Chapter 1 Flashcards
(154 cards)
1
Q
What is the “Constitution” of a molecule?
A
*The order in which the atoms of a molecule are connected.
*This connectivity of the molecule is commonly referred to as simply the molecules “structure”.
2
Q
What is the correct structure of Ozone, O3?
A
3
Q
What is Condensed Structure Formulae?
A
Lewis structures in which many (or all) covalent bonds and electron pairs are omitted.
4
Q
What is Bond-line formulas?
A
*Shows bonds between atoms (except those that go from C to H).
Shows atoms that are not carbon and hydrogen (heteroatoms)
5
Q
How would this molecule look in bond-line Structure?
A
6
Q
What do curved arrows mean?
A
Curved arrows are used to track the flow of electrons in chemical reactions.
7
Q
Draw the curved arrow(s) representing the electron transfers.
A
8
Q
Draw the curved arrow(s) representing the electron transfers.
A
9
Q
Are most A-B reactions reversible?
A
Yes
10
Q
What differentiates a Strong acid from a weak acid in strength?
A
Strong acid equilibrium favour products and weak acid equilibrium favour reactants.
11
Q
The stronger the acid, the ______ the Ka. The stronger the acid, the ______ the pKa.
A
The stronger the acid, the larger the Ka. The stronger the acid, the smaller the pKa.
12
Q
pH = ____
A
pH = -log [H+]
13
Q
What are the most common type of Organic Acids?
A
Carboxylic Acids (with carboxyl functional group: C==O + C–OH (-COOH)
14
Q
What is and atomic orbital?
A
*A three-dimensional region around the nucleus where an electron is most likely to be found.
*Because the Heisenberg uncertainty principle states that both the precise location and the exact momentum of an atomic particle cannot be simultaneously determined
*We can never say precisely where an electron is—we can only describe its probable location.
15
Q
What’s the shape of an s orbital.
A
s atomic orbital is a sphere with the nucleus at its center
16
Q
Is an 2s orbital closer or further to the nucleus than an 1s orbital?
A
further
17
Q
What are the two types of properties of electrons? (“behavioural”)
A
Electrons have particle-like and wave-like properties.
18
Q
What is a node? (of a 2s orbital)
A
*2s orbital has a region where the probability of finding an electron falls to zero.
*Nodes occur because electrons have both particle-like and wave-like properties. A node is a consequence of the wave-like properties of an electron.
19
Q
What’s the shape of an p orbital.
A
*p orbitals have two lobes.
*Generally, the lobes are depicted as teardrop shaped, but computer-generated representations reveal that they are shaped more like doorknobs
*(Notice that in this context, + and - indicate the phase of the orbital; they do not indicate charge.)
20
Q
Where is the node of a p orbital?
A
The node of the p orbital is a plane—called a nodal plane—that passes through the center of the nucleus, between its two lobes.
21
Q
How many degenerated p-orbital is there?
A
*Three
*Each corresponding lobe is perpendicular to another
*(This picture shows the three degenerated 2p orbitals
22
Q
What is different from 1p orbitals to 2p orbitals? What about the energies?
A
*The energy of a 2p orbital is slightly greater than that of a 2s orbital because the average location of an electron in a 2p orbital is farther away from the nucleus.
*Drawing would be bigger tear drops than 1p orbital
23
Q
Draw a 3s orbital.
A
24
Q
What is the Molecular Orbital Theory?
A
*Like an atomic orbital, which describes the volume of space around an atom’s nucleus where an electron is likely to be found, a molecular orbital describes the volume of space around a molecule where an electron is likely to be found.
*And like atomic orbitals, molecular orbitals, too, have specific sizes, shapes, and energies.
*An atomic orbital surrounds an atom.
A molecular orbital surrounds a molecule.
25
Explain the formation of a sigma bond.
Is energy released or absorbed?
*Imagine a meeting of two separate H atoms. As the 1s atomic orbital of one hydrogen atom approaches the 1s atomic orbital of the other hydrogen atom, the orbitals begin to overlap. The atoms continue to move closer, and the amount of overlap increases until the orbitals combine to form a molecular orbital. The covalent bond that is formed when the two s orbitals overlap is called a sigma (S) bond. A S bond is cylindrically symmetrical—the electrons in the bond are symmetrically distributed about an imaginary line connecting the nuclei of the two atoms joined by the bond.
