Chemistry: The Central Science Chapter 16: Acid-Base Equilibria 16.1: Acids and Bases: A Brief Review Acids have a sour taste and cause certain dyes to change color Base have a bitter taste and feel slippery During the 1880sSvante Arrhenius linked acid behavior with presence of H + and base behavior with the presence of OH - in aqueous solution o An acid is a substance that, when dissolved in water, increases the concentration of H + ions o A base is a substance that, when dissolved in water, increases the concentration of OH - ions 16.2: Brønsted Lowry Acids and Bases The Arrhenius concepts has the limitations of being restricted to aqueous solutions In 1923, Johannes Brønsted and Thomas Lowry proposed a more general definition of acids and bases o Their concept is based on the fact that acid-base reactions involve the transfer of H + ions from one substance to another The H + Ion in Water o An H + ion is simply a proton with no surrounding valence electron o H + react to the nonbonding electron pairs and form hydronium ion, H 3 O + Hydronium ion and further react with other water molecule, forming larger clusters of hydrated hydrogen ions, such as H 5 O 2 + and H 9 O 4 + Proton-Transfer Reactions o Brønsted and Lowry proposed definitions of acids and bases in terms of their ability to transfer protons An acid is a substance (molecule or ion) that donates a proton to another substance A base is a substance that accepts a proton o Brønsted-Lowry acid donates proton (H + ) o Brønsted-Lowry base accepts proton (H + ) o The concept of proton transfer also applies to reactions that do not occur in aqueous solution o Substances that are capable of acting as either acid or base is called amphiprotic
Acts as a base when combined with something more strongly acidic Acts as an acid when combined with something more strongly basic Conjugate Acid-Base Pairs o An acid and a base that differ only in the presence or absence of a proton are called a conjugate acid-base pair Every acid has a conjugate base, formed by removing a proton from the acid Every base has associated with it a conjugate acid, formed by adding a proton to the base o E.g. HX + H 2 O X - + H 3 O + HX donates a proton to H 2 O therefore HX is Brønsted-Lowry acid and H 2 O is the Brønsted-Lowry base In reverse reaction the H 3 O + donates a proton to X - ion, so H 3 O + is the acid while X - is the base When HX donates a proton, it leaves behind a substance X - which acts as a base o X - is the conjugate base When H 2 O accepts a proton, it forms a substance H 3 O + which acts as acid o H 3 O + is the conjugate acid Relative Strengths of Acids and Bases o The stronger an acid, the weaker is its conjugate base; the stronger a base, the weaker is its conjugate acid Strong acid completely transfers its protons to water; its conjugate base has a negligible tendency to abstract protons in aqueous solution A weak acid only partially dissociates in solution; the conjugate base of a weak acid is a weak base A substance with negligible acidity has strong base as its conjugate base o Proton-transfer reactions can be think of as a tug-of-war between two bases to abstract protons In every acid-base reaction the position of the equilibrium favors transfer of the proton from the stronger acid to the stronger base to form the weaker acid and the weaker base 16.3: The Autoionization of Water
One of the most important chemical properties of water is its ability to act as either a Brønsted acid or a Brønsted base. Water can donate a proton to another water molecule o This process is called the autoionization of water No individual molecule remains ionize for long The Ion Product of Water o Because the autoionization of water is an equilibrium process, its equilibriumconstant expression can be written K c = [H 3 O + ][OH - ] o Equilibrium constant for the autoionization of water is denote K w and is called the ion-product constant for water At 25 C, K w equals 1.0 x 10-14 o o A solution in which [H + ] = [OH - ] is neutral In acidic solutions [H + ] exceeds [OH - ] In basic solutions [OH - ] exceeds [H + ] 16.4: The ph Scale The molar concentration of H + in an aqueous solution is usually expressed in term of ph o ph = log[h + ] o Only number on the right of the decimal point are the significant figures in a logarithm The ph decreases as [H + ] increases o Change in [H + ] by a factor of 10 causes the ph to change by 1 o ph = 7 is neutral o ph < 7 is acidic o ph > 7 is basic If [H + ] is part of a kinetic rate law, then changing its concentration will change the rate o In biological systems many reactions involve protons transfers and have rates the depend on [H + ] Speeds of these reactions are crucial poh and Other p Scales o poh = log[oh - ] o ph + poh = 14.00 (at 25 C) Measuring ph o Can be measured with a ph meter, a device that generate voltage in solution
