Chemical Bonding PDF

10: CHEMICAL BONDING  CHAPTER OVERVIEW 10: Chemical Bonding How do atoms make compounds? Typically they join together in such a way that they lose their identities as elements and adopt a new identity as a compound. These joins are called chemical bonds. But how do atoms join together? Ultimately,...

10: CHEMICAL BONDING  CHAPTER OVERVIEW 10: Chemical Bonding How do atoms make compounds? Typically they join together in such a way that they lose their identities as elements and adopt a new identity as a compound. These joins are called chemical bonds. But how do atoms join together? Ultimately, it all comes down to electrons. Before we discuss how electrons interact, we need to introduce a tool to simply illustrate electrons in an atom. 10.1: Bonding Models and AIDs Drugs 10.2: Representing Valence Electrons with Dots 10.3: Lewis Structures of Ionic Compounds- Electrons Transferred 10.4: Covalent Lewis Structures- Electrons Shared 10.5: Writing Lewis Structures for Covalent Compounds 10.6: Resonance - Equivalent Lewis Structures for the Same Molecule 10.7: Predicting the Shapes of Molecules 10.8: Electronegativity and Polarity - Why Oil and Water Do not Mix 10: Chemical Bonding is shared under a CC BY-NC-SA 4.0 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 1 10.1: Bonding Models and AIDs Drugs 10.1: Bonding Models and AIDs Drugs is shared under a not declared license and was authored, remixed, and/or curated by LibreTexts. 10.1.1 https://chem.libretexts.org/@go/page/396621 10.2: REPRESENTING VALENCE ELECTRONS WITH DOTS  LEARNING OBJECTIVE Draw a Lewis electron dot diagram for an atom or a monatomic ion. In almost all cases, chemical bonds are formed by interactions of valence electrons in atoms. To facilitate our understanding of how valence electrons interact, a simple way of representing those valence electrons would be useful. A Lewis electron dot diagram (or electron dot diagram, or a Lewis diagram, or a Lewis structure) is a representation of the valence electrons of an atom that uses dots around the symbol of the element. The number of dots equals the number of valence electrons in the atom. These dots are arranged to the right and left and above and below the symbol, with no more than two dots on a side. (The order in which the positions are used does not matter.) For example, the Lewis electron dot diagram for hydrogen is simply H⋅ Because the side is not important, the Lewis electron dot diagram could also be drawn as follows: ˙ H or ⋅H or H. The electron dot diagram for helium, with two valence electrons, is as follows: He: By putting the two electrons together on the same side, we emphasize the fact that these two electrons are both in the 1s subshell; this is the common convention we will adopt, although there will be exceptions later. The next atom, lithium, has an electron configuration of 1s22s1, so it has only one electron in its valence shell. Its electron dot diagram resembles that of hydrogen, except the symbol for lithium is used: Li⋅ Beryllium has two valence electrons in its 2s shell, so its electron dot diagram is like that of helium: Be: The next atom is boron. Its valence electron shell is 2s22p1, so it has three valence electrons. The third electron will go on another side of the symbol: ˙ B: Again, it does not matter on which sides of the symbol the electron dots are positioned. For carbon, there are four valence electrons, two in the 2s subshell and two in the 2p subshell. As usual, we will draw two dots together on one side, to represent the 2s electrons. However, conventionally, we draw the dots for the two p electrons on different sides. As such, the electron dot diagram for carbon is as follows: ˙ ⋅C: With N, which has three p electrons, we put a single dot on each of the three remaining sides: ˙ ⋅N:. For oxygen, which has four p electrons, we now have to start doubling up on the dots on one other side of the symbol. When doubling up electrons, make sure that each side has no more than two electrons. ¨ ⋅O:. Fluorine and neon have seven and eight dots, respectively: ¨: :F. ¨ :Ne:.. With the next element, sodium, the process starts over with a single electron because sodium has a single electron in its highest-numbered shell, the n = 3 shell. By going through the periodic table, we see that the Lewis electron dot diagrams of atoms will never have more than 10.2.1 https://chem.libretexts.org/@go/page/47524 eight dots around the atomic symbol.  EXAMPLE 10.2.1: LEWIS DOT DIAGRAMS What is the Lewis electron dot diagram for each element? a. aluminum b. selenium Solution a. The valence electron configuration for aluminum is 3s23p1. So it would have three dots around the symbol for aluminum, two of them paired to represent the 3s electrons: ˙ Al : 2. The valence electron configuration for selenium is 4s24p4. In the highest-numbered shell, the n = 4 shell, there are six electrons. Its electron dot diagram is as follows: ˙ ⋅S e:..  