Chapter 9. Chemical Bonding I: Basic Concepts

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1 Chapter 9 Chemical Bonding I: Basic Concepts This is the first of two chapters on bonding. Upon completion of Chapter 9, the student should be able to: 1. Identify the valence electrons for all representative elements. 2. Rationalize why alkali metals and alkaline earth metals usually form cations and oxygen and the halogens usually form anions using Lewis dot symbols in the discussion. 3. Use Lewis dot symbols to show the formation of both ionic and molecular compounds. 4. Define lattice energy, Coulomb s law and the Born-Haber cycle. 5. Demonstrate how the Born-Haber cycle is an application of Hess s law and use the Born-Haber cycle to determine lattice energy for an ionic solid. 6. Identify covalent compounds, the type of covalent bonds present, and the number of lone pairs of electrons using Lewis structures. 7. Relate types of bonds to bond length and bond strength. 8. Compare and contrast various properties expected for ionic compounds versus covalent compounds. 9. Identify ionic, polar covalent and (nonpolar) covalent bonds using the concepts of electronegativity. 10. Predict the relative changes in electronegativity with respect to position on the periodic table. 11. Use the concept of electronegativity to rationalize oxidation numbers. 12. Use Lewis dot and the octet rule to write Lewis structures of compounds and ions. 13. Apply the concept of formal charge to predict the most likely Lewis structure of a compound. 14. Explain how Lewis structures are inadequate to explain observed bond length (bond types) in some compounds and how the concept of resonance must be invoked. 15. Recall several common examples in which the octet rule fails. 16. Demonstrate, using Lewis structures, the formation of a coordinate covalent (dative) bond. 17. Use Lewis structures and bond energies to predict heats of reaction. 18. Rationalize why enthalpy change for breaking chemical bonds is positive and the formation of chemical bonds is negative.

2 Section 9.1 Lewis Dot Symbols In order to use Lewis dot symbols correctly, our students must first understand what valence electrons are. It is well to review that concept first before proceeding. We should also be aware that Lewis dot symbols are best reserved for row two elements. Lewis dot symbols can be used for transition metals in some cases, but in general, it is not advisable to attempt to use a simple model like Lewis dot on complex molecules. Section 9.2 The Ionic Bond Coulomb s law states kq Q F = 1 2 r 2 when F is the force, either attractive or repulsive, depending upon whether the charges are similar, k is a constant, Q1 and Q2 are the charges and r is the distance between centers of charge. If we have a cation being attracted to an anion, then Coulomb s law describes the force of attraction. It follows then that a cation with a positive two charge should have a stronger attraction for an anion than a cation with a positive one charge assuming their ionic radii are the same. We saw earlier that cations become smaller as their charges increase (Fe 2+ versus Fe 3+, for example); therefore, because the r 2 term is in the denominator, the smaller cation would be expected to have a stronger force of attraction. The same will be true for anions with a minus one versus minus two charge assuming the radii are the same. However, note that anions get larger as their negative charge increases thus the r 2 of Coulomb s law acts to decrease the force of attraction. For some students, it may be easier for them to follow the process if the reaction of calcium and oxygen atoms is written as follows: Ca + O Ca O The two valence electrons on the calcium atom are paired and correspond to the 4s 2 electrons. The same would be true for the oxygen where one pair of electrons corresponds to the 2s 2 electrons and the other pair is the pair of electrons that are in the same 2p orbital. We will see later that it really doesn t matter where the valence electrons come from; they all must be accounted for when we do Lewis dot symbols anyway.

3 Section 9.3 Lattice Energy of Ionic Compounds Lattice energy is the energy required to separate one mole of solid ionic compound into gaseous ions. This quantity must always be endothermic. We can use a variation on the law of conservation of energy known as the Born-Haber cycle to determine lattice energy. The Born-Haber cycle uses the concepts of sublimation, energy of dissociation, ionization energy, and electron affinity in determining lattice energy. In that respect, the Born-Haber cycle is a good review of several concepts; however, it tends to be challenging for many students. Section 9.4 The Covalent Bond Covalent bonds are those bonds where electrons are shared. The most equal sharing of electrons occurs in homonuclear diatomic molecules where neither of the identical atoms would have a greater attraction for electrons than the other atom would. This perfectly equal sharing results in a pure covalent bond. The other extreme is the ionic bond that is discussed in Section 9.2 where an electron is given up by one atom to form a cation and accepted by another atom to form an anion. In between these two extremes are polar covalent bonds where electrons spend more time around one atom than the other. There is yet one other covalent bond which is discussed in Section 9.9. The coordinate covalent bond or dative bond is formed when one atom donates a pair of electrons to another atom, which has fewer than eight electrons about it to form a chemical bond. The example your author uses is NH3 donating a pair of electrons to BF3. One uses the octet rule, surrounding each atom with eight electrons, to draw Lewis structures of molecules. Lewis structures give rise to single, double, and triple bonds and also to non-bonding pairs of electrons. Single bonds are longer but weaker than double bonds, which are longer and weaker than triple bonds. We will see in Chapter 10 that lone pairs of non-bonding electrons can have an influence on the structure and reactivity of molecules. Section 9.5 Electronegativity Electronegativity is the ability of an atom to draw electrons around itself within a given molecule. Note that there is a bit of similarity between electronegativity and electron affinity, but electronegativity is within a molecule while electron affinity is for isolated atoms. It is interesting to note that the concept of electronegativity was developed by Linus Pauling, the only person to be the sole recipient of two Nobel Prizes the Nobel Prize in Chemistry and the Nobel Peace Prize. The larger the difference of electronegativity between atoms, the more polar the bond is between the atoms. It is suggested

