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Introduction to Ionic Mechanisms Part I: Fundamentals of Bronsted-Lowry Acid-Base Chemistry

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Introduction to Ionic Mechanisms Part I: Fundamentals of Bronsted-Lowry Acid-Base Chemistry INTRODUCTION TO IONIC MECHANISMS PART I: FUNDAMENTALS OF BRONSTED-LOWRY ACID-BASE CHEMISTRY HYDROGEN ATOMS AND PROTONS IN ORGANIC MOLECULES - A hydrogen atom that has lost its only electron is sometimes referred to as a proton. That is because once the electron is lost, all that remains is the nucleus, which in the case of hydrogen consists of only one proton. The large majority of organic reactions, or transformations, involve breaking old bonds and forming new ones. If a covalent bond is broken heterolytically, the products are ions. In the following example, the bond between carbon and oxygen in the t-butyl alcohol molecule breaks to yield a carbocation and hydroxide ion. H3C CH3 H3C OH H3C + OH CH3 H3C A tertiary Hydroxide carbocation ion The full-headed curved arrow is being used to indicate the movement of an electron pair. In this case, the two electrons that make up the carbon-oxygen bond move towards the oxygen. The bond breaks, leaving the carbon with a positive charge, and the oxygen with a negative charge. In the absence of other factors, it is the difference in electronegativity between the two atoms that drives the direction of electron movement. When pushing arrows, remember that electrons move towards electronegative atoms, or towards areas of electron deficiency (positive, or partial positive charges). The electron pair moves towards the oxygen because it is the more electronegative of the two atoms. If we examine the outcome of heterolytic bond cleavage between oxygen and hydrogen, we see that, once again, oxygen takes the two electrons because it is the more electronegative atom. Hydrogen is left with only a positive charge. In other words, it becomes a proton. H3C H3C H H3C O H H3C O + H3C H3C An alkoxide Hydrogen ion, ion or proton BRONSTED-LOWRY ACIDS AND “ACIDIC PROTONS” - A hydrogen bonded to a very electronegative atom makes for a highly polar bond. The dipole moment favors electron density around the more electronegative atom, leaving the hydrogen with a partial positive charge. This bond is different from other bonds in the molecule because of its propensity to break into a negative ion and a positive hydrogen ion. This propensity is driven by the tendency of the more electronegative atom to take up the electrons that make up the covalent bond. In fact, one can write resonance structures for such molecule that show the bond in question already broken. The t-butyl alcohol molecule can be used again to illustrate this point. H3C H3C H3C O H H3C O H H3C H3C t-butyl alcohol (I) (II) The greatest contributor to the hybrid is obviously structure I because it is neutral. Structure II has charge separation and therefore is a minor contributor. However, the significance of structure II is that it shows the negative character of the oxygen and the positive character of the hydrogen, and therefore the polarity of this bond. A better representation of the hybrid could be structure III, which shows the oxygen with partial negative character, and the hydrogen with partial positive character. H C 3 δ− δ+ H3C O H (III) H3C In the Bronsted-Lowry theory of acids and bases, an acid is a hydrogen ion donor, or proton donor, and a base is a hydrogen ion acceptor, or proton acceptor. Hydrogen atoms that have a substantial degree of partial positive charge (i.e. low electron density around them) are commonly referred to as acidic protons. In the example above, the hydrogen bonded to oxygen is considered to be acidic, and the molecule as a whole is considered a Bronsted acid because it has a propensity to release a hydrogen ion, or proton. ACID-BASE REACTIONS AS “PROTON” TRANSFERS - When a Bronsted acid (or simply acid) reacts with a Bronsted base (or simply base) a proton is transferred from the acid to the base. This results in formation of another acid, called the conjugate acid, and another base, called the conjugate base. For example, when hydroxide ion (a base) reacts with hydrogen chloride (an acid), a new acid (water) is formed. Water is then the conjugate acid of hydroxide ion. Likewise, a new base (chloride ion) is formed. Chloride ion is then the conjugate base of hydrogen chloride. The reaction is an equilibrium process because the new acid and the new base can react together to revert to the original reactants. Therefore we can also say that hydroxide ion is the conjugate base of water, and that hydrogen chloride is the conjugate acid of chloride ion. This relativity of concepts is characteristic of the Bronsted-Lowry