EXERCISE 10.2.1 What is the Lewis electron dot diagram for each element? a. phosphorus b. argon Answer a ˙ ⋅P:. Answer b ¨ :Ar:.. SUMMARY Lewis electron dot diagrams use dots to represent valence electrons around an atomic symbol. Lewis electron dot diagrams for ions have less (for cations) or more (for anions) dots than the corresponding atom. 10.2: Representing Valence Electrons with Dots is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by Marisa Alviar- Agnew & Henry Agnew. 10.2.2 https://chem.libretexts.org/@go/page/47524 10.3: LEWIS STRUCTURES OF IONIC COMPOUNDS- ELECTRONS TRANSFERRED  LEARNING OBJECTIVES State the octet rule. Define ionic bond. Draw Lewis structures for ionic compounds. In Section 4.7, we demonstrated that ions are formed by losing electrons to make cations, or by gaining electrons to form anions. The astute reader may have noticed something: many of the ions that form have eight electrons in their valence shell. Either atoms gain enough electrons to have eight electrons in the valence shell and become the appropriately charged anion, or they lose the electrons in their original valence shell; the lower shell, now the valence shell, has eight electrons in it, so the atom becomes positively charged. For whatever reason, having eight electrons in a valence shell is a particularly energetically stable arrangement of electrons. The octet rule explains the favorable trend of atoms having eight electrons in their valence shell. When atoms form compounds, the octet rule is not always satisfied for all atoms at all times, but it is a very good rule of thumb for understanding the kinds of bonding arrangements that atoms can make. It is not impossible to violate the octet rule. Consider sodium: in its elemental form, it has one valence electron and is stable. It is rather reactive, however, and does not require a lot of energy to remove that electron to make the Na+ ion. We could remove another electron by adding even more energy to the ion, to make the Na2+ ion. However, that requires much more energy than is normally available in chemical reactions, so sodium stops at a 1+ charge after losing a single electron. It turns out that the Na+ ion has a complete octet in its new valence shell, the n = 2 shell, which satisfies the octet rule. The octet rule is a result of trends in energies and is useful in explaining why atoms form the ions that they do. Now consider an Na atom in the presence of a Cl atom. The two atoms have these Lewis electron dot diagrams and electron configurations: Na ⋅ ¨l : ⋅C.. 1 2 5 [N e] 3s [N e] 3s 3p For the Na atom to obtain an octet, it must lose an electron; for the Cl atom to gain an octet, it must gain an electron. An electron transfers from the Na atom to the Cl atom: ¨ Na ⋅ ↷ ⋅Cl :.. resulting in two ions—the Na+ ion and the Cl− ion: + ¨ − Na :Cl :.. 2 6 [N e] [N e] 3s 3p Both species now have complete octets, and the electron shells are energetically stable. From basic physics, we know that opposite charges attract. This is what happens to the Na+ and Cl− ions: + ¨l :− → N a+ C l− Na + :C or N aCl.. where we have written the final formula (the formula for sodium chloride) as per the convention for ionic compounds, without listing the charges explicitly. The attraction between oppositely charged ions is called an ionic bond, and it is one of the main types of chemical bonds in chemistry. Ionic bonds are caused by electrons transferring from one atom to another. In electron transfer, the number of electrons lost must equal the number of electrons gained. We saw this in the formation of NaCl. A similar process occurs between Mg atoms and O atoms, except in this case two electrons are transferred: The two ions each have octets as their valence shell, and the two oppositely charged particles attract, making an ionic bond: 2− 2+ ¨ 2+ 2− Mg + [:O :] Mg O or M gO.. 10.3.1 https://chem.libretexts.org/@go/page/47526 Remember, in the final formula for the ionic compound, we do not write the charges on the ions. What about when an Na atom interacts with an O atom? The O atom needs two electrons to complete its valence octet, but the Na atom supplies only one electron: ¨ Na ⋅ ↷ ⋅O :. The O atom still does not have an octet of electrons. What we need is a second Na atom to donate a second electron to the O atom: These three ions attract each other to give an overall neutral-charged ionic compound, which we write as Na2O. The need for the number of electrons lost being equal to the number of electrons gained explains why ionic compounds have the ratio of cations to anions that they do. This is required by the law of conservation of matter as well.  