?
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18 The Lewis structure for nitrogen is N True False
18 The Lewis structure for nitrogen is N True False
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Recall that atoms tend towards having the electron configuration of a noble gas.For most atoms, that means having 8 valence electrons. The Octet Rule also applies to molecular compounds. In covalent bonding, an atom will share electrons in an effort to
attempt to obtain 2 valence electrons).
A pair of valence electrons that is not shared between atoms is called an unshared pair, also known as a lone pair or a nonbonding pair.
Exceptions to the Octet Rule H needs 2e- Be needs 4e- B needs 6e-
An electron dot structure such as H:H represents the shared pair of electrons of the covalent bond by two dots.
H + H H H
Hydrogen atom Hydrogen atom Hydrogen molecule Shared pair
H H
1s 1s
Hydrogen molecule
A structural formula represents the covalent bonds by dashes and shows the arrangement of covalently bonded
is represented by one dash.
A 2 B 1
A 2 B 1
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The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example.
F F F F F F +
# # #
OR
Fluorine atom Fluorine molecule Fluorine atom
1s 2s 2p 1s 2s 2p Fluorine molecule
In a water molecule, each hydrogen and oxygen atom attains a noble-gas configuration by sharing electrons.
The water molecule has two unshared, or lone, pairs of electrons. 2 H + O --> O H or O H H H
Hydrogen atoms Oxygen atom Water molecule
1s 2p 2s 1s 1s
O H H
Water molecule
In the ammonia molecule, NH3, each atom attains a noble- gas configuration by sharing electrons. This molecule has one unshared pair of electrons.
3 H + N --> N H or N H H H H
Hydrogen atom Nitrogen atom Ammonia molecule 1s 2p 2s 1s 1s
H N H
1s Ammonia molecule
H
electronegative element (excluding hydrogen).
by single bonds.
P has an electronegativity of 2.1 and Cl has an electronegativity of 3.0 P will be the central atom. The Cl atoms will surround the P atom. The single bonds are shown as single lines.
(two) of electrons.
give each one 8 (a full shell), or just 2 electrons for hydrogen.
hydrogen)? Have you used up all the valence electrons? Have you used too many electrons?
The N atom has 5 valence electrons and each of the three H atoms has 1 so the total number of valence electrons is,
5 + 3(1) = 8
electrons in the polyatomic ion or molecule.
electronegative element (excluding hydrogen because it can only have one bond).
by single bonds. H can never be the central atom so N must be The H atoms will surround the N atom. The single bonds are shown as single lines.
Each H already has two electrons, so that's done. But we have to add electrons to N to make 8.
(two) electrons. Now add electons to the outer atoms to give each one a full shell (2 in the case of H).
atom.
Does each atom have a full outer shell ?
electrons you started with? Have you used too many electrons?
20 How many total valence electrons does H
2O have?
A B C D
20 How many total valence electrons does H
2O have?
A B C D
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21 Which element in H 2O is the least electronegative?
A B
21 Which element in H 2O is the least electronegative?
A B
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22 Which of the following is the correct Lewis Structure for H2O?
A B C D
(except H only gets 2).
22 Which of the following is the correct Lewis Structure for H2O?
A B C D
(except H only gets 2).
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23
A B C
Which of the following is the correct Lewis Structure for C2H6?
D
23
A B C
Which of the following is the correct Lewis Structure for C2H6?
D
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If you are drawing the Lewis Structure for an ION... A negative ion has extra electrons, add the charge of the ion to your valence electron count. ClO2
A positive ion is missing electrons, subtract the charge of the ion to your valence electron count. NH4
+ has 1(5) + 4(1) -1 = 8 electrons
24
A B
How many valence electrons does CO3
2- have?
C D
24
A B
How many valence electrons does CO3
2- have?
C D
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25
A B
How many valence electrons does H3O+ have?
C D
25
A B
How many valence electrons does H3O+ have?
C D
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The "Formal Charge" method tells us how the electrons are distributed within a molecule. For example, depending on how the electrons are shared, some atoms may have more electrons than others resulting in a semi-charged state for that atom. O P O O O FC for P: 5 - 4= +1 (count each bond as one) FC for each O: 6 -7= -1 (count each bond as one)
Note: The charges must add to the charge of the molecule. So for PO4
3-
1 P atom x +1 = +1 + 4 O atoms x -1 = -4 +1 + -4 = -3
Formal Charge = # of valence electrons - # of electrons atom possesses within the lewis structure.