*As the two orbitals begin to overlap, energy is released because the electron in each atom is attracted to its own nucleus and to the nucleus of the other atom
26
In the formation of bonds, the decreased and increased energy means...
*The attraction between the electrons of and atom to the nuclei of another causes a decrease of energy until the nuclei are so close that they start to repulse one another.
*The repulsion of the two approaching nuclei causes a large increase in energy. The minimum energy (maximum stability) is achieved when the nuclei are a particular distance apart. This distance is the bond length of the new covalent bond
*aka: Minimum energy corresponds to maximum stability.
*At the end energy is released when a covalent bond forms and breaking the bond requires precisely the same amount of energy its used to be formed
27
What is "bond dissociation energy"
*a measure of bond strength—is the
energy required to break a bond or the energy released when a bond is formed.
*Every covalent bond has a characteristic bond length and bond dissociation energy.
28
What is the conservation principle related to orbitals in molecular formations?
Orbitals are conserved, meaning the number of molecular orbitals formed must equal the number of atomic orbitals combined.
29
In the context of the formation of an H ¬ H bond, how many molecular orbitals are discussed when two atomic orbitals are combined?
When describing the formation of an H ¬ H bond, two atomic orbitals are combined, but only one molecular orbital is discussed.
30
Even though only one molecular orbital is discussed in the formation of an H ¬ H bond, where is the other molecular orbital, and why is it not discussed?
The other molecular orbital is present, but it doesn't contain any electrons. It exists due to the wave-like properties of electrons.
31
What property of electrons leads to the formation of two molecular orbitals when two atomic orbitals combine?
The wave-like properties of electrons cause two atomic orbitals to form two molecular orbitals.
32
How can two atomic orbitals combine in the context of bonding molecular orbitals, and what is the term used to describe the constructive combination of two s atomic orbitals?
Two atomic orbitals can combine in an additive (constructive) manner, similar to how two light waves or sound waves can reinforce each other. The constructive combination of two s atomic orbitals is called a s (sigma) bonding molecular orbital.
33
How can two atomic orbitals combine in a way that cancels each other, and what is this cancellation similar to in the context of light waves and sound waves?
Two atomic orbitals can combine in a destructive way, canceling each other. This cancellation is similar to the darkness resulting from the cancellation of two light waves or the silence resulting from the cancellation of two sound waves.
34
What is the term used to describe the destructive combination of two s atomic orbitals, and how is an antibonding orbital indicated in notation?
The destructive combination of two s atomic orbitals is called a σ* antibonding molecular orbital. An antibonding orbital is indicated by an asterisk (*), and chemists read it as "star." Therefore, σ* is read as "sigma star."
35
How are the σ bonding molecular orbital and the σ antibonding molecular orbital represented in a molecular orbital (MO) diagram, and what do the horizontal lines in an MO diagram signify?*
The σ bonding molecular orbital and the σ* antibonding molecular orbital are shown in the molecular orbital (MO) diagram. In an MO diagram, the energies of both the atomic orbitals and the molecular orbitals are represented as horizontal lines. The bottom line represents the lowest energy level, and the top line represents the highest energy level.
36
When two atomic orbitals overlap, how many molecular orbitals are formed, and what is the relative energy relationship between these molecular orbitals and the atomic orbitals?
When two atomic orbitals overlap, two molecular orbitals are formed—one lower in energy and one higher in energy than the atomic orbitals. The energies of these molecular orbitals are represented in relation to the energies of the atomic orbitals in the MO diagram.
37
Where are electrons in the bonding molecular orbital most likely to be found, and what is the significance of this location in terms of atomic nuclei attraction?
Electrons in the bonding molecular orbital are most likely to be found between the nuclei, where they can more easily attract both nuclei simultaneously. This increased electron density between the nuclei is what binds the atoms together.
38
Why are electrons in the antibonding molecular orbital likely to be found anywhere except between the nuclei, and what impact does this have on the exposure of positively charged nuclei to each other?
Electrons in the antibonding molecular orbital are likely to be found anywhere except between the nuclei because a nodal plane lies between the nuclei. This leaves the positively charged nuclei more exposed to each other, detracting from the formation of a bond.
