Chapter 12 Chemical Bonding by Christopher Hamaker

Chapter 12 Chemical Bonding by Christopher Hamaker
© 2011 Pearson Education, Inc. Chapter 12

Chemical Bonding Chemical Bond - the force of attraction between any two atoms in a compound. Interactions involving valence electrons are responsible for the chemical bond.

Chemical Bond Concept Recall that an atom has core and valence electrons. Core electrons are found close to the nucleus. Valence electrons are found in the most distant s and p energy subshells. It is valence electrons that are responsible for holding two or more atoms together in a chemical bond. Chapter 12

Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that participate in chemical bonding. Group # of valence e- e- configuration 1A 1 ns1 2A 2 ns2 3A 3 ns2np1 4A 4 ns2np2 5A 5 ns2np3 6A 6 ns2np4 7A 7 ns2np5

Lewis symbol (Lewis structure) - a way to represent atoms (and their bonds) using the element symbol and valence electrons as dots.

Octet Rule The octet rule states that atoms bond in such a way that each atom acquires eight electrons in its outer shell. There are two ways in which an atom may achieve an octet: Transfer of electrons from one atom to another Sharing one or more pairs of electrons Chapter 12

Types of Bonds Ionic bonds are formed when a complete transfer of electrons between atoms occurs, forming ionic compounds. Covalent bonds are formed when two atoms share electrons to form molecular compounds. Chapter 12

Describing Ionic Bonds
An ionic bond is a chemical bond formed by the electrostatic attraction between positive and negative ions. This type of bond involves the transfer of electrons from one atom (usually a metal) to another (usually a nonmetal). The number of electrons lost or gained by an atom is determined by its need to be “isoelectronic” with a noble gas. 2

Ionic Bonds An ionic bond is formed by the attraction between positively charged anions and negatively charged anions. This electrostatic attraction is similar to the attraction between opposite poles on two magnets. Chapter 12

Ionic Bonds, Continued The ionic bonds formed from the combination of anions and cations are very strong and result in the formation of a rigid, crystalline structure. The structure for NaCl, ordinary table salt, is shown here. Chapter 12

Let’s examine the formation of NaCl Na + Cl  NaCl
IONIC BONDING Let’s examine the formation of NaCl Na + Cl  NaCl Chlorine has a high electron affinity. When chlorine gains an electron, it gains the Ar configuration. Sodium has a low ionization energy it readily loses this electron . When Sodium loses the electron, it gains the Ne configuration. Na  Na+ + e-

Essential Features of Ionic Bonding
Atoms with low I.E. and low E.A. tend to form positive ions. Atoms with high I.E. and high E.A. tend to form negative ions. Ion formation takes place by electron transfer. The ions are held together by the electrostatic force of the opposite charges. Reactions between metals and nonmetals (representative) tend to be ionic.

Formation of Cations Cations are formed when an atom loses valence electrons to become positively charged. Most main group metals achieve a noble gas electron configuration by losing their valence electrons, and are isoelectronic with a noble gas. Magnesium (Group IIA/2) loses its two valence electrons to become Mg2+. A magnesium ion has 10 electrons (12 – 2 = 10 e-) and is isoelectronic with neon. Chapter 12

Formation of Cations, Continued
We can use electron dot formulas to look at the formation of cations. Each of the metals in Period 3 forms cations by losing one, two, or three electrons, respectively. Each metal atom becomes isoelectronic with the preceding noble gas, neon. Chapter 12

Such noble gas configurations and the corresponding ions are particularly stable. The atom that loses the electron becomes a cation (positive). The atom that gains the electron becomes an anion (negative). 2

Electron Configurations of Ions
As metals lose electrons to form cations and establish a “noble gas” configuration, the electrons are lost from the valence shell first. For example, magnesium generally loses two electrons from its 3s subshell to look like neon. [Ne]3s2 [Ne] 2

Electron Configurations of Ions
Transition metals also lose electrons from the valence shell first, which is not the last subshell to fill according to the aufbau sequence. For example, zinc generally loses two electrons from its 4s subshell to adopt a “pseudo”-noble gas configuration. [Ar]4s23d10 [Ar]3d10 2