EXAMPLE 10.3.1: SYNTHESIS OF CALCIUM CHLORIDE FROM ELEMENTS With arrows, illustrate the transfer of electrons to form calcium chloride from Ca atoms and Cl atoms. Solution A Ca atom has two valence electrons, while a Cl atom has seven electrons. A Cl atom needs only one more to complete its octet, while Ca atoms have two electrons to lose. Thus we need two Cl atoms to accept the two electrons from one Ca atom. The transfer process looks as follows: The oppositely charged ions attract each other to make CaCl2.  EXERCISE 10.3.1 With arrows, illustrate the transfer of electrons to form potassium sulfide from K atoms and S atoms. Answer SUMMARY The tendency to form species that have eight electrons in the valence shell is called the octet rule. The attraction of oppositely charged ions caused by electron transfer is called an ionic bond. The strength of ionic bonding depends on the magnitude of the charges and the sizes of the ions. 10.3: Lewis Structures of Ionic Compounds- Electrons Transferred is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 10.3.2 https://chem.libretexts.org/@go/page/47526 10.4: Covalent Lewis Structures- Electrons Shared  Learning Objectives Define covalent bond. Illustrate covalent bond formation with Lewis electron dot diagrams. Ionic bonding typically occurs when it is easy for one atom to lose one or more electrons and another atom to gain one or more electrons. However, some atoms won’t give up or gain electrons easily. Yet they still participate in compound formation. How? There is another mechanism for obtaining a complete valence shell: sharing electrons. When electrons are shared between two atoms, they make a bond called a covalent bond. Let us illustrate a covalent bond by using H atoms, with the understanding that H atoms need only two electrons to fill the 1s subshell. Each H atom starts with a single electron in its valence shell: H⋅ ⋅ H The two H atoms can share their electrons: H : H We can use circles to show that each H atom has two electrons around the nucleus, completely filling each atom’s valence shell: Because each H atom has a filled valence shell, this bond is stable, and we have made a diatomic hydrogen molecule. (This explains why hydrogen is one of the diatomic elements.) For simplicity’s sake, it is not unusual to represent the covalent bond with a dash, instead of with two dots: H–H Because two atoms are sharing one pair of electrons, this covalent bond is called a single bond. As another example, consider fluorine. F atoms have seven electrons in their valence shell: These two atoms can do the same thing that the H atoms did; they share their unpaired electrons to make a covalent bond. Note that each F atom has a complete octet around it now: We can also write this using a dash to represent the shared electron pair: There are two different types of electrons in the fluorine diatomic molecule. The bonding electron pair makes the covalent bond. Each F atom has three other pairs of electrons that do not participate in the bonding; they are called lone pair electrons. Each F atom has one bonding pair and three lone pairs of electrons. 10.4.1 https://chem.libretexts.org/@go/page/47528 Covalent bonds can be made between different elements as well. One example is HF. Each atom starts out with an odd number of electrons in its valence shell: The two atoms can share their unpaired electrons to make a covalent bond: We note that the H atom has a full valence shell with two electrons, while the F atom has a complete octet of electrons.  Example 10.4.1: Use Lewis electron dot diagrams to illustrate the covalent bond formation in HBr. Solution HBr is very similar to HF, except that it has Br instead of F. The atoms are as follows: The two atoms can share their unpaired electron:  Exercise 10.4.1 Use Lewis electron dot diagrams to illustrate the covalent bond formation in Cl2. Answer When working with covalent structures, it sometimes looks like you have leftover electrons. You apply the rules you learned so far, and there are still some electrons that remain unattached. You can't just leave them there. So where do you put them? Multiple Covalent Bonds Some molecules are not able to satisfy the octet rule by making only single covalent bonds between the atoms. Consider the compound ethene, which has a molecular formula of C H. The carbon atoms are bonded together, with each carbon also bonded 2 4 to two hydrogen atoms. two C atoms = 2 × 4 = 8 valence electrons four H atoms = 4 × 1 = 4 valence electrons total of 12 valence electrons in the molecule If the Lewis electron dot structure was drawn with a single bond between the carbon atoms and with the octet rule followed, it would look like this: 10.4.2 https://chem.libretexts.org/@go/page/47528  Figure 10.4.1 : Incorrect dot structure of ethene. (CK12 License) This Lewis structure is incorrect because it contains a total of 14 electrons. However, the Lewis structure can be changed by eliminating the lone pairs on the carbon atoms and having to share two pairs instead of only one pair. Figure 10.4.2 : Correct dot structure for ethene. (CK12 