The best Lewis structure will have the formal charge = 0 on each
electronegative atoms should carry a charge as they have the greater attraction for electrons! Each bond is counted as one in a formal charge calculation as each atom forming part of the bond contributes just one electron to that bond. [ O - H ]-1 FC on O = 6-7 = -1 FC on H = 1-1 = 0
O H
The oxygen is more electronegative so it makes sense that it carries the negative charge.
Example: Below are two possible lewis structure for the phosphate ion, PO4
3-. Which Lewis structure is considered to more closely
represent the actual molecule based on formal charge calculations?
O P O O O O P O O O
Structure 2 is superior as all formal charges = 0 whereas in structure 1, the P carries a +1 charge and each oxygen carries a -1 charge Structure 1 Structure 2
26 Which of the following would be the formal charge
A +1 B 0 C -1 D -2 E -3
26 Which of the following would be the formal charge
A +1 B 0 C -1 D -2 E -3
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27 In which of the following molecules would N carry a
non-zero formal charge?
A
HCN
B
NH3
C
NO3-
D
NO2-
E
NH4+
27 In which of the following molecules would N carry a
non-zero formal charge?
A
HCN
B
NH3
C
NO3-
D
NO2-
E
NH4+
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Draw the Lewis dot structure for the sulfate ion, SO4 2-, and find the formal charge on each atom. FC on S = 6-4 = +2 FC on O = 6-7 = -1
Draw the Lewis dot structure for the hydronium ion, H3O+ and find the formal charge on each atom. FC on O = 6-5 = +1 FC on H = 1-1 = 0 * note how in this case the more electronegative atom (O) is carrying a + charge relative to H. This demonstrates the theory is imperfect.
C N Cl F O S B P I H
C O O
Si Se Xe
Draw a Lewis Structure We ran out of electrons, but carbon does not have an octet yet!
Now What?
Slide for Answer
Atoms form double or triple covalent bonds if they can attain a noble gas structure by sharing two pairs or three pairs of electrons. A bond that involves two shared pairs of electrons is a double covalent bond. A bond formed by sharing three pairs of electrons is a triple covalent bond.
Carbon Dioxide, CO2
1 (4) + 2 (6) = 16 e- This leaves 12 electrons, 6 pairs
work with; how many have we used?
BONDS between C and O. Instead of sharing only 1 pair, a double bond shares 2 pairs. So one pair is taken away from each atom and replaced with another bond.
It requires more energy to break double and triple bonds compared to single bonds. Triple bonds are the strongest of the three.
Bond Type Bond Energy
C C C C C C 348 kJ 614 kJ 839 kJ
Bond Type Bond Energy
N N 163 kJ 418 kJ 941 kJ N N N N
Type of Bond Electrons shared Bond Strength Bond Length 2 4 6 weak intermediate strong long intermediate short
28 As the number of bonds between a pair of atoms
increases, the distance between the atoms:
remains unchanged
varies, depending on the atoms
28 As the number of bonds between a pair of atoms
increases, the distance between the atoms:
remains unchanged
varies, depending on the atoms
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29 As the number of bonds between a pair of
atoms increases, the strength of the bond between the atoms:
varies, depending on the atoms
29 As the number of bonds between a pair of
atoms increases, the strength of the bond between the atoms:
varies, depending on the atoms
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30 As the number of bonds between a pair of atoms
increases, the energy of the bond between the atoms:
varies, depending on the atoms
30 As the number of bonds between a pair of atoms
increases, the energy of the bond between the atoms:
varies, depending on the atoms
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31 How many electrons are shared by two
atoms to create a single bond?
31 How many electrons are shared by two
atoms to create a single bond?
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32 How many electrons are shared by two
atoms to create a double bond?
32 How many electrons are shared by two
atoms to create a double bond?
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33 How many electrons are shared by two
atoms to create a triple bond?
33 How many electrons are shared by two
atoms to create a triple bond?
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34 Using Lewis structure drawings, determine which
molecule below would have the shortest bond length between atoms?