Formation of Anions Anions are formed when an atom gains electrons and becomes negatively charged. Most nonmetals achieve a noble gas electron configuration by gaining electrons to become isoelectronic with a noble gas. Chlorine (Group VIIA/17) gains one valence electron and becomes Cl–. A chloride ion has 18 electrons ( = 18 e-) and is isoelectronic with argon. Chapter 12

Formation of Anions, Continued
We can also use electron dot formulas to look at the formation of anions. The nonmetals in Period 3 gain one, two, or three electrons, respectively, to form anions. Each nonmetal ion is isoelectronic with the following noble gas, argon. Chapter 12

Consider the transfer of valence electrons from a sodium atom to a chlorine atom. e- The resulting ions are electrostatically attracted to one another. The attraction of these oppositely charged ions for one another is the ionic bond. 2

Ionic Radii The radius of a cation is smaller than the radius of its starting atom. The radius of an anion is larger than the radius of its starting atom. Chapter 12

Predicting Formulas Predict the formula of the ionic compounds formed from the following combination of ions: 1. sodium and sulfur 2. magnesium and oxygen 3. aluminum and oxygen 4. barium and fluorine

Covalent Bonds When two nonmetals bond, they often share electrons since they have similar attractions for them. This sharing of valence electrons is called the covalent bond. These atoms will share sufficient numbers of electrons in order to achieve a noble gas electron configuration (that is, eight valence electrons). 2

Covalent Bonds Covalent bonds are formed when two nonmetal atoms share electrons, and the shared electrons in the covalent bond belong to both atoms. When hydrogen chloride (HCl) is formed, the hydrogen atom shares its one valence electron with the chlorine. This gives the chlorine atom eight electrons in its valence shell, making it isoelectronic with argon. The chlorine atom shares one of its valence electrons with the hydrogen, giving it two electrons in its valence shell, and making it isoelectronic with helium. Chapter 12

not a part of a massive three dimensional crystal structure.
Covalent Compounds Covalent compounds are usually formed from nonmetals. Molecules - compounds characterized by covalent bonding. not a part of a massive three dimensional crystal structure.

Instead, each atom gets a noble gas configuration by sharing electrons.
Each Hydrogen atom now has two electrons around it and has a He configuration The shared electrons pair is a Covalent Bond

Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? 7e- 7e- 8e- 8e- F F + F Lewis structure of F2 lone pairs F single covalent bond single covalent bond F

Bond Length When a covalent bond is formed, the valence shells of the two atoms overlap with each other. In HCl, the 1s energy sublevel of the hydrogen atom overlaps with the 3p energy sublevel of the chlorine atom. The mixing of sublevels draws the atoms closer together. The distance between the two atoms is smaller than the sum of their atomic radii and is the bond length. Chapter 12

Lengths of Covalent Bonds
Bond Lengths Triple bond < Double Bond < Single Bond

Bond Energy Energy is released when two atoms form a covalent bond.
H(g) + Cl(g)  HCl(g) + heat Conversely, energy is needed to break a covalent bond. The energy required to break a covalent bond is referred to as the bond energy. The amount of energy required to break a covalent bond is the same as the amount of energy released when the bond is formed. HCl(g) + heat  H(g) + Cl(g) Chapter 12

Bond energy - the amount of energy required to break a bond holding two atoms together.
triple bond > double bond > single bond Bond length - the distance separating the nuclei of two adjacent atoms. single bond > double bond > triple bond

Electron Dot Formulas of Molecules
In Section 6.8, we drew electron dot formulas for atoms. The number of dots around each atom is equal to the number of valence electrons the atom has. We will now draw electron dot formulas for molecules (also called Lewis structures). A Lewis structure shows the bonds between atoms and helps us visualize the arrangement of atoms in a molecule. Chapter 12

Lewis Structures You can represent the formation of the covalent bond in H2 as follows: H . : + This uses the Lewis dot symbols for the hydrogen atom and represents the covalent bond by a pair of dots. 2

Lewis Structures The shared electrons in H2 spend part of the time in the region around each atom. : H In this sense, each atom in H2 has a helium configuration. 2

. : : + H Cl H Cl Lewis Structures
The formation of a bond between H and Cl to give an HCl molecule can be represented in a similar way. H . : Cl + : H Cl Thus, hydrogen has two valence electrons about it (as in He) and Cl has eight valence electrons about it (as in Ar). 2