License) A double covalent bond is a covalent bond formed by atoms that share two pairs of electrons. The double covalent bond that occurs between the two carbon atoms in ethane can also be represented by a structural formula and with a molecular model as shown in the figure below. Figure 10.4.3 : (A) The structural model for C H consists of a double covalent bond between the two carbon atoms and single 2 4 bonds to the hydrogen atoms. (B) Molecular model of C H. 2 4 A triple covalent bond is a covalent bond formed by atoms that share three pairs of electrons. The element nitrogen is a gas that composes the majority of Earth's atmosphere. A nitrogen atom has five valence electrons, which can be shown as one pair and three single electrons. When combining with another nitrogen atom to form a diatomic molecule, the three single electrons on each atom combine to form three shared pairs of electrons. Figure 10.4.4 : Triple bond in N. 2 Each nitrogen atom follows the octet rule with one lone pair of electrons, and six electrons that are shared between the atoms. Summary Covalent bonds are formed when atoms share electrons. Lewis electron dot diagrams can be drawn to illustrate covalent bond formation. Double bonds or triple bonds between atoms may be necessary to properly illustrate the bonding in some molecules. Contributions & Attributions Anonymous by request 10.4: Covalent Lewis Structures- Electrons Shared is shared under a mixed license and was authored, remixed, and/or curated by Marisa Alviar- Agnew & Henry Agnew. 10.4.3 https://chem.libretexts.org/@go/page/47528 10.5: Writing Lewis Structures for Covalent Compounds  Learning Objectives Draw Lewis structures for covalent compounds. The following procedure can be used to construct Lewis electron structures for more complex molecules and ions.  How-to: Constructing Lewis electron structures 1. Determine the total number of valence electrons in the molecule or ion. Add together the valence electrons from each atom. (Recall that the number of valence electrons is indicated by the position of the element in the periodic table.) If the species is a polyatomic ion, remember to add or subtract the number of electrons necessary to give the total charge on the ion. For CO32−, for example, we add two electrons to the total because of the −2 charge. 2. Arrange the atoms to show specific connections. When there is a central atom, it is usually the least electronegative element in the compound. Chemists usually list this central atom first in the chemical formula (as in CCl4 and CO32−, which both have C as the central atom), which is another clue to the compound’s structure. Hydrogen and the halogens are almost always connected to only one other atom, so they are usually terminal rather than central. 3. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond. In H2O, for example, there is a bonding pair of electrons between oxygen and each hydrogen. 4. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen). These electrons will usually be lone pairs. 5. If any electrons are left over, place them on the central atom. We will explain later that some atoms are able to accommodate more than eight electrons. 6. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet. This will not change the number of electrons on the terminal atoms. 7. Final check Always make sure all valence electrons are accounted for and that each atom has an octet of electrons, except for hydrogen (with two electrons). The central atom is usually the least electronegative element in the molecule or ion; hydrogen and the halogens are usually terminal. Now let’s apply this procedure to some particular compounds, beginning with one we have already discussed.  Example 10.5.1: Water Write the Lewis Structure for H2O. Solution Solutions to Example 10.4.1 Steps for Writing Lewis Structures Example 10.5.1 1. Determine the total number of valence electrons in the molecule or Each H atom (group 1) has 1 valence electron, and the O atom (group ion. 16) has 6 valence electrons, for a total of 8 valence electrons. 10.5.1 https://chem.libretexts.org/@go/page/47529  Steps for Writing Lewis Structures Example 10.5.1 2. Arrange the atoms to show specific connections. HOH Because H atoms are almost always terminal, the arrangement within the molecule must be HOH. Placing one bonding pair of electrons between the O atom and each H 3. Place a bonding pair of electrons between each pair of adjacent atom gives atoms to give a single bond. 4. Beginning with the terminal atoms, add enough electrons to each H -O- H atom to give each atom an octet (two for hydrogen). with 4 electrons left over. Each H atom has a full valence shell of 2 electrons. Adding the remaining 4 electrons to the oxygen (as two lone pairs) gives the following structure: 5. If any electrons are left over, place them on the central atom. 6. If the central atom has fewer electrons than an octet, use lone pairs from terminal atoms to form multiple (double or triple) bonds to the Not necessary. central atom to achieve an octet. The Lewis structure gives oxygen an octet and each hydrogen 2 7. Final check. electrons.  