o It can also be measured using indicators, although less precise 16.5: Strong Acids and Bases Strong acids and bases are strong electrolytes, existing in aqueous solution entirely as ions Strong Acids o Seven most common strong acids: monoprotic acids (HCl, HBr, HI, HNO 3, HClO 3, and HClO 4 ) and diprotic acid (H 2 SO 4 ) o The equilibrium arrow is not used because the reaction goes to completion o ph of a solution of a strong monoprotic acid is in proportion to the molarity of the acid Strong Bases o Ionic hydroxides of the alkali metals and the heavier alkaline earth metals are strong bases o Strongly basic solutions are also created by certain substances that react with water to form OH - Most common of these contain the oxide ion 16.6: Weak Acids Weak acids partially ionized in aqueous solution E.g. OR o Subscript a denotes that it is an equilibrium constant for the ionization of acid, so K a is called that acid-dissociation constant The larger the value of K a, the stronger the acid Calculating K a from ph o In many cases the small magnitude of K a allow us to use approximations to simplify the problem o Proton-transfer reactions are generally very rapid The measured or calculated ph for a weak acid always represents an equilibrium condition Percent Ionization o A way to measure acid strength is percent ionization
o For any acid that the concentration that ionized equals the concentration of H + that forms, assuming that the autoionization of water is negligible Using K a to Calculate ph o K a can be used to calculate for ph by using the concepts from chapter 15 (such as I.C.E.) to calculate for [H + ] then for ph o As a general rule, if the quantity of x is more than about 5% of the initial value, it is better to use the quadratic formula o Properties of acid solution the relate directly to the concentration of H + are much less evident for a solution of weak acid than for a solution of a strong acid of the same concentration Rate of reaction with an active metal is much faster for strong acid than weak acid o The percent ionization decreases as the concentration increases The concentration of H + is not directly proportional to the concentration of the weak acid Doubling the concentration of a weak acid does not double the concentration of H + Polyprotic Acids o Polyprotic Acids acids that have more than one ionizable H atom o The ionization occurs in successive steps o It is always easier to remove the first proton from a poolyprotic acid than to remove the second The values of K a become successively smaller as successive protons are removed K a values for successive losses of protons form these acids usually differ by a factor of at least 10 3 o Most of the H + in the solution comes from the first ionization reaction As long as successive K a values differ by a factor of 10 3 or more, it is possible to obtain a satisfactory estimate of the ph of polyprotic acid solutions by treating them as if they were monoprotic acids, considering only K a1 16.7: Weak Bases Weak base reacts with water, forming the conjugate acid and OH - ions E.g. o Equilibrium-constant expression is
K b to denote that this equilibrium constant is for the ionization of weak base; the constant is called the base-dissociation constant The constant K b always refers to the equilibrium in which a base reacts with H 2 O to form the corresponding conjugate acid and OH - Base need to contain one or more lone pairs of electrons because a lone pair is necessary to form the bond with H + Types of Weak Bases o Weak bases fall into two categories Neutral substances that have an atom with a lone pair that can serve as proton acceptor Most of these bases contain nitrogen atom These substances include ammonia and a related class of compounds called amines Weak base that consists of anions of weak acids (basically, the conjugate base) 16.8: Relationship between K a and K b The product of the acid-dissociation constant for an acid and the base-dissociation constant for its conjugate base equals the ion-product constant for water o 16.9: Acid-Base Properties of Salt Solutions Salt solutions can be acidic or basic o The acid-base properties of salt solutions are due to the behavior of their constituent cations and anions Many ions are able to react with water to generate H + or OH - ; the reaction called hydrolysis o The ph of an aqueous solution can be predicted qualitatively by considering the ions of which the salt is composed An Anion s Ability to React with Water o In general, an anion, X -, in solution can be considered the conjugate base of an acid o Anions that still have ionizable protons, such as HSO 3 -, are amphiprotic A Cation s Ability to React with Water