39
How does the energy level of the bonding molecular orbital compare to that of the antibonding molecular orbital, and what is the relationship between electron density and stability in molecular orbitals?
The bonding molecular orbital is lower in energy and more stable than the antibonding molecular orbital. Higher electron density between the nuclei contributes to greater stability.
40
How are electrons assigned to molecular orbitals, and what principles govern their occupation in these orbitals?
Electrons are assigned to molecular orbitals using the same rules as those used for assigning electrons to atomic orbitals:
*Electrons always occupy available orbitals with the lowest energy (the aufbau principle).
*No more than two electrons can occupy a molecular orbital (the Pauli exclusion principle).
41
According to the molecular orbital diagram, why is H2+ predicted to be less stable than H2, and why is He2 predicted not to exist?
H2+ is predicted to be less stable than H2 because it has only one electron in the bonding molecular orbital. He2 is predicted not to exist because the four electrons of He2 would fill both the lower energy bonding MO and the higher energy antibonding MO, canceling the advantage gained by the electrons in the bonding MO.
42
What type of bond is formed when two p atomic orbitals overlap, and how is this overlap characterized in terms of the orientation of the orbitals?
When two p atomic orbitals overlap, a pi (π) bond is formed. This overlap is characterized by the side of one orbital overlapping the side of the other.
43
What is constructive overlap, and how does it relate to the formation of a pi bonding molecular orbital in the context of p orbitals?
Constructive overlap occurs when two in-phase p atomic orbitals overlap (e.g., blue lobes overlap with blue lobes, and green lobes overlap with green lobes). This overlap forms a pi (π) bonding molecular orbital.
44
What is destructive overlap, and how does it contribute to the formation of a π* ("pi star") antibonding molecular orbital in the context of p orbitals?*
Destructive overlap occurs when two out-of-phase p atomic orbitals overlap (e.g., blue lobes overlap with green lobes). This overlap forms a π* (pi star) antibonding molecular orbital.
45
How many nodal planes does the pi bonding molecular orbital have, and what is the significance of these nodal planes in its structure?
The pi bonding molecular orbital has one nodal plane—a nodal plane that passes through both nuclei.
46
What does the Valence-Shell Electron-Pair Repulsion (VSEPR) model aim to predict, and what is the underlying principle of this model in terms of electron density around an atom?
*The VSEPR model predicts molecular geometry based on minimizing electron repulsion between regions of electron density around an atom.
*The model is founded on the principle that atoms share electrons by overlapping their atomic orbitals, and, due to electron pair repulsion, bonding and lone-pair electrons around an atom are positioned as far apart as possible.
47
How do Lewis structures and the VSEPR model complement each other in understanding the structure of a molecule?
*Lewis structures provide a first approximation of the structure of a simple molecule by showing how atoms share electrons.
*The VSEPR model complements this by offering insight into the shape of the molecule, considering the minimization of electron repulsion and the positioning of electron pairs as far apart as possible.
48
Why do organic chemists often find the VSEPR model helpful when visualizing chemical reactions, and what aspect of chemical change does it primarily focus on?
*Because it allows them to visualize chemical reactions in terms of changes occurring in the bonds of reacting molecules.
*The model primarily focuses on molecular geometry and the arrangement of electron pairs around atoms.
49
Despite its usefulness, why is the VSEPR model considered inadequate for some molecules, and what limitation does it have in terms of molecular orbitals?
The VSEPR model is considered inadequate for some molecules because it does not allow for the consideration of antibonding molecular orbitals.
50
In this book, how does the use of both the Molecular Orbital (MO) model and the VSEPR model differ, and what factors influence the choice between these models when describing a molecule?
Both the MO and VSEPR models are used in this book. The choice between these models depends on which one provides the best description of the molecule under discussion.
(VSEPR for almost whole chap 1 [1.7-1.13])
51
How is the bonding in methane (CH4) characterized, and what are the key features of the four covalent C¬H bonds in methane?
*The bonding in methane is characterized by four covalent C¬H bonds.
*All four identical bonds have the same length (1.10 Å), and the bond angles are identical at 109.5°.
52
What compound is a methane but with two carbons? What is the type of bond in between the two carbons?
*ethane
*there is a carbon–carbon single bond between the two carbons.
