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16.6: Molecular Structure, Bonding, and Acid-Base Behavior

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    Introduction

    When looking at acid base behavior there are several principles one needs to consider before proceeding. First, all reactions are in theory reversible, and so one should always look at the back reaction, that is the reaction of the conjugates, even if it does not occur. Is the conjugate stable? Does it have a stable or reactive structure? One should also look at charge distribution as it clearly influences reactivity. So you are sort of looking at two things, how stable a species is (what structures and feature lead to nonreactivity) and how unstable a species is (what structures and features lead to an unstable species that is going to react).

    The second thing is sort of common sense, but worth stating. The stronger the bond that holds hydrogen to an atom, the weaker the acid. That is, weakly bonded hydrogens are more acidic than strongly bonded ones.

    Binary compounds of hydrogen

    Binary compounds are compounds between two different types of elements, and can consist of more than two atoms. Binary compounds with hydrogen can be acidic (proton donors), basic (proton acceptors), or not show acidic or basic properties. The behavior is defined by the direction of the dipole moment, as indicated in figure 16.6.1, where acidic compounds occur when hydrogen is attached to something more electronegative and basic compounds occur when the bond is to something less electronegative.

    aREBc16fig1.JPG

    Figure\(\PageIndex{1}\): Rule of thumb showing acidity/basicity as a function of polarity. Note, basic hydrides are when you have an ionic compound.

    Basic Hydrides

    Basic hydrides occur when hydrogen is attached to a highly electropositive element where the difference in electronegativity is great enough to form an ionic compound (review section 8.7.2). That is, when it forms the hydride ion H-. Soluble hydrides are strong bases (review section 16.5.3.1), and the basicity can be viewed as a property of fact that the small hydride ion is very reactive and grabs a proton from water.

    \[H^-(aq) + H_2O(aq) \rightarrow H_2(g) +OH^-(aq)\]

    It needs to be emphasized that hydride is a reactive anion, and so any water soluble hydride will release hydride that can react with water to form hydrogen and hydroxide. These are strong bases.

    Acidic Hydrides

    These occur when a hydrogen atom is attached to a nonmetal, where the more electronegative nonmetal pulls electron density from the hydrogen making it more positive and thus potentially more reactive. But there is no one factor alone that defines the acid strength of a hydride, although there are some periodic trends we can see, and these can be understood in terms of trends in electronegativity and bond strength, noting the latter can often be correlated to bond lengths and the sizes of atoms.

    Periodic Trend for Acidic Hydrides

    Figure 16.6.2 shows the ionization constants for non-metal binary hydrides and two periodic trends become obvious. Note, the hydrides in red have large ionization constants and so are considered as strong acids.

    acidstrength-MH.png

    Figure\(\PageIndex{2}\):Periodic trends acid strength in terms of equilibrium constants

    1. Going across a period the acid strength increases as there is an increase in electronegativity and the molecule gets more polar, with the hydrogen getting a larger partial positive charge. Note, ammonia is a base and actually accepts protons, and methane is symmetric

    2. Going down a group the acid strength increases because the bond strength decreases as a function of increasing size of the nonmetal, and this has a larger effect than the electronegativity. In fact HF is a weak acid because it is so small that the hydrogen-fluorine bond is so strong that it is hard to break. Remember, the weaker the bond, the strong the acid strength. This is further illustrated in figure 16.6.3 where the weakest bond has produces the strongest acid.

     

    Relative Acid Strength HF << HCl < HBr < HI
    H–X Bond Energy (kJ/mol) 570   432   366   298
    Ka 10-3   107   109   1010

    Table\(\PageIndex{1}\) Acid strength as function of bond energy

    Binary Compounds of Oxygen

    Binary compounds of oxygen involving metals are basic and known as basic anhydrides, while binary compounds of oxygen with nonmetals are acidic and known as acid anhydrides.