4 that a difference of 2.0 or more in electronegativities will result in an ionic bond. The most electronegative element is fluorine; therefore, fluorine does not wish to share its electrons with other atoms. It is also for this reason that the oxidation number of fluorine is always minus one when it combines with other elements to form compounds. Fluorine will only form single bonds (no double or triple bonds) for this same reason. Section 9.6 Writing Lewis Structures When writing Lewis structures, one counts the total number of valence electrons from all the atoms in the molecule and attempts to arrange them so that all atoms, with the exception of hydrogen, which will only have two, have eight electrons surrounding them in either single, double or triple bonds or as lone pairs. Since hydrogen and fluorine can only form single bonds, they are always terminal atoms in Lewis structures. Be sure when counting valence electrons that the charge on the species being examined is accounted for. If the species is a cation, then the positive charge of the ion is subtracted from the total number of valence electrons to give the correct number of electrons to be used in the Lewis structure. If the species is an anion, then the negative charge is added to the total number of valence electrons. Section 9.7 Formal Charge and Lewis Structure Formal charge is a bookkeeping method used to assist in assigning arrangement of atoms in Lewis structures. Formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure. The sum of formal charges of all atoms within a Lewis structure must be equal to the charge on the species. That is, the sum of formal charges will be zero for neutral compounds, positive for cations, and negative for anions. Three guidelines for use of formal charges are: 1. Lewis structure with no formal charges are preferred for neutral molecules 2. Structures with the smallest possible formal charges are most likely 3. In general, negative formal charges are found on the most electronegative atoms Section 9.8 The Concepts of Resonance Your author uses the analogy of a rhinoceros as being a cross between a griffin, a mythical animal part eagle and part lion and a unicorn, another mythical animal that is horselike with a

5 single horn in the center of its forehead. This analogy is used to discuss resonance forms for such molecules as ozone or benzene where Lewis structures do not adequately describe what is known to be the true structure. Thus, the two Lewis structures of benzene are like the griffin and the unicorn. Neither one actually exist but a combination of the two are required to describe what we know is the real structure of benzene. Section 9.9 Exceptions to the Octet Rule Since there are so many exceptions to the octet rule, perhaps it would be better to call it the octet generalization. The value of Lewis structures and the octet rule are that they are simple to use and describe many common compounds. There are several examples of exceptions to the octet rule. Such molecules as those with odd number of valence electrons, (NO, for example), that form an incomplete octet like BF3 and those that use the expanded octet such as SF6, are all examples of violations to the octet rule. Section 9.10 Bond Energy When thermochemical information about specific compounds is not known, one can use bond dissociation energy to estimate enthalpies of reaction. It should be understood that using average bond energies will result in only an estimate and will not agree with H s found using actual thermochemical values. Your author uses the equation: H = BE (reactants) - BE (products) and which is correct; however, this equation may give rise to confusion for your students because the term refers to final minus initial states. It should seem that H o should equal the sum of the bond energies of the products minus the sum of the bond energies of the reactants, but this will result in H o with the wrong sign. One way to get around this dilemma is to do the following: a) determine the type and number of bonds broken and total energy required to break these bonds b) determine the type and number of bonds formed and total energy released when these bonds are formed Since breaking bonds requires energy, it must be an endothermic process so that H for bond breaking (part a above) must have a positive sign while bond formation releases energy thus that process (part b above) must be exothermic and have a negative H. The overall H is just the addition of those two values or

6 H = Hbond breaking + Hbond forming keeping in mind that the first term is positive and the second is negative.

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