A O2 B F2 C Cl2 D
CO
E
I2
34 Using Lewis structure drawings, determine which
molecule below would have the shortest bond length between atoms?
A O2 B F2 C Cl2 D
CO
E
I2
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35 Which of the following molecules would have the
longest C-O bond length? Use Lewis structures.
A CO B CO2 C H2CO D CH3OH E The lengths are all the same
35 Which of the following molecules would have the
longest C-O bond length? Use Lewis structures.
A CO B CO2 C H2CO D CH3OH E The lengths are all the same
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If you run out of electrons before the central atom has an octet……form multiple bonds until it does.
Oxygen molecule
1s 2s 2p 1s 2s 2p
O + O --> O O or O O O O
Oxygen atom Oxygen atom Oxygen molecule
Oxygen molecule
C N Cl F O S B P I H
C
Si Se Xe
Draw a Lewis Structure Carbon has the lower electronegativity, so we will consider it the "central" atom... O
Slide for Answer
In carbon monoxide, oxygen has a stable configuration but the carbon does not.
1s 2p 2s 2s 1s 2p
C + O # # # C O
Carbon atom Oxygen atom Carbon monoxide
C O
Carbon monoxide molecule
A coordinate covalent bond is a covalent bond in which one atom contributes both bonding electrons. In a structural formula, you can show coordinate covalent bonds as arrows that point from the atom donating the pair of electrons to the atom receiving them. In a coordinate covalent bond, the shared electron pair comes from one of the bonding atoms. Carbon has 4 valence electrons, oxygen has 6.
C N Cl F O S B P I H Si Se Xe
Draw a Lewis Structure
F F Slide for Answer F
A molecule is a neutral group of atoms joined together by covalent
A diatomic molecule is a molecule consisting of two atoms. Certain elements do not exist as single atoms; they always appear as pairs. When atoms turn into ions, this NO LONGER HAPPENS!
Hydrogen Nitrogen Oxygen Fluorine Chlorine Bromine Iodine
H H N N O O
H2 N2 O2
36 On the periodic table below, mark which elements
exist as diatomic molecules. Note the pattern.
There are three types of ions or molecules that do not follow the octet rule: #1 Ions or molecules with an odd number of electrons #2 Ions or molecules with less than an octet #3 Ions or molecules with more than eight valence electrons (an expanded octet)
Though relatively rare and usually quite unstable and reactive, there are ions and molecules with an odd number of electrons. NO is an example:
Beryllium (Be) - this metal is shown to form molecular compounds, rather than ionic compounds as expected; only needs 4 electrons to be stable Boron (B) - only needs 6 electrons to be stable Memorize these exceptions
The only way PCl5 exists is if phosphorus has 10 electrons around it. This is called an expanded octet. Atoms on the third energy level or higher are allowed to expand their octet to 10 or 12 electrons. These atoms are larger and can accommodate more electrons.
How many electrons do these central atoms have around them?
Draw the Lewis dot structure for sulfur hexaflouride, SF6:
Draw the Lewis dot structure for the xenon tetrafluoride, XeF4.
Draw the Lewis dot structure for boron trifluoride, BF3:
Draw the Lewis dot structure for the iodine tricholoride, ICl3.
37
Boron and Beryllium
Boron and Helium
Boron, Beryllium, and Hydrogen
Boron, Beryllium, Hydrogen and Helium
Boron, Beryllium, Hydrogen, Helium and Oxygen [*] Which of the following need fewer than 8 valence electrons to be stable?
37
Boron and Beryllium
Boron and Helium
Boron, Beryllium, and Hydrogen
Boron, Beryllium, Hydrogen and Helium
Boron, Beryllium, Hydrogen, Helium and Oxygen [*] Which of the following need fewer than 8 valence electrons to be stable?
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38 The correct lewis structure for BeCl2 is
Cl - Be - Cl
True False
38 The correct lewis structure for BeCl2 is
Cl - Be - Cl
True False
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39 Elements in the first two rows of the periodic table
cannot have expanded octets because their atoms do not have enough space.
True False
39 Elements in the first two rows of the periodic table
cannot have expanded octets because their atoms do not have enough space.