Example 10.5.2 Write the Lewis structure for the C H 2O molecule Solution Solutions to Example 10.4.2 Steps for Writing Lewis Structures Example 10.5.2 Each hydrogen atom (group 1) has 1 valence electron, carbon (group 1. Determine the total number of valence electrons in the molecule or 14) has 4 valence electrons, and oxygen (group 16) has 6 valence ion. electrons, for a total of [(2)(1) + 4 + 6] = 12 valence electrons. 2. Arrange the atoms to show specific connections. Because carbon is less electronegative than oxygen and hydrogen is normally terminal, C must be the central atom. Placing a bonding pair of electrons between each pair of bonded atoms gives the following: 3. Place a bonding pair of electrons between each pair of adjacent atoms to give a single bond. 6 electrons are used, and 6 are left over. 10.5.2 https://chem.libretexts.org/@go/page/47529  Steps for Writing Lewis Structures Example 10.5.2 Adding all 6 remaining electrons to oxygen (as three lone pairs) gives the following: 4. Beginning with the terminal atoms, add enough electrons to each atom to give each atom an octet (two for hydrogen). Although oxygen now has an octet and each hydrogen has 2 electrons, carbon has only 6 electrons. Not necessary. 5. If any electrons are left over, place them on the central atom. There are no electrons left to place on the central atom. To give carbon an octet of electrons, we use one of the lone pairs of 6. If the central atom has fewer electrons than an octet, use lone pairs electrons on oxygen to form a carbon–oxygen double bond: from terminal atoms to form multiple (double or triple) bonds to the central atom to achieve an octet. Both the oxygen and the carbon now have an octet of electrons, so this is an acceptable Lewis electron structure. The O has two bonding 7. Final check pairs and two lone pairs, and C has four bonding pairs. This is the structure of formaldehyde, which is used in embalming fluid.  Exercise 10.5.1 Write Lewis electron structures for CO2 and SCl2, a vile-smelling, unstable red liquid that is used in the manufacture of rubber. Answer CO2. Answer SCl2. The United States Supreme Court has the unenviable task of deciding what the law is. This responsibility can be a major challenge when there is no clear principle involved or where there is a new situation not encountered before. Chemistry faces the same challenge in extending basic concepts to fit a new situation. Drawing of Lewis structures for polyatomic ions uses the same approach, but tweaks the process a little to fit a somewhat different set of circumstances. Writing Lewis Structures for Polyatomic Ions (CK-12) Recall that a polyatomic ion is a group of atoms that are covalently bonded together and which carry an overall electrical charge. The ammonium ion, NH , is formed when a hydrogen ion (H ) attaches to the lone pair of an ammonia (NH ) molecule in a + 4 + 3 coordinate covalent bond. 10.5.3 https://chem.libretexts.org/@go/page/47529  Figure 10.5.3 : The ammonium ion. (CK12 License) When drawing the Lewis structure of a polyatomic ion, the charge of the ion is reflected in the number of total valence electrons in the structure. In the case of the ammonium ion: 1 N atom = 5 valence electrons 4 H atoms = 4 × 1 = 4 valence electrons subtract 1 electron for the 1+ charge of the ion total of 8 valence electrons in the ion It is customary to put the Lewis structure of a polyatomic ion into a large set of brackets, with the charge of the ion as a superscript outside of the brackets.  Exercise 10.5.2 Draw the Lewis electron dot structure for the sulfate ion. Answer (CK12 License) Exceptions to the Octet Rule (BC Campus) As important and useful as the octet rule is in chemical bonding, there are some well-known violations. This does not mean that the octet rule is useless—quite the contrary. As with many rules, there are exceptions, or violations. There are three violations to the octet rule. Odd-electron molecules represent the first violation to the octet rule. Although they are few, some stable compounds have an odd number of electrons in their valence shells. With an odd number of electrons, at least one atom in the molecule will have to violate the octet rule. Examples of stable odd-electron molecules are NO, NO2, and ClO2. The Lewis electron dot diagram for NO is as follows: Although the O atom has an octet of electrons, the N atom has only seven electrons in its valence shell. Although NO is a stable compound, it is very chemically reactive, as are most other odd-electron compounds. Electron-deficient molecules represent the second violation to the octet rule. These stable compounds have less than eight