53
How are bonds represented in a perspective formula, and what is the significance of solid lines, solid wedges, and hatched wedges in such a representation?
In a perspective formula, bonds in the plane of the paper are drawn as solid lines, while a bond protruding toward the viewer is represented as a solid wedge, and a bond projecting away from the viewer is drawn as a hatched wedge. Solid lines must be adjacent to each other.
54
What information does the potential map of methane provide regarding the charge distribution on carbon and hydrogen atoms, and what conclusion can be drawn from the absence of red and blue areas in the map?
The potential map of methane shows that neither carbon nor hydrogen carries much of a charge. There are no red areas representing partially negatively charged atoms, nor blue areas representing partially positively charged atoms. The absence of partially charged atoms indicates that methane is a nonpolar molecule. This absence is explained by the similar electronegativities of carbon and hydrogen, leading to the relatively equal sharing of bonding electrons.
55
Despite having only two unpaired valence electrons, how does carbon form four covalent bonds, and what explanation is provided for this phenomenon?
Carbon forms four covalent bonds despite having only two unpaired valence electrons by promoting one of the electrons in its 2s orbital into its empty 2p orbital. This promotion results in carbon having four unpaired valence electrons, allowing it to form four covalent bonds and complete its octet.
56
Despite using an s orbital and three p orbitals to form the four C ¬ H bonds in methane, why are they considered identical, and what concept explains this phenomenon?
The four C ¬ H bonds in methane are considered identical because carbon utilizes hybrid atomic orbitals. The concept that explains this phenomenon is hybridization.
57
What are hybrid orbitals, and how are they formed through the process of hybridization?
Hybrid orbitals are mixed orbitals resulting from the combination of atomic orbitals. They are formed through a process called hybridization, which involves combining atomic orbitals.
58
What is the result of combining the one s and three p orbitals of the second shell, and how are these resulting orbitals apportioned?
Combining the one s and three p orbitals of the second shell results in four equal orbitals, each being one part s and three parts p. These mixed orbitals are called sp3 orbitals.
59
What does the term "sp3" imply, and what is the composition of each sp3 orbital in terms of s and p character?
The term "sp3" implies that one s orbital and three p orbitals were mixed. Each sp3 orbital has 25% s character and 75% p character.
60
Why are the four sp3 orbitals considered degenerate, and what does degeneracy mean in this context?
The four sp3 orbitals are degenerate, meaning they all have the same energy level.
61
What does an sp3 orbital looks like?
*because of overlap, and 25% s character, 75% p character, lobe sizes look a bit different than a usual p orbital
*the s orbital adds to one lobe of the p orbital and subtracts from the other lobe
62
What is the primary role of the larger lobe of the sp3 orbital?
to form covalent bonds
63
How does the stability of an sp3 orbital compare to a p orbital and an s orbital?
The stability of an sp3 orbital reflects its composition; it is more stable (lower energy) than a p orbital but not as stable as an s orbital (higher energy).
64
What is the spatial arrangement of the four sp3 orbitals in a tetrahedral carbon structure?
The four sp3 orbitals point toward the corners of a regular tetrahedron, forming a pyramid with four faces, each being an equilateral triangle.
65
How are the four C ¬ H bonds in methane formed in relation to the sp3 orbitals of carbon and the s orbital of hydrogen?
Each of the four C ¬ H bonds in methane is formed from the overlap of an sp3 orbital of carbon with the s orbital of a hydrogen.
66
What is the bond angle in methane?
The bond angle in methane is 109.5°
67
Why is it called a tetrahedral bond angle in methane?
Because it corresponds to the angle between any two lines pointing from the center to the corners of a tetrahedron.
68
What type of carbon is found in methane? Why is it called like that?
The carbon in methane is called a tetrahedral carbon because it forms covalent bonds using four equivalent sp3 orbitals, resulting in a tetrahedral arrangement.
69
Why is hybrid orbital theory mentioned in the context of organic compounds, and what does it provide?
Hybrid orbital theory is mentioned because it offers a theoretical framework to explain the bonding in organic compounds, particularly those involving tetrahedral carbons.
*It helps us understand how atoms bond in organic molecules
70
To how many atoms is each carbon ion ethane bonded to?
Each carbon in ethane (CH3CH3) is bonded to four other atoms.
71
What is the type of bond between the each atoms of ethane?