    Basic Anhydrides

    In section 16.5.2.4 we noted that soluble metal oxides are strong bases and release the oxide ion (O-2) into water and that pulls a proton from the water 

    \[ O^{-2}(aq)+H_2O(l)  \rightarrow 2OH^-(aq) \\ \; \\ and \\ \; \\ CaO + H_2O \rightarrow Ca(OH)_2  \]

    So a metal oxide like CaO can be viewed as Ca(OH)2 with the water removed.  Note we used the word "soluble" and many oxides are insoluble, as we shall see in the next chapter (for example, the equilibrium constant for the solubility of aluminum oxide is 1.9x10-33 and so it is not a strong base because it does not dissolve).  But soluble metal oxides are strong bases because the oxide pulls a proton from the water.

    Acid Anhydrides

    Acid anhydrides react with water to form oxyacids. This is an important class of acids, where the oxygens are attached to a nonmetal and the acidic hydrogens are attached to the oxygen. This also explains the origins of acid rain, and as nonmetal oxides are formed in the combustion process and can then combine with atmospheric water.

    \[\begin{align}\text{nonmetal oxide + water} & \rightleftharpoons \text{oxyacid}\\CO_{2}(g)+H_{2}O(l) & \rightleftharpoons H_{2}CO_{3}(aq) \\ SO_{2}(g)+H_{2}O(l) & \rightleftharpoons H_{2}SO_{3}(aq) \\ SO_{3}(g)+H_{2}O(l) & \rightleftharpoons H_{2}SO_{4}(aq) \\ N_2O_3(g) + H_2O(l)  & \rightleftharpoons 2HNO_2(aq)  \\ N_2O_5(l) + H_2O(l) & \rightleftharpoons 2HNO_3(aq) \end{align}\]

    Note, 79% of the atmosphere is nitrogen and the dominant combustion biproduct of nitrogen is NO, with smaller amounts of N2O and NO2. The NO is further oxidized to NO2, and the NO and NO2, combine to form N2O3, which is the acid anhydride of nitrous acid. So nonmetal oxides produce oxyacids in nature and these are an important class of acids that you need to be familiar with.

     

    Oxyacids

    Oxyacids are compounds of the general formula HnXOm, where X is a nonmetal and the acidic hydrogens are attached to oxygen (not the nonmetal).  As seen above, these can be formed by the reaction of nonmetal oxides with water.  There are two trends on the acid strength of Oxyacids that you need to be familiar with.

    1. Homologous structure (nonmetals of the same periodic group that have identical numbers of H and O)
    2. Structures with the same central atom that have varying numbers of oxygens.

    These, along with the binary hydrides are often referred to as inorganic acids, in contrast to the carboxylic acids (see below), which are often referred to as organic acids. 

    Oxyacids with Homologous Structures

    These are structures of the same formula but different nonmetals, like hypochlorous, hypobromous and hypoiodous acids (HClO, HBrO & HIO, with the generic structure of HXO, where X is a halogen). Now note, we write the acidic hydrogen first because it is like the cation of an analogous salt, but it is really covalently bonded to the oxygen and not the nonmetal (X), so the structure is of the form H-O-X. What is important to note here is that the acidic bonds that need to be broken are all between hydrogen and oxygen, and so are of the same basic bond length, and the strength is influenced by the electronegativity of X's ability to pull electron density towards it. This is called an inductive effect as the nonmetal's electronegativity induces the effect across the oxygen and makes the oxygen-hydrogen bond more polar. Thus the trend is increasing strength in the direction of increasing electronegativity.

     

    HOX Electronegativity of X Ka
    HOCl 3.0 4.0 × 10−8
    HOBr 2.8 2.8 × 10−9
    HOI 2.5 3.2 × 10−11

    Table\(\PageIndex{2}\): Relationship of electronegativity to acid ionization constant. Note this is an inductive effect, and unlike the binary hydride, the hydrogen is not bonded to the nonmetal, with the result that the trend is the opposite of the trend in table 16.6.1.