True False
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C N Cl F O S B P I H Si Se Xe
Draw a Lewis Structure and use that to determine the VSEPR number For the central oxygen: Electron domains = 3 Bonding domains = 2 Unpaired electrons = 1 Its VSEPR number is 3 2 1
O O O Slide for Answer
Consider the Lewis structure we would draw for ozone, O3: We would expect the double bond to have a shorter bond length than the single bond. However, the true, observed structure of ozone shows that both O-O bonds are the same length. How can this be?
O O O O O O [*]
One Lewis structure cannot accurately depict a molecule like ozone. Therefore, we use multiple structures, called resonance structures, to describe the molecule. Ozone has two resonance structures.
[*]
The actual ozone molecule is a synthesis of these two resonance structures. The bond length for both outer oxygen atoms falls somewhere between the single and double bond length. O O O O O O
Resonance structure Resonance structure
Ozone molecule
[*]
The nitrate ion, NO31- also requires resonance structures to explain its covalent bonding. There are three resonance structures for the nitrate ion:
[*]
Draw the Lewis dot structure for SO3:
[*]
40 How many resonance structures can be drawn for
the carbonate ion, CO32- ?
[*]
40 How many resonance structures can be drawn for
the carbonate ion, CO32- ?
[*]
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The benzene molecule is a regular hexagon of carbon atoms with a hydrogen atom bonded to each one. There are two resonance structures for benzene.
Benzene, C6H6, is obtained from the distillation of fossil fuels. More than 4 billion pounds of benzene is produced annually in the United States. Because benzene is a carcinogen, its use is closely regulated.
[*]
In truth, the shared pairs of electrons do not always remain between adjacent C atoms. They are not localized. Instead, the electrons are said to be delocalized, meaning that they they can move around the 6-carbon ring. Benzene is commonly depicted as a hexagon with a circle inside to signify the delocalized electrons in the ring... we will talk more about this at the end of the year when we study
# # # #
[*]
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According to VSEPR theory, the molecules will adopt a shape/geometry so as to reduce the repulsion between the bonded electrons. Click here to view a PhET simulation
The VSEPR number of a molecule is a three digit number that can be used to determine a molecule's shape. Here's how you find it:
atom, if applicable.
electron-domains around the central atom.
Electron domains are either shared pairs of electrons
Multiple bonds (i.e. double or triple bonds) count as
lone pairs around the central atom.
second and third.
Bonding domains are single, double or triple bonds. Each pair of electrons that are not involved in bonds counts as one lone pair.
41 How many electron domains does CH4 have?
A 1 B 2 C 3 D 4 E 5
41 How many electron domains does CH4 have?
A 1 B 2 C 3 D 4 E 5
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42 How many electron domains does H2O have?
A 1 B 2 C 3 D 4 E 5
42 How many electron domains does H2O have?
A 1 B 2 C 3 D 4 E 5
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43 How many electron domains does CO2 have?
A 1 B 2 C 3 D 4 E 5
C O O
43 How many electron domains does CO2 have?
A 1 B 2 C 3 D 4 E 5
C O O
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C N Cl F O S B P I H Si Se Xe
Draw a Lewis Structure and use that to determine the VSEPR number
Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Its VSEPR number is 4 4 0
C N Cl F O S B P I H Si Se Xe
Draw a Lewis Structure and use that to determine the VSEPR number
N F F F
Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 3 Lone pairs of electrons = 1 Its VSEPR number is 4 3 1
Slide for Answer
C N Cl F O S B P I H Si Se Xe
Draw a Lewis Structure and use that to determine the VSEPR number
F Si F F F
Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Its VSEPR number is 4 4 0
Slide for Answer
C N Cl F O S B P I H Si Se Xe
Draw a Lewis Structure and use that to determine the VSEPR number
O P O O O
Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 4 Bonding domains = 4 Lone pairs of electrons = 0 Its VSEPR number is 4 4 0
Slide for Answer
F C N Cl F O S
Draw a Lewis Structure and use that to determine the VSEPR number
F I F F F
Check to make sure that each atom has a full outer shell. Now calculate the VSEPR #. Electron domains = 6 Bonding domains = 5 Lone pairs of electrons = 1 Its VSEPR number is 6 5 1
Slide for Answer
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According to VSEPR theory, the repulsion between electron pairs causes molecular shapes to adjust so that the valence-electron pairs stay as far apart as possible. The shape of a molecule plays an important role in determining its chemical and physical properties. To determine a molecule's shape, i.e. its molecular geometry, we must first determine its electron-domain geometry.