Single bonds
72
What type of orbitals do the carbons in ethane use to bond to four atoms each?
*sp3 orbitals just like in methane
*One sp3 orbital of one carbon of ethane overlaps an sp3 orbital of the other carbon to form the C ¬ C bond.
73
Is ethane a polar molecule?
*No, the potential map shows that ethane, like methane, is a nonpolar molecule.
74
What does the MO diagram illustrate regarding the overlap of sp3 orbitals in carbon compounds?
the end-on overlap of an sp3 orbital of one carbon with an sp3 orbital of another carbon, forming a cylindrically symmetrical bond known as a sigma (σ) bond.
75
What type of bond is formed through end-on overlap in the MO diagram?
End-on overlap results in the formation of a sigma (σ) bond.
75
Are all single bonds in organic compounds sigma bonds?
Yes
76
According to Figure 1.9, where is the electron density concentrated in the σ bonding molecular orbital?
between the nuclei.
77
What is the significance of the small back lobes (nonoverlapping green lobes) in Figure 1.9?
they are a result of the concentrated electron density between the nuclei in the σ bonding MO.
78
How many bonds do the carbons in ethene (ethylene) form with each other?
The carbons in ethene form a double bond.
79
What's the other name for ethene?
ethylene
80
How many total bonds does each carbon in ethene form?
Each carbon in ethene forms a total of four bonds.
81
Despite forming four bonds, why is each carbon in ethene bonded to only three atoms? (ethene)
One of the bonds formed by each carbon is a double bond with the other carbon, and the remaining three bonds are single bonds with other atoms or groups.
82
How many atomic orbitals does each carbon hybridize to bond with three atoms in ethene?
Each carbon hybridizes three atomic orbitals to bond with three atoms in ethene. (+ one unhybridized p orbital)
83
What is the specific hybridization that results from the hybridization of an s orbital and two p orbitals? And how many hybrid orbitals are formed after sp2 hybridization?
*The hybridization resulting from the hybridization of an s orbital and two p orbitals is called sp2 hybridization.
*After sp2 hybridization, three hybrid orbitals are formed
84
How many degenerate sp2 orbitals does each carbon atom have after hybridization?
After hybridization, each carbon atom has three degenerate sp2 orbitals.
85
Why do the three sp2 orbitals in ethene lie in a plane?
To minimize electron repulsion, as they get as far from each other as possible.
86
In what direction are the axes of the three sp2 orbitals directed?
The axes of the three sp2 orbitals are directed toward the corners of an equilateral triangle.
87
Where is the carbon nucleus located in relation to the plane formed by the sp2 orbitals?
The carbon nucleus is at the center of the equilateral triangle formed by the sp2 orbitals.
88
What is the result of the orientation of the sp2 orbitals in terms of bond angles?
The result of the orientation of the sp2 orbitals is that the bond angles are all close to 120°.
89
How is the unhybridized p orbital placed in the three sp2 hybridized orbitals of ethene?
Perpendicular to the triangular plane formed by the sp2 orbitals.
90
How are the atoms in ethene placed in 3D space.
They all lie in the same plane
91
What makes the atoms in ethene lie in the same plane? (Aka to achieve maximum overlap)
*The two p orbitals that overlap to form the p bond must be parallel to each other
*This forces the triangle formed by one carbon and two hydrogens to lie in the same plane as the triangle formed by the other carbon and two hydrogens.
92
What can you tell from the potential map of ethene?
that it is a nonpolar molecule with a slight accumulation of negative charge (the pale orange area) above the two carbons.
93
Describe double bonds versus single bonds when it comes to the amount of electrons in their formation.
Single bond: two electrons
Double bonds: four electrons
94
Draw ethyne
95
What's another name for ethyne?
acetylene
96
How many bonds can an ethyne carbon make?
four
97
How many atoms can an ethyne carbon be attached to?
two
98
What does the orbitals look like around a carbon of ethyne?
*The two sp orbitals point in opposite directions.
*The two unhybridized p orbitals are perpendicular to each other and to the sp orbitals. (The smaller lobes of the sp orbitals are not shown.)
99
How many orbitals are hybridized in a carbon atom in ethyne?
two sp orbitals and two unhybridized p orbitals
100
What can you tell from the potential map of ethyne?