    HClO>HBrO>HIO (table 16.6.2)
    HI>HBr>HCl>>HF (table 16.6.1)

    Oxyacids with Same Central Atom

    Here we are looking at the trend for acids like perchloric, chloric, chlorous and hypochlous, and see the greater the number of oxygens the stronger the acid. This can be explained in several ways. The most valid deals with the inductive effect, where oxygen is the second most electronegative element, and so the more oxygens attached to the nonmetal of the acid, the greater their pull on electron density across the OH bond, and the more polar the bond. This can be seen in figure 16.6.3.

     

    cbfd507c1a19b98db1f5201e5ab728ed.jpg

    Figure\(\PageIndex{3}\): Increasing number of oxygens increases Ka as evidenced by the decreased electron density on the acidic hydrogen (which is most blue in HClO4 ). Note, Ka =10-pKa , and so the larger pKa , the smaller Ka .

    Another facet deals with the stability of the conjugate anion, that is, how strong is the back reaction. For example, in hypochlorite (ClO-), the negative charge is spread over two nuclei and is thus more concentrated than in other oxyanions of chlorine, like perchlorate (ClO4-), where it is spread out over 5 nuclei. This increased charge density makes the back reaction stronger in hypochlorite than perchlorate, and so the forward reaction would be weak, making perchlorate stronger in the forward direction.

     

    99db2ab846ba9253de6cfe75d12abf9c.jpg

    Figure\(\PageIndex{4}\): This figure shows two things. First, the larger the ion the more dispersed the charge and thus the less the charge density, making the perchlorate the more stable anion. The second argument this image puts forth is that the greater the number of resonance structures the more dispersed the charge density, which is correct, but I would argue against the resonance structures as drawn. These structures follow the lowest formal charge of the chlorine, but invoke pi bonds when there are no unhybridized p orbitals to make them. This is an interesting issue that chemist argue about, and represent a shortcoming of Lewis dot structures.

    Carboxylic Acids

    Many organic acids have the carboxylic acid functional group (COOH) that is important to be able to identify. In gen chem 1 you were required to memorize several of these, like acetic acid, oxalic acid and phthalic acid.

    Various_carboxylic_acids.jpg

    Figure\(\PageIndex{5}\): Some organic acids that show the caboxylic acid group in red. Note, Adipic acid is a diprotic acid because it has two carboxylic groups. The hydrogens in black are strongly held and not acidic, while the ones in red are acidic.

    Resonance Stabilization

    Resonance structures are a shortcoming of Lewis dot structures that treat a bond as an electron pair shared between two nuclei, and if electrons are shared between more than two nuclei, you need to draw a resonance structure. If the species is an ion, and the charge can be distributed across multiple nuclei, the charge density is lowered and species is more stable. Figure 16.6.5 shows this effect for the carboxylate ion of the back reaction, where in truth the charge of the extra electron is delocalized over three nuclei (the carbon and the two oxygens) via a pi bond that is represented by both resonance structures. This means the ion is more stable (less reactive), making the back reaction weaker, and thus the forward reaction stronger.

    carbres1.gif

    Figure\(\PageIndex{5}\): Illustration of effect of resonance structures on carboxylic acid/carboxylate reactions.

    It is important to remember that when looking at chemical structure of an acid, you also need to look at its conjugate base. In this example, the carboxylate group on the product side (conjugate base of the weak carboxylic acid) is stabilized by charge delocalization, making the back reaction weaker, and thus contributes to increasing acid strength (the forward reaction).

    Robert E. Belford (University of Arkansas Little Rock; Department of Chemistry). The breadth, depth and veracity of this work is the responsibility of Robert E. Belford, rebelford@ualr.edu. You should contact him if you have any concerns. This material has both original contributions, and content built upon prior contributions of the LibreTexts Community and other resources, including but not limited to:


    This page titled 16.6: Molecular Structure, Bonding, and Acid-Base Behavior is shared under a CC BY-NC-SA 3.0 license and was authored, remixed, and/or curated by Robert Belford.

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