Lewis Structures Formulas such as these are referred to as Lewis electron-dot formulas or Lewis structures. An electron pair is either a bonding pair (shared between two atoms) or a lone pair (an electron pair that is not shared). bonding pair lone pair : H Cl 2

Guidelines for Electron Dot Formulas
Calculate the total number of valence electrons by adding all of the valence electrons for each atom in the molecule. Divide the total valence electrons by 2 to find the number of electron pairs in the molecule. Surround the central atom with four electron pairs. Use the remaining electron pairs to complete the octet around the other atoms. The only exception is hydrogen, which only needs two electrons. Chapter 12

Guidelines for Electron Dot Formulas, Continued
Electron pairs that are shared by atoms are called bonding electrons. The other electrons complete octets and are called nonbonding electrons, or lone pairs. If there are not enough electron pairs to provide each atom with an octet, move a nonbonding electron pair between two atoms that already share an electron pair. Chapter 12

Electron Dot Formula for H2O
Count the total number of valence electrons: oxygen has six and each hydrogen has one for a total of eight electrons [6 + 2(1) = 8 e-]. The number of electron pairs is 4 (8/2 = 4). Place eight electrons around the central oxygen atom. We can then place the two hydrogen atoms in any of the four electron pair positions. Notice there are two bonding and two nonbonding electron pairs. Chapter 12

Electron Dot Formula for H2O, Continued
4. To simplify, we represent bonding electron pairs with a single dash line called a single bond. 5. The resulting structure is referred to as the structural formula of the molecule. Chapter 12

Electron Dot Formula for SO3
Count the total number of valence electrons: each oxygen has six and sulfur has six for a total of 24 electrons [3(6) + 6 = 24 e-]. This gives us 12 electron pairs. Place four electron pairs around the central sulfur atom and attach the three oxygens. We started with 12 electron pairs and have eight left. Place the remaining electron pairs around the oxygen atoms to complete each octet. One of the oxygens does not have an octet, so move a nonbonding pair from the sulfur to provide two pairs between the sulfur and the oxygen. Chapter 12

Multiple Bonds In the molecules described so far, each of the bonds has been a single bond, that is, a covalent bond in which a single pair of electrons is shared. It is possible to share more than one pair. A double bond involves the sharing of two pairs between atoms. or C : H 2

Multiple Bonds Triple bonds are covalent bonds in which three pairs of electrons are shared between atoms. C or H : ::: 2

Writing Lewis Dot Formulas
The following rules allow you to write electron-dot formulas even when the central atom does not follow the octet rule. To illustrate, we will draw the structure of PCl3, phosphorus trichloride. 2

Step 1: Total all valence electrons in the molecular formula. That is, total the group numbers of all the atoms in the formula. 5 e- (7 e-) x 3 26 e- total For a polyatomic anion, add the number of negative charges to this total. For a polyatomic cation, subtract the number of positive charges from this total. 2

Step 2: Arrange the atoms radially, with the least electronegative atom in the center. Place one pair of electrons between the central atom and each peripheral atom. Cl Cl P Cl 2

Step 3: Distribute the remaining electrons to the peripheral atoms to satisfy the octet rule. : Cl Cl P Cl 2

Step 4: Distribute any remaining electrons to the central atom. If there are fewer than eight electrons on the central atom, a multiple bond may be necessary. : : : : : : Cl Cl : P : : Cl : 2

Try drawing Lewis dot formulas for the following covalent compound. SCl2 20 e- total Cl S 16 e- left : 4 e- left : 2

Try drawing Lewis dot formulas for the following covalent compound. COCl2 24 e- total 18 e- left : 0 e- left Cl C O 2

Note that the carbon has only 6 electrons. One of the oxygens must share a lone pair. COCl2 24 e- total 18 e- left : 0 e- left Cl C O 2

Note that the carbon has only 6 electrons. One of the oxygens must share a lone pair. COCl2 24 e- total 18 e- left : O : 0 e- left C Note that the octet rule is now obeyed. : : Cl Cl 2

Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – N is less electronegative than F, put N in center Step 2 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5) 5 + (3 x 7) = 26 valence electrons Step 3 – Draw single bonds between N and F atoms and complete octets on N and F atoms. Step 4 - Check, are # of e- in structure equal to number of valence e- ? 3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons F N

Resonance The two shared electron pairs constitute a double bond.
The double bond can be placed between the sulfur and any of the three oxygen atoms. The structural formula can be shown as any of the structures below. This phenomenon is called resonance. Chapter 12

A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + - What are the resonance structures of the carbonate (CO32-) ion? O C - O C - O C -

Delocalized Bonding: Resonance
According to theory, one pair of bonding electrons is spread over the region of all three atoms. O This is called delocalized bonding, in which a bonding pair of electrons is spread over a number of atoms. 2

Exceptions to the Octet Rule
Although many molecules obey the octet rule, there are exceptions where the central atom has more than eight electrons. Generally, if a nonmetal is in the third period or greater it can accommodate as many as twelve electrons, if it is the central atom. These elements have unfilled “d” subshells that can be used for bonding. 2

Let’s look at BF3 Let’s look at NO
Lewis Structures and Exceptions to the Octet Rule 1. Incomplete Octet - less then eight electrons around an atom other than H. Let’s look at BF3 2. Odd Electron - if there is an odd number of valence electrons it isn’t possible to give every atom eight electrons. Let’s look at NO 3. Expanded Octet - elements in 3rd period and below may have 10 and 12 electrons around it.

The Incomplete Octet Be – 2e- 2H – 2x1e- 4e- BeH2 H Be B – 3e- 3F – 3x7e- 24e- 3 single bonds (3x2) = 6 9 lone pairs (9x2) = 18 Total = 24 F B BF3

Odd-Electron Molecules N – 5e- O – 6e- 11e- NO N O The Expanded Octet (central atom with principal quantum number n > 2) S F S – 6e- 6F – 42e- 48e- 6 single bonds (6x2) = 12 18 lone pairs (18x2) = 36 Total = 48 SF6

For example, the bonding in phosphorus pentafluoride, PF5, shows ten electrons surrounding the phosphorus. : F : : F : : : F : P F : : : : F : 2

In xenon tetrafluoride, XeF4, the xenon atom must accommodate two extra lone pairs. : F : F : : : Xe : F : F : : : 2

Write the Lewis structures of SH6 and SF4
Expanded octet is the most common exception. Write the Lewis structures of SH6 and SF4

Electron Dot Formula for NH4+
The total number of valence electrons is 5 – 4(1) – 1 = 8 e-. We must subtract one electron for the positive charge. We have four pairs of electrons. Place four electron pairs around the central nitrogen atom and attach the four hydrogens. Enclose the polyatomic ion in brackets and indicate the charge outside the brackets. Chapter 12

Electron Dot Formula for CO32-
The total number of valence electrons is 4 + 3(6) + 2 = 24 e-. We must add one electron for the negative charge. We have 12 pairs of electrons. Place four electron pairs around the central carbon atom and attach the three oxygens. Use the remaining electron pairs to give the oxygen atoms their octets. One oxygen does not have an octet. Make a double bond and enclose the ion in brackets. Chapter 12

Let's Practive Lewis Structures
Using the guidelines presented, write Lewis structures for the following: 1. H2O 2. NH3 3. CO2 4. NH4+ 5. CO32- 6. N2

Example Video dot structures I: single bonds (6:57 min)
dot structures II: multiple bonds (5:59 min) Resonance (11:51 min) © 2011 Pearson Education, Inc. Chapter 4

Polar Covalent Bonds A polar covalent bond is one in which the bonding electrons spend more time near one of the two atoms involved. When the atoms are alike, as in the H-H bond of H2 , the bonding electrons are shared equally (a nonpolar covalent bond). When the two atoms are of different elements, the bonding electrons need not be shared equally, resulting in a “polar” bond. 2

Polar Covalent Bonding and Electronegativity
1 The Polar Covalent Bond Ionic bonding involves the transfer of electrons. Covalent bonding involves the sharing of electrons. Polar covalent bonding - bonds made up of unequally shared electron pairs.