Recall: Electron domains are either shared pairs of electrons
Bonding domains are single, double or triple bonds. Each pair of electrons that are not involved in bonds counts as one lone pair. To determine the electron-domain geometry, look at the first number and use the following chart...
The EDG (2,3,4,5,or 6) gives us the general shape of the molecule, as shown here. However, these domains do not have to be bonds. The molecular geometry tells us if there is a bond or lone pair of electrons present, thereby specializing the general shape. Let's take a closer look...
Linear
Two atoms around a central one will form a linear shape with bond angles of 180o
There is only one molecular geometry for linear electron-domain: linear molecular geometry (220).
trigonal planar
Three atoms around a central one will form a trigonal planar shape with bond angles of 120o
There are two molecular geometries: · Trigonal planar, if all the electron domains are bonding (330) · Bent, if one of the domains is a nonbonding pair (321)
120 trigonal planar (330) 117 bent (321)
It is very important to note that unbonded pairs of electrons repel more strongly than bonded electrons thereby shrinking the bond angle between atoms
Four atoms around a central one will form a tetrahedral shape with bond angles of 109.5o tetrahedral
There are three molecular geometries: Tetrahedral, if all are bonding pairs (440) Trigonal pyramidal, if one is a nonbonding pair (431) Bent, if there are two nonbonding pairs (422)
tetrahedral (440) trigonal pyramidal (431) bent (422)
109.5 107 104.5
Again, note the decrease in bond angle as the number of high repelling unbonded pairs of electrons increase.
Five atoms around a central one will form a trigonal bipyramidal shape with bond angles of 120o and 90o trigonal bipyramidal
Trigonal bipyramidal Seesaw T-shaped Linear
Trigonal Bipyramidal (550) See-Saw (541) T-Shape (532) Linear (523)
There are four molecular geometries for the trigonal bipyramidal electron domain geometry:
Six atoms around a central one will form an octahedral shape with bond angles of 90o
Square Planar Octahedral Square Pyramidal
Octahedral (660) Square Pyramidal (651) Square Planar (642)
There are only three molecular geometries for the octahedral electron domain geometry:
Using VSEPR numbers, you can determine molecular geometry. VSEPR numbers are a set of 3 numbers. 1) the total number of electron domains 2) the number of bonding domains* 3) the number of unshared pairs
Electron-domain geometry has the same name as the first shape. (*Remember that multiple bonds count as ONE domain)
Draw the Lewis structure for ammonia, NH3. What are the VSEPR numbers for NH3? 4,3,1 What is the electron-domain geometry of NH3? tetrahedral What is the molecular shape of NH3?# # # What is the N-H bond angle in the molecule? 107 What is the formal charge on the N atom? 5-5 = 0
triangular pyramidal
Draw the Lewis structure for ClF3. What are the VSEPR numbers for ClF3? 5,3,2 What is the electron-domain geometry of ClF3? trigonal bipyramidal What is the molecular shape of ClF3? T What would be the Cl-F bond angle(s)? 180, 90 What would be the formal charge on Cl? 7-7 = 0
44 The methane molecule (CH4) has which geometry?
44 The methane molecule (CH4) has which geometry?
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45 Give the VSEPR number for this molecule.
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45 Give the VSEPR number for this molecule.
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46 Give the VSEPR number for this molecule.
46 Give the VSEPR number for this molecule.
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47 Give the VSEPR number for
this molecule.
47 Give the VSEPR number for
this molecule.
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48 Which compound below contains an atom that is
surrounded by more than an octet of electrons?
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48 Which compound below contains an atom that is
surrounded by more than an octet of electrons?
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49 Which of the following molecules would have a bent
shape?
A SO2 B SO3 C CH4 D
C2H2
E
HF
49 Which of the following molecules would have a bent
shape?
A SO2 B SO3 C CH4 D
C2H2
E
HF
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50 Which of the following molecules would have a 104.5
degree bond angle between atoms?
A H2S B CF3Cl C CO2 D
PCl3
E
NO3-
50 Which of the following molecules would have a 104.5
degree bond angle between atoms?