*Because the two unhybridized p orbitals on each carbon are perpendicular to each other, they create regions of high electron density above and below and in front of and in back of the internuclear axis of the molecule
*The overall result can be seen in the potential map for ethyne—the negative charge accumulates in a cylinder that wraps around the egg-shaped molecule.
101
What does a triple bond consist of?
A triple bond consists of
one σ bond and two π bonds.
102
By how many electrons are the carbons in ethyne held together?
6
103
What's the order in which we can determine which orbitals are used in bonds of a molecule?
1) Determine the most probable arrangement of the electrons (respecting valence electrons: type of bonds and lone pairs)
2) Determine the bond angles in 3D space
104
How many bonds does a carbon with a positive charge form in the methyl cation?
three bonds.
105
What orbitals does the positively charged carbon in the methyl cation hybridize?
it hybridizes three orbitals—an s orbital and two p orbitals.
106
What type of orbitals does the positively charged carbon use to form its three covalent bonds in the methyl cation?
sp2 orbitals
107
What remains empty in the positively charged carbon in the methyl cation?
The unhybridized p orbital
108
How do the positively charged carbon and the three atoms bonded to it are arranged in the methyl cation?
*They lie in a plane in the methyl cation.
*The unhybridized p orbital stands perpendicular to the plane.
109
How is the carbon atom in the methyl radical hybridized?
sp2
110
What is the key difference in electron count between the methyl radical and the methyl cation?
The methyl radical has one more electron than the methyl cation.
111
Where does the unpaired electron in the methyl radical reside, and how is its electron density distributed?
In the p orbital, with half of the electron density in each lobe.
112
How do the ball-and-stick models of the methyl cation and the methyl radical compare?
They look the same
113
Why do the potential maps of the methyl cation and the methyl radical appear quite different?
due to the additional electron in the methyl radical, which affects the distribution of electron density
114
How many pairs of bonding electrons and lone pairs does the negatively charged carbon in methyl anion have?
Three pairs of bonding electrons and one lone pair
115
What is the orbitals and lone pairs arrangement in the methyl anion?
Four pairs of electrons are farthest apart when the four orbitals contain- ing the bonding and lone-pair electrons point toward the corners of a tetrahedron.
116
What hybridization does the negatively charged carbon in the methyl anion have?
sp3 (tetrahedral)
117
What do the orbitals in the methyl anion do?
*three of carbon’s sp3 orbitals each overlap the s orbital of a hydrogen, and the fourth sp3 orbital holds the lone pair.
118
What can you say from the methyl anion's potential map?
Electron density is concentrated around the lone pair
119
How many covalent bonds does the nitrogen atom in ammonia form?
three covalent bonds
120
What is the electronic configuration of nitrogen in ammonia?
it has three unpaired valence electrons.
121
How many lone pairs does the nitrogen atom in ammonia have?
one
122
What is the goal of nitrogen in forming the three covalent bonds in ammonia?
to achieve an outer shell of eight electrons, completing its octet.
123
What are the hybrid orbitals used by nitrogen in ammonia?
The s orbital and three p orbitals hybridize to form four degenerate sp3 orbitals.
124
Why does nitrogen in ammonia hybridized four orbitals?
because it forms three bonds and has one lone pair.
125
From the overlap of what orbitals are the N-H bonds in ammonia formed?
of an sp3 orbital of nitrogen with the s orbital of a hydrogen.
126
Why is the tetrahedral bond angle in ammonia (107.3) a bit smaller than the "normal" bond angle (109.5)?
*Because of its lone pair
*A lone pair occupies more space around the atom than does a bonding pair that is attracted by two nuclei and relatively confined between them. Consequently, a lone pair exerts more electron repulsion, which squeezes the N¬H bonds together and decreases the bond angle.
127
What are the bond angles in ammonium? NH4+
109.5°, just like the bond angles in methane
128
How many covalent bonds does the oxygen atom in water make and how many lone pairs does it have?
two covalent bonds and two lone pairs
129
What does the electronic configuration of water shows?
That it has two unpaired valence electrons and therefore doesn't need to promote any electron to form the two covalent bonds in order to complete its octet.
130
What is the hybridization in water molecule?
like carbon and nitrogen, the one s and three p orbitals hybridize to form four degenerate sp3 orbitals
131
Why does oxygen in water hybridized four orbitals?
because it forms two bonds and has two lone pairs.