Polar Covalent Bonds Covalent bonds result from the sharing of valence electrons. Often, the two atoms do not share the electrons equally. That is, one of the atoms holds onto the electrons more tightly than the other. When one of the atoms holds the shared electrons more tightly, the bond is polarized. A polar covalent bond is one in which the electrons are not shared equally. Chapter 12

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F electron rich region electron poor region e- poor e- rich F H d+ d-

Electronegativity Each element has an innate ability to attract valence electrons. Electronegativity is the ability of an atom to attract electrons in a chemical bond. Linus Pauling devised a method for measuring the electronegativity of each of the elements. Fluorine is the most electronegative element. Chapter 12

Polar Covalent Bonds Electronegativity is a measure of the ability of an atom in a molecule to draw bonding electrons to itself. In general, electronegativity increases from the lower-left corner to the upper-right corner of the periodic table. The current electronegativity scale, developed by Linus Pauling, assigns a value of 4.0 to fluorine and a value of 0.7 to cesium. 2

Electronegativity, Continued
Electronegativity increases as you go left to right across a period. Electronegativity increases as you go from bottom to top in a family. Chapter 12

The Electronegativities of Common Elements

Electronegativity Differences
The electronegativity of H is 2.1; Cl is 3.0. Since there is a difference in electronegativity between the two elements (3.0 – 2.1 = 0.9), the bond in H – Cl is polar. Since Cl is more electronegative, the bonding electrons are attracted toward the Cl atom and away from the H atom. This will give the Cl atom a slightly negative charge and the H atom a slightly positive charge. Chapter 12

Polar Covalent Bonds The absolute value of the difference in electronegativity of two bonded atoms gives a rough measure of the polarity of the bond. When this difference is small (less than 0.5), the bond is nonpolar. When this difference is large (greater than 0.5), the bond is considered polar. If the difference exceeds approximately 1.8, sharing of electrons is no longer possible and the bond becomes ionic. 2

Classification of bonds by difference in electronegativity
Bond Type Covalent  2 Ionic 0 < and <2 Polar Covalent Increasing difference in electronegativity Covalent share e- Polar Covalent partial transfer of e- Ionic transfer e-

The greater the difference in electronegativity between two atoms, the greater the polarity of a bond. Which would be more polar, a H-F bond or a H-Cl bond? H-F … = 1.9 H-Cl… = 0.9 Therefore, the HF bond is more polar than the HCl bond.

Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and the NN bond in H2NNH2. Cs – 0.7 Cl – 3.0 3.0 – 0.7 = 2.3 Ionic H – 2.1 S – 2.5 2.5 – 2.1 = 0.4 Polar Covalent N – 3.0 N – 3.0 3.0 – 3.0 = 0 Covalent

Delta (δ) Notation for Polar Bonds
We use the Greek letter delta, d, to indicate a polar bond. The negatively charged atom is indicated by the symbol d–, and the positively charged atom is indicated by the symbol d+. This is referred to as delta notation for polar bonds. d+ H – Cl d– Chapter 12

Delta (δ) Notation for Polar Bonds, Continued
The hydrogen halides HF, HCl, HBr, and HI all have polar covalent bonds. The halides are all more electronegative than hydrogen and are designated with a d–. Chapter 12

The electrons spend more time with fluorine. This sets up a polar bond
somewhat positively charged somewhat negatively charged These two electrons are not shared equally. The electrons spend more time with fluorine. This sets up a polar bond A truly covalent bond can only occur when both atoms are identical. Electronegativity is used to determine if a bond is polar and who gets the electrons the most.

Example Video Electronegativity (11:38 min)
molecular polarity (13:46 min) © 2011 Pearson Education, Inc. Chapter 4

Nonpolar Covalent Bonds
What if the two atoms in a covalent bond have the same or similar electronegativities? The bond is not polarized and it is a nonpolar covalent bond. If the electronegativity difference is less than 0.5, it is usually considered a nonpolar bond. The diatomic halogen molecules have nonpolar covalent bonds. Chapter 12

H2, N2, O2, F2, Cl2, Br2, I2 Features of Covalent Bonds
Covalent bonds tend to form between atoms with similar tendency to gain or lose electrons. The diatomic elements have totally covalent bonds (totally equal sharing.) H2, N2, O2, F2, Cl2, Br2, I2 Each fluorine has eight electrons around it. Ne’s configuration.