A H2S B CF3Cl C CO2 D
PCl3
E
NO3-
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51 The molecular shape and geometry of the nitrate ion
(NO3-) would be:
A bent B linear C trigonal planar D
trigonal bipyramidal
E
tetrahedral
51 The molecular shape and geometry of the nitrate ion
(NO3-) would be:
A bent B linear C trigonal planar D
trigonal bipyramidal
E
tetrahedral
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According to carbon's orbital diagram, it should only be able to form two bonds... __ __ __ __ __ 1s 2s 2p
But we know carbon forms 4 bonds, not 2!!!
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Scientists propose that the outermost s and p orbitals are actually combined to create 4 "hybrid" orbitals of equal energy. Carbon __ ___ ___ ___ ___
1s sp3 hybrid orbitals
This explained how carbon could form 4 bonds
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To predict the hybridization involved in a compound, simply look at the first VSEPR numbers, this tells you how many electron domains(orbitals) need to be hybridized. For example: = 4 electron domains sp3 Carbon requires 4 hybrid orbital so it hybridizes it's outermost "s"
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Example: Find the hybridization of the N atom in NH3? VSEPR # = 4 so the hybridization is sp3
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Example: What is the hybridization of C in CO2? VSEPR # = 2 Only the s orbital and 1 p orbital are needed to be hybridized so the hydridization is sp Note: The other 2 p orbitals not involved in hybridization are used to form the double bonds (called Pi bonds)
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52 Which of the following would require sp2 hybridization?
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A BF3 B H2O C PCl3 D
F2
E
N2
52 Which of the following would require sp2 hybridization?
[*]
A BF3 B H2O C PCl3 D
F2
E
N2
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53 What would be the hybridization found on O in OF2?
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A sp B sp2 C sp3 D
s2p3
E
s3p3
53 What would be the hybridization found on O in OF2?
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A sp B sp2 C sp3 D
s2p3
E
s3p3
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Though atoms often form compounds by sharing electrons, the electrons are not always shared equally. In a covalent bond, one atom has a greater ability to pull the shared pair toward it.
Identical atoms will have an electronegativity difference of ZERO. As a result, the bond is NONPOLAR.
Bond Type Non-Polar Covalent Polar Covalent Ionic Electronegativity Difference very small or zero about 0.2 to 1.6 above 1.7 (between metal & non-metal)
Therefore, the fluorine end of the molecule has more electron density than the hydrogen end.
H F
We use the symbol to designate a dipole (2 poles). The "+" end is on the more positive end of the molecule and the arrow points towards the more negative end.
Compound Bond Electronegativity Dipole length (A0) Difference Moment (D) HF 0.92 1.9 1.82 HCl 1.27 0.9 1.08 HBr 1.41 0.7 0.82 HI 1.61 0.4 0.44 Bond lengths, Electronegativity, Differences and Dipole Moments of the Hydrogen Halides
But just because a molecule possesses polar bonds does not mean the molecule as a whole will be polar.
For instance, in the case of CO2: The polar bond is shown as a dipole, the arrow points to the more negative atom. Dipoles add as vectors.
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By adding the individual bond dipoles,
moment for the molecule. For a molecule to be polar, it must a) contain one or more dipoles AND b) have these polar bonds arranged asymmetrically
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In other words, if all the dipoles are symmetrical, they will cancel each
NONPOLAR. Many molecules with lone pairs of electrons will be POLAR.
These are some examples of polar & nonpolar molecules. What are their VSEPR numbers?
330, nonpolar 440, nonpolar 440, polar 431, polar 110(?), polar
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Slide for Answer Slide for Answer Slide for Answer Slide for Answer
54 Which of these are polar molecules?
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54 Which of these are polar molecules?
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55 Sulfur trioxide (SO3) is polar.
True False
55 Sulfur trioxide (SO3) is polar.
True False
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56 Hydrogen sulfide gas (H2S) is non-polar.
True False
56 Hydrogen sulfide gas (H2S) is non-polar.
True False
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57 Which of the following contains polar bonds but is a
non-polar molecule?
A CH4 B CS2 C H2S D
CF4
E
All of these are polar
57 Which of the following contains polar bonds but is a
non-polar molecule?
A CH4 B CS2 C H2S D
CF4
E
All of these are polar
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