132
What are the overlaps in the orbitals of water molecules for the formation of bonds?
*Each of the two O¬H bonds is formed by the overlap of an sp3 orbital of oxygen with the s orbital of a hydrogen.
*A lone pair occupies each of the two remaining sp3 orbitals.
133
What is the bond angle in water and why is it like that?
104.5°; its even smaller than the bond angles in NH3 (107.3°) because oxygen has two relatively diffuse lone pairs, whereas nitrogen has only one.
134
What are hydrogen halides?
*HF, HCl, HBr, and HI
135
What does the electronic configuration of a halogen look like?
A halogen has only one unpaired valence electron, so it forms only one covalent bond.
136
How many bonds can hydrogen halide make? Why aren't bond angles important with hallides?
they have only one bond and, therefore, no bond angles
137
What is the hybridization in halogens valence electrons?
138
What overlap result in the formation of a hydrogen-halogen bond?
of an sp3 orbital of the halogen with the s orbital of hydrogen.
139
In the case of fluorine, to which electron shell does the sp3 orbital used in bond formation belong?
second electron shell.
140
For chlorine, to which electron shell does the sp3 orbital belong?
third electron shell.
141
Why is the average electron density less in a 3sp3 orbital than in a 2sp3 orbital?
because the average distance from the nucleus is greater for an electron in the third shell than for an electron in the second shell.
142
What effect does the decrease in electron density as the electron shells gets bigger have on the hydrogen–halogen bond?
The decrease in electron density results in the hydrogen–halogen bond becoming longer and weaker.
143
How does the size (atomic weight) of the halogen impact the hydrogen–halogen bond?
As the size (atomic weight) of the halogen increases, the hydrogen–halogen bond becomes longer and weaker.
144
Describe the orbitals used in bonding and the bond angles in BeH2.
*Beryllium (Be) does not have any unpaired valence electrons.
*Therefore, it cannot form any bonds unless it promotes an electron. After promoting an electron from the s orbital to an empty p orbital and hybridizing the two orbitals that now each contain an unpaired electron, two sp orbitals result.
145
What is bond order?
*the number of covalent bonds shared by two atoms.
*A single bond has a bond order of one, a double bond has a bond order of two, and a triple bond has a bond order of three.
146
What is the easiest way to determine the hybridization of carbon, nitrogen, or oxygen? What are exceptions to these rules?
* it's to count the number of p bonds it forms:
■ If it forms no p bonds, it is sp3 hybridized.
■ If it forms one p bond, it is sp2 hybridized.
■ If it forms two p bonds, it is sp hybridized.
*carbocations and carbon radicals, which are sp2 hybridized—not because they form a p bond, but because they have an empty or a half-filled p orbital
147
The greater the electron density in the region of overlap,
the _____ the bond.
stronger
148
Why is a C ¬ H σ bond shorter than a C ¬ C σ bond?
*because the s orbital of hydrogen is closer to the nucleus than is the sp3 orbital of carbon.
*Consequently, the nuclei are closer together in a bond formed by sp3–s overlap than they are in a bond formed by sp3–sp3 overlap.
149
Why do both length and strength of a C-H bond depend on the hybridization of the carbon?
*The more s character in the orbital used by carbon to form the bond, the shorter and stronger the bond— again because an s orbital is closer to the nucleus than is a p orbital.
*Thus, a C¬H bond formed by an sp carbon (50% s) is shorter and stronger than a C¬H bond formed by an sp2 carbon (33.3% s), which, in turn, is shorter and stronger than a C¬H bond formed by an sp3 carbon (25% s).
150
*The more s character in the orbital, the ________ the bond.
*The more s character in the orbital, the _________ the bond.
*The more s character in the orbital, the ________ the bond angle.
*A pi bond is _______ than a sigma bond.
*shorter,stronger,larger
*weaker
151
Draw the polarity arrows:
152
Why is the dipole moment stronger in water than in a single O-H bond? And similarly in a ammonia versus a single N-H bond?
because the dipoles of the two O ¬ H bonds reinforce each other; the lone-pair electrons also contribute to the dipole moment.
153
What is the formula for Formal Charge?
FC = #Ve - (lines+dots)