Coordinate Covalent Bonds
When bonds form between atoms that both donate an electron, you have: A . : B + It is, however, possible that both electrons are donated by one of the atoms. This is called a coordinate covalent bond. A : B + 2

Coordinate Covalent Bonds
A covalent bond resulting from one atom donating a lone pair of electrons to another atom is called a coordinate covalent bond. A good example of a molecule with a coordinate covalent bond is ozone, O3. Chapter 12

Hydrogen Bonding H N O F :
Hydrogen bonding is a force that exists between a hydrogen atom covalently bonded to a very electronegative atom, X, and a lone pair of electrons on a very electronegative atom, Y. To exhibit hydrogen bonding, one of the following three structures must be present. H N O F : Only N, O, and F are electronegative enough to leave the hydrogen nucleus exposed. 2

Hydrogen Bonds The bond between H and O in water is very polar.
Therefore, the oxygen is partially negative, and the hydrogens are partially positive. As a result, the hydrogen atom on one molecule is attracted to the oxygen atom on another. This intermolecular attraction is referred to as a hydrogen bond. Chapter 12

Examples of hydrogen bonding:
H2O NH3 HF

Intermolecular Forces
Hydrogen Bond The hydrogen bond is a special dipole-dipole interaction between they hydrogen atom in a polar N-H, O-H, or F-H bond and an electronegative O, N, or F atom. A H … B or A & B are N, O, or F

Hydrogen Bonding Molecules exhibiting hydrogen bonding have abnormally high boiling points compared to molecules with similar van der Waals forces. For example, water has the highest boiling point of the Group VI hydrides. Similar trends are seen in the Group V and VII hydrides. 2

Why is the hydrogen bond considered a “special” dipole-dipole interaction?
Decreasing molar mass Decreasing boiling point

Example Video Hydrogen Bonding (1:40 min)

Shapes of Molecules Electron pairs surrounding an atom repel each other. This is referred to as Valence Shell Electron Pair Repulsion (VSEPR) theory. The electron pair geometry indicates the arrangement of bonding and nonbonding electron pairs around the central atom. The molecular shape indicates the arrangement of atoms around the central atom as a result of electron repulsion. Chapter 12

Predicting Molecular Geometry
The following rules and figures will help discern electron pair arrangements. Draw the Lewis structure Determine how many electrons pairs are around the central atom. Count a multiple bond as one pair. Arrange the electrons pairs. 2

Arrangement of Electron Pairs About an Atom
Linear 3 pairs Trigonal planar 4 pairs Tetrahedral 5 pairs Trigonal bipyramidal 6 pairs Octahedral

Tetrahedral Molecules
Methane, CH4, has four pairs of bonding electrons around the central carbon atom. The four bonding pairs (and, therefore, atoms) are repelled to the four corners of a tetrahedron. Thus, the electron pair geometry is tetrahedral. The molecular shape is also tetrahedral. Chapter 12

Trigonal Pyramidal Molecules
In ammonia, NH3, the central nitrogen atom is surrounded by three bonding pairs and one nonbonding pair. Thus, the electron pair geometry is tetrahedral and the molecular shape is trigonal pyramidal. Chapter 12

Bent Molecules In water, H2O, the central O atom is surrounded by two nonbonding pairs and two bonding pairs. Thus, the electron pair geometry is tetrahedral and the molecular shape is bent. Chapter 12

Linear Molecules In carbon dioxide, CO2, the central C atom is bonded to each oxygen by two electron pairs (a double bond). According to VSEPR theory, the electron pairs will repel each other and they will be at opposite sides of the C atom. Thus, the electron pair geometry and the molecular shape are both linear. Chapter 12

0 lone pairs on central atom 2 atoms bonded to central atom
Cl Be 2 atoms bonded to central atom

Summary of VSEPR Theory
Chapter 12

Draw Lewis structure for molecule. Count number of lone pairs on the central atom and number of atoms bonded to the central atom. Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO2 and SF4? S F S O AB4E AB2E distorted tetrahedron bent

Two electron pairs (linear arrangement). : You have two double bonds, or two electron groups about the carbon atom. Thus, according to the VSEPR model, the bonds are arranged linearly, and the molecular shape of carbon dioxide is linear. Bond angle is 180o. 2