electrons around an atom in the molecule. The most common examples are the covalent compounds of beryllium and boron. For example, beryllium can form two covalent bonds, resulting in only four electrons in its valence shell: Boron commonly makes only three covalent bonds, resulting in only six valence electrons around the B atom. A well-known example is BF3: 10.5.4 https://chem.libretexts.org/@go/page/47529 The third violation to the octet rule is found in those compounds with more than eight electrons assigned to their valence shell. These are called expanded valence shell molecules. Such compounds are formed only by central atoms in the third row of the periodic table or beyond that have empty d orbitals in their valence shells that can participate in covalent bonding. One such compound is PF5. The only reasonable Lewis electron dot diagram for this compound has the P atom making five covalent bonds: Formally, the P atom has 10 electrons in its valence shell.  Example 10.5.3: Octet Violations Identify each violation to the octet rule by drawing a Lewis electron dot diagram. a. ClO b. SF6 Solution a. With one Cl atom and one O atom, this molecule has 6 + 7 = 13 valence electrons, so it is an odd-electron molecule. A Lewis electron dot diagram for this molecule is as follows: b. In SF6, the central S atom makes six covalent bonds to the six surrounding F atoms, so it is an expanded valence shell molecule. Its Lewis electron dot diagram is as follows:  Exercise 10.5.3: Xenon Difluoride Identify the violation to the octet rule in XeF2 by drawing a Lewis electron dot diagram. Answer The Xe atom has an expanded valence shell with more than eight electrons around it. 10.5.5 https://chem.libretexts.org/@go/page/47529 Summary Lewis dot symbols provide a simple rationalization of why elements form compounds with the observed stoichiometries. A plot of the overall energy of a covalent bond as a function of internuclear distance is identical to a plot of an ionic pair because both result from attractive and repulsive forces between charged entities. In Lewis electron structures, we encounter bonding pairs, which are shared by two atoms, and lone pairs, which are not shared between atoms. Lewis structures for polyatomic ions follow the same rules as those for other covalent compounds. There are three violations to the octet rule: odd-electron molecules, electron-deficient molecules, and expanded valence shell molecules. 10.5: Writing Lewis Structures for Covalent Compounds is shared under a mixed license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 10.5.6 https://chem.libretexts.org/@go/page/47529 10.5.7 https://chem.libretexts.org/@go/page/47529 10.5.8 https://chem.libretexts.org/@go/page/47529 10.6: Resonance - Equivalent Lewis Structures for the Same Molecule  Learning Objectives Explain the concept of resonance and how it works with within molecules. Resonance There are some cases in which more than one viable Lewis structure can be drawn for a molecule. An example is the ozone (O 3 ) molecule in Figure 10.6.1. There are a total of 18 electrons in the structure and so the following two structures are possible. Figure 10.6.1 : Resonance forms of ozone. Note the use of the double-headed arrow. The structure on the left (10.6.1) can be converted to the structure on the right by a shifting of electrons without altering the positions of the atoms. It was once thought that the structure of a molecule such as O consisted of one single bond and one double bond which then 3 shifted back and forth as shown above. However, further studies showed that the two bonds are identical. Any double covalent bond between two given atoms is typically shorter than a single covalent bond. Studies of the O and other similar molecules 3 showed that the bonds were identical in length. Interestingly, the length of the bond is in between the lengths expected for an O−O single bond and a double bond. Resonance is the use of two or more Lewis structures to represent the covalent bonding in a molecule. One of the valid structures is referred to as a resonance structure. It is now understood that the true structure of a molecule which displays resonance is that of an average or a hybrid of all the resonance structures. In the case of the O molecule, each of the covalent bonds between O atoms 3 are best thought of as being "one and a half" bonds, as opposed to either a pure single bond or a pure double bond. This "half-bond" can be shown as a dotted line in both the Lewis structure and the molecular model (Figure 10.6.2). Figure 10.6.2 : "Half-bond" model of ozone molecule. This is a better description of the electronic structure of ozone than either of the resonance structures in Figure 10.6.1. Many polyatomic ions also display resonance. In some cases, the true structure may be an average of three valid resonance structures, as in the case of the nitrate ion, NO in Figure 10.6.3. − 3 Figure 10.6.3 : Resonance structure of nitrate anion. The bond lengths between the central N atom and each O atom are identical and the bonds can be approximated as being equal to one and one-third bonds. Summary Resonance structures are averages of different Lewis structure possibilities. Bond lengths are intermediate between covalent bonds and covalent double bonds. 