o Polyatomic cations whose formulas contain one or more protons can be considered the conjugate acid of the weak bases o Most metal ions can also react with water to decrease the ph of an aqueous solution Ions of alkali metals and of the heavier alkaline earth metals do not react with water and therefore do not affect the ph Combined Effect of Cation and Anion in Solution o To summarize, anion: An anion that is the conjugate base of a strong acid will not affect the ph An anion that is the conjugate base of a weak acid will cause an increase in ph o Cation: A cation that is the conjugate acid of a weak base will cause a decrease in ph A cation that is the conjugate acid of a strong base will not affect the ph Other metal ions will cause a decrease in ph When a solution contains both the conjugate base of a weak acid and the conjugate acid of a weak base, the ion with the larger equilibrium constant, K a or K b, will have the greater influence on the ph 16.10: Acid-Base Behavior and Chemical Structure Factors That Affect Acid Strength o A molecule containing H will transfer a proton only if the H X bond is polarized in the way that the electron goes toward the X In the reverse of this, the H atom possesses a negative change and behaves as a proton acceptor Nonpolar H X bonds produce neither acidic or basic aqueous solution o Very strong bonds are less easily dissociated than weaker ones o The stability of the conjugate base, X - also affects the ease with which a hydrogen atom ionizes from HX The greater the stability of the conjugate base, the stronger is the acid Binary Acids o The bond strength decreases and the acidity increases down a group o Acidity increases as the electron negativity increases, as it generally does moving from left to right Oxyacids
o Acids in which OH groups and possibly additional oxygen atoms are bound to a central atom are called oxyacids o OH group can be present in both acid and base If the central atom is a metal, the electronegativity of O pulls the electron from the central atom, forming OH - and behave as base When the central atom is nonmetal, the bond to O is covalent and the substance does not readily lose OH - Generally, as the electronegativity of the central atom increases, so will the acidity of the substance o Electron density is drawn toward the central atom, the O H bond becomes weaker and more polar, thereby favoring the loss of H + o Because the conjugate base is usually an anion, its stability generally increases as the electronegativity of the central atom increases The strength of an acid will increase as additional electronegative atoms bond to the central atom o It pull electron density from the O H bond o It also helps stabilize the conjugate base by increasing its ability to spread out its negative charge o For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom o For oxyacids that have the same central atom, acid strength increases as the number of oxygen atoms attached to the central atom increases o In series of oxyacids, the acidity increases as the oxidation number of the central atom increases Carboxylic Acids o Carboxyl group is often writted as COOH, with one double C O bond and a single C OH bond Acids that contain a carboxy group are called carboxylic acids, and they form the largest category of organic acids o Two factors contribute to the acidic behavior of carboxylic acids Addition of oxygen atom attached to the carbon atom Conjugate base of carboxylic acid can exhibit resonance, which contributes further to the stability of the anion by spreading the negative charge over several atoms
16.11: Lewis Acids and Bases For a substance to be a proton acceptor, it must have an unshared pair of electrons for binding the proton G. N. Lewis proposed a definition of acid and base that emphasized the shared electron pair: o A Lewis acid is an electron-pair acceptor, and a Lewis base is an electron-pair donor o A Lewis base is an electron-pair donor In the Lewis theory, a base can donate its electron pair to something other than H + o The Lewis definition therefore greatly increases the number of species that can be considered acids o Lewis theory allows us to treat a wider variety of reactions as acid-base reactions o Metals such as Fe 3+ have vacant orbitals that accept the electron pairs Some compounds with multiple bonds can behave as Lewis acids Hydrolysis of Metal Ions o Because metal ions are positively charged, they attract the unshared electron pairs of water molecules The interaction, referred to as hydration, causes salts to dissolve in water o When a water molecule interacts with the positively charged metal ion, electron density is drawn from the oxygen This flow of electron density causes the O H bond to become more polarized E.g. Acid-dissociation constants for hydrolysis reactions generally increase with increasing charge and decreasing radius of the ion