Three electron pairs (trigonal planar arrangement). Cl C : O The three groups of electron pairs are arranged in a trigonal plane. Thus, the molecular shape of COCl2 is trigonal planar. Bond angle is 120o. 2

Three electron pairs (trigonal planar arrangement). O : Ozone has three electron groups about the central oxygen. One group is a lone pair. These groups have a trigonal planar arrangement. 2

Three electron pairs (trigonal planar arrangement). O : Since one of the groups is a lone pair, the molecular geometry is described as bent or angular. 2

Three electron pairs (trigonal planar arrangement). O : Note that the electron pair arrangement includes the lone pairs, but the molecular geometry refers to the spatial arrangement of just the atoms. 2

Four electron pairs (tetrahedral arrangement). : :Cl: H N : : O H : : :Cl C Cl: : : :Cl: : Four electron pairs about the central atom lead to three different molecular geometries. 2

Four electron pairs (tetrahedral arrangement). : :Cl: H N : : O H C :Cl : : Cl: :Cl: : tetrahedral 2

Four electron pairs (tetrahedral arrangement). : :Cl: H N : : O H C :Cl : : Cl: :Cl: : tetrahedral trigonal pyramid 2

Four electron pairs (tetrahedral arrangement). : :Cl: : H N : C O :Cl : : : Cl: H :Cl: H : tetrahedral trigonal pyramid bent 2

Five electron pairs (trigonal bipyramidal arrangement). : F : : F : : F P This structure results in both 90o and 120o bond angles. 2

Other molecular geometries are possible when one or more of the electron pairs is a lone pair. SF4 ClF3 XeF2 Let’s try their Lewis structures. 2

Other molecular geometries are possible when one or more of the electron pairs is a lone pair. F ClF3 XeF2 F : S F F see-saw 2

Other molecular geometries are possible when one or more of the electron pairs is a lone pair. S F : Cl F : XeF2 see-saw T-shape 2

Other molecular geometries are possible when one or more of the electron pairs is a lone pair. S F : Cl F : F : Xe : : F see-saw T-shape linear 2

Six electron pairs (octahedral arrangement). :F: : :F : F: : S : F: : :F :F: : This octahedral arrangement results in 90o bond angles. 2

Six electron pairs (octahedral arrangement). IF5 XeF4 Six electron pairs also lead to other molecular geometries. 2

Six electron pairs (octahedral arrangement). F F F XeF4 I F F : square pyramid 2

Six electron pairs (octahedral arrangement). : I F : F F Xe F F : square pyramid square planar 2

Nonpolar Molecules with Polar Bonds
CCl4 has polar bonds, but the overall molecule is nonpolar. Using VSEPR theory, the four chlorine atoms are at the four corners of a tetrahedron. The chlorines are each δ–, while the carbon is δ+. The net effect of the polar bonds is zero, so the molecule is nonpolar. Chapter 12

Diamond Versus Graphite
Why is diamond colorless and hard, while graphite is black and soft when both are pure carbon? Diamond has a three-dimensional structure, whereas graphite has a two-dimensional structure. The layers in graphite are able to slide past each other easily. Chapter 12

Example Video VSEPR Theory: Introduction (20:30 min)

Chapter Summary Chemical bonds hold atoms together in molecules.
Atoms bond in such a way as to have eight electrons in their valence shell. This is called the octet rule. There are two types of bonds: ionic and covalent. Ionic bonds are formed between a cation and an anion. Covalent bonds are formed from the sharing of electrons. Chapter 12

Chapter Summary, Continued
Electron dot formulas help us visualize the arrangements of atoms in a molecule. Electrons are shared unequally in polar covalent bonds. Electronegativity is a measure of the ability of an atom to attract electrons in a chemical bond. Electronegativity increases from left to right and from bottom to top on the periodic table. Chapter 12

Chapter Summary, Continued
VSEPR theory can be used to predict the shapes of molecules. The electron pair geometry indicates the arrangement of bonding and nonbonding pairs around a central atom. The molecular shape indicates the arrangement of atoms in a molecule. Chapter 12

Chapter 12 Chemical Bonding by Christopher Hamaker

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