acid-base theory. + H Cl O H O H H + Cl δ+ δ− hydroxide hydrogen water chloride ion chloride ion BASE1 + ACID1 ACID2 + BASE2 conjugate acid-base pairs Conjugate acid-base pairs differ only by a proton. Other examples of conjugate acid-base pairs are H2O / H3O+ and NH3 / NH4+. The above reaction also shows the direction of electron movement. In acid-base reactions, electron movement always originates at the base and moves towards the acidic proton in the acid. The base is the electron-rich species. We can identify bases because they usually have an atom with unshared electron pairs (lone pairs). Sometimes this atom also carries a negative charge, but this is not a requirement. Likewise, we can identify the acid because it is the molecule that has acidic protons (hydrogens that carry a strong partial positive charge). The following are examples of other acids and bases. The acidic protons are shown in red, and the basic atoms in blue. Keep in mind that the concept of acid or base is always relative in the Bronsted theory. Some molecules such as water can act as acids or as bases, depending on who they are reacting with. We will expand on this later. − + − O δ + − O δ δ + δ δ δ O Bronsted acids − + H F H3C δ δ H H H O S O H O H O (CH3COOH) (H2SO4) H H Bronsted bases N H N H O H3C O CH3 O H H H H THE SCALE OF ACIDITY: VALUES pKa - Many acid-base reactions take place in water, one of the most universal solvents. Water also has the dual capability of acting as a proton donor or as a proton acceptor. It makes sense, then, to develop a scale of acidity based on the behavior of the substance of interest towards water. Since most acid-base reactions are equilibrium processes, the equilibrium constant of the reaction between an acid (or base) and water forms the basis for the pKa scale. Most general and organic chemistry textbooks contain adequate discussions of this parameter. What is of interest to us here is the relationship between pKa and acidity. The equation pKa = -log Ka = log(1/Ka) shows that such relationship is inverse. The stronger the acid (i.e. the higher its acidity constant Ka), the lower its pKa value, and viceversa. Tables of pKa values usually show the acids and their conjugate bases arranged by order of decreasing (or increasing) acidity. Chemistry students should become proficient at reading and using data presented in such tables. PREDICTING EQUILIBRIUM IN ACID-BASE REACTIONS - Equilibrium in acid-base reactions always favors the weaker side. In the following example the pKa values for the substances acting as acids are shown under their structures. Equilibrium favors the left side because the substances on the left are the weaker acid and the weaker base. CN + NH3 HCN + NH2 38 9.2 weaker stronger acid acid We can arrive at the same conclusion looking at the bases. The strength of bases is measured by the pKa of their conjugate acids. To understand how this works, we must remember that the relative strengths of the acid and the base in a conjugate pair hold an inverse relationship. The stronger the acid, the weaker its conjugate base, and viceversa. Therefore, the relationship between the pKa of an acid and the strength of its conjugate base is direct: the stronger the base, the higher the pKa value of its conjugate acid. We can look at the same reaction again, except that now we’re focusing on the bases, and arrive at the same conclusion that equilibrium favors the left side. CN + NH3 HCN + NH2 pKa of pKa of conj. acid = 9.2 conj. acid = 38 weaker base stronger base The weaker acid and the weaker base are always on the same side. If you arrive at a different conclusion, something is not right. IDENTIFYING ACIDIC PROTONS - The most general principle ruling acid strength can be stated thus: strong acids have relatively stable conjugate bases. In general, the more stable the conjugate base, the stronger the acid. An important thing to remember is that stability and reactivity are inverse. The more stable a substance is, the less reactive it is, and viceversa. Therefore, another way of stating the rule above is by saying that strong acids have weak conjugate bases. HCl and H3O+ are strong acids. Accordingly, the corresponding conjugate bases, Cl- and H2O, are weak (very stable). Chloride ion is stable because the negative charge resides on a very electronegative atom. Water molecule is one of the most stable substances known. How do we know which proton is the most acidic in a molecule (such as acetic acid) that contains more than one type of proton? Remember that the higher the degree of positive character on the proton, the more acidic it is. Examination of a pKa table reveals some trends for acidic protons. The following guidelines can be used to predict acidity.
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