10.6: Resonance - Equivalent Lewis Structures for the Same Molecule is shared under a CK-12 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 10.6.1 https://chem.libretexts.org/@go/page/47531 10.7: PREDICTING THE SHAPES OF MOLECULES  LEARNING OBJECTIVE Determine the shape of simple molecules. Molecules have shapes. There is an abundance of experimental evidence to that effect—from their physical properties to their chemical reactivity. Small molecules—molecules with a single central atom—have shapes that can be easily predicted. The basic idea in molecular shapes is called valence shell electron pair repulsion (VSEPR). VSEPR says that electron pairs, being composed of negatively charged particles, repel each other to get as far away from one another as possible. VSEPR makes a distinction between electron group geometry, which expresses how electron groups (bonds and nonbonding electron pairs) are arranged, and molecular geometry, which expresses how the atoms in a molecule are arranged. However, the two geometries are related. There are two types of electron groups: any type of bond—single, double, or triple—and lone electron pairs. When applying VSEPR to simple molecules, the first thing to do is to count the number of electron groups around the central atom. Remember that a multiple bond counts as only one electron group. Any molecule with only two atoms is linear. A molecule whose central atom contains only two electron groups orients those two groups as far apart from each other as possible—180° apart. When the two electron groups are 180° apart, the atoms attached to those electron groups are also 180° apart, so the overall molecular shape is linear. Examples include BeH2 and CO2: Figure 10.7.1: Beryllium hydride and carbon dioxide bonding. The two molecules, shown in the figure below in a "ball and stick" model. Figure 10.7.2: Beryllium hydride and carbon dioxide models. (CK12 Licence) A molecule with three electron groups orients the three groups as far apart as possible. They adopt the positions of an equilateral triangle— 120° apart and in a plane. The shape of such molecules is trigonal planar. An example is BF3: Figure 10.7.3: Boron trifluoride bonding. (CK12 Licence) Some substances have a trigonal planar electron group distribution but have atoms bonded to only two of the three electron groups. An example is GeF2: Figure 10.7.4: Germanium difluoride bonding. From an electron group geometry perspective, GeF2 has a trigonal planar shape, but its real shape is dictated by the positions of the atoms. This shape is called bent or angular. A molecule with four electron groups about the central atom orients the four groups in the direction of a tetrahedron, as shown in Figure 10.7.1 Tetrahedral Geometry. If there are four atoms attached to these electron groups, then the molecular shape is also tetrahedral. Methane (CH4) is an example. 10.7.1 https://chem.libretexts.org/@go/page/47532  Figure 10.7.5: Tetrahedral structure of methane. (CK12 Licence) This diagram of CH4 illustrates the standard convention of displaying a three-dimensional molecule on a two-dimensional surface. The straight lines are in the plane of the page, the solid wedged line is coming out of the plane toward the reader, and the dashed wedged line is going out of the plane away from the reader. Figure 10.7.6: Methane bonding. (CK12 Licence) NH3 is an example of a molecule whose central atom has four electron groups, but only three of them are bonded to surrounding atoms. Figure 10.7.7: Ammonia bonding. (CK12 Licence) Although the electron groups are oriented in the shape of a tetrahedron, from a molecular geometry perspective, the shape of NH3 is trigonal pyramidal. H2O is an example of a molecule whose central atom has four electron groups, but only two of them are bonded to surrounding atoms. Figure 10.7.8: Water bonding. Although the electron groups are oriented in the shape of a tetrahedron, the shape of the molecule is bent or angular. A molecule with four electron groups about the central atom, but only one electron group bonded to another atom, is linear because there are only two atoms in the molecule. Double or triple bonds count as a single electron group. The Lewis electron dot diagram of formaldehyde (CH2O) is shown in Figure 10.7.9. Figure 10.7.9: Lewis Electron Dot Diagram of Formaldehyde. The central C atom has three electron groups around it because the double bond counts as one electron group. The three electron groups repel each other to adopt a trigonal planar shape. Figure 10.7.10: Formaldehyde bonding. (The lone electron pairs on the O atom are omitted for clarity.) The molecule will not be a perfect equilateral triangle because the C–O double bond is different from the two C–H bonds, but both planar and triangular describe the appropriate approximate shape of this molecule. Table 10.7.1 summarizes the shapes of molecules based on the number of electron groups and surrounding atoms. 10.7.2 https://chem.libretexts.org/@go/page/47532  Table 10.7.1: Summary of Molecular Shapes Number of Electron Groups on Central Atom Number of Bonding Groups Number of Lone Pairs Electron Geometry Molecular Shape 2 2 0 linear linear 3 3 0 trigonal planar trigonal planar 3 2 1 trigonal planar bent 4 4 0 tetrahedral tetrahedral 4 3 1 tetrahedral trigonal pyramidal 4 2 2 tetrahedral bent  EXAMPLE 10.7.1 What is the approximate shape of each molecule? a. PCl3 b. NOF Solution The first step is to draw the Lewis structure of the molecule. For PCl , the electron dot diagram is as follows: 3 The lone electron pairs on the Cl atoms are omitted for clarity. The P atom has four electron groups with three of them bonded to surrounding atoms, so the molecular shape is trigonal pyramidal. The electron dot diagram for NOF is as follows: The N atom has three electron groups on it, two of which are bonded to other atoms. The molecular shape is bent.  EXERCISE 10.7.1 What is the approximate molecular shape of CH 2 Cl 2 ? Answer Tetrahedral  EXERCISE 10.7.2 Ethylene (C H ) has two central atoms. Determine the geometry around each central atom and the shape of the overall molecule. 2 4 (Hint: hydrogen is a terminal atom.) Answer Trigonal planar about both central C atoms. SUMMARY The approximate shape of a molecule can be predicted from the number of electron groups and the number of surrounding atoms. 10.7: Predicting the Shapes of Molecules is shared under a mixed 3.0 license and was authored, remixed, and/or curated by Marisa Alviar-Agnew & Henry Agnew. 10.7.3 https://chem.libretexts.org/@go/page/47532 10.8: Electronegativity and Polarity - Why Oil and Water Do not Mix  Learning Objectives Explain how polar compounds differ from nonpolar compounds. Determine if a molecule is polar or nonpolar. Given a pair of compounds, predict which would have a higher melting or boiling point. Bond Polarity The ability of an atom in a molecule to attract shared electrons is called electronegativity. When two atoms combine, the difference between their electronegativities is an indication of the type of bond that will form. If the difference between the electronegativities of the two atoms is small, neither atom can take the shared electrons completely away from the other atom, and the bond will be covalent. If the difference between the electronegativities is large, the more electronegative atom will take the bonding electrons completely away from the other atom (electron transfer will occur), and the bond will be ionic. This is why metals (low electronegativities) bonded with nonmetals (high electronegativities) typically produce ionic compounds. A bond may be so polar that an electron actually transfers from one atom to another, forming a true ionic bond. How do we judge the degree of polarity? Scientists have devised a scale called electronegativity, a scale for judging how much atoms of any element attract electrons. Electronegativity is a unitless number; the higher the number, the more an atom attracts electrons. A common scale for electronegativity is shown in Figure 10.8.1. Figure 10.8.1 : Electronegativities of the Elements. Electronegativities are used to determine the polarity of covalent bonds. The polarity of a covalent bond can be judged by determining the difference of the electronegativities of the two atoms involved in the covalent bond, as summarized in the following table: difference of the electronegativities of the two atoms involved in the covalent bond Electronegativity Difference Bond Type 0–0.4 pure covalent 0.5–2.0 polar covalent >2.0 likely ionic Nonpolar Covalent Bonds A bond in which the electronegativity difference is less than 1.9 is considered to be mostly covalent in character. However, at this point we need to distinguish between two general types of covalent bonds. A nonpolar covalent bond is a covalent bond in which the bonding electrons are shared equally between the two atoms. In a nonpolar covalent bond, the distribution of electrical charge is balanced between the two atoms. 10.8.1 https://chem.libretexts.org/@go/page/47534  Figure 10.8.2 : A nonpolar covalent bond is one in which the distribution of electron density between the two atoms is equal. The two chlorine atoms share the pair of electrons in the single covalent bond equally, and the electron density surrounding the Cl 2 molecule is symmetrical. Also note that molecules in which the electronegativity difference is very small (