Acids and Bases

Acids and Bases Overview of Acids and Bases 3) The third factor for determining the acidity is called Induction. Acid Strength and pKa Choosing a Base to Deprotonate a Given Compound Induction or inductive effect is the transmission of charge through atoms. Arhrenius Acids and Bases: The Ka values of ethyl alcohol (1) and ethyl amine (2) are shown below: Electronegative atoms pull the negative charge and help to stabilize the conjugate base. Acids dissociate in water to form protons (H+) Bases dissociate in water to form (-OH) O O H+ HCl Examles: Na+ NaOH H Ka1 = Cl- + NH = 10-16 > Ka2 = OH H O Br OH = 10-38 NH2 pKa = 4.75 pKa = 2.85 (Deprotonate means remove the most acidic proton) O O Cl OH F OH pKa = 2.80 Suppose you need to deprotonate the following compound: First, find its pKa value. OH pKa = 2.65 pKa = 25 H Write the chemical equation of the compound reacting with a conjugate base (A-): -OH + The more electronegative, the more electron withdrawing - the stronger the acid. The alcohol is a stronger acid because its Ka is higher. Restrictions: 1) only in water, 2) only -OH is a base To simplify the numbers, -logKa is used, which is the pKa of the given acid: O Cl What about organic molecules and non-aqueous solutions? O pKa = - logKa Acids are proton donors. No restiction to being in water Bases are proton acceptors. No restriction to -OH being the base pKa1 = - logKa1 = 16 OH OH Brønsted Acids and Bases: OH OH Cl Cl + Base AH pKa must be > 25 ! C. A. C. B. Pick a compound from a pKa table with pKa > 25. Amines are used often (pKa = 38). Increasing Acidity So, the stronger the acid, the lower the pKa value. (It is the oppsoite for Ka.) A- + The pKa of the conjugate acid must be higher than 25, i.e weaker acid is formed. The closer the electronegative atom to the hydrogen, the stronger the acid. < pKa2 = - logKa2 = 38 H pKa = 25 Acid O Cl O H + Na+NH2- Na+ + NH2-H pKa = 25 + Base H H O + N H Cojugate Acid Acid O Below are pKa values of some common functional groups: This explains the great difference of pKa values of alkanes, alkenes, and alkynes. OH Cojugate Base H 51 Curved-arrow mechanism of acid-base reactions: H 44 H 25 1) Start the arrow from the lone pair/negative charge of the base OH 19 OH 16 10 HCl 4.75 H pKa = 25 -7 Increasing Acidity and attack the most acidic hydrogen of the acid. H + H SP - 50 % S character + H SP2 - 33 % S character + H SP3 - 25 % S character pKa = 44 2) Break the bond of the hydrogen and move the electrons to the other atom. What makes the Acid Strong ? H Remember - curved arrows show movement of electrons. pKa = 50 The strength of the acid is dictated by the stability of the conjugate base that forms * Both Ahhrenius and Brønsted definitions are from Proton Prospective. when the acid dissociates. The stronger the acid, the more stable its conjugate base. + S orbital is more electronegative than P, so the more S character, the stronger the acid. Therefore, for acidity: alkynes > alkenes > alkanes A Lewis acid is an electron-pair acceptor. Lewis base is an electron-pair donor. Base O Stable A- means more acid dissociation, more proton, thus stronger acid. The priority order Atom, Resonance, Induction, Orbital for determining the acidity There are a few factors to be considered when assessing the stability of the conjugate base. of a compound is known as ARIO. 1) The first and the key factor is the Atom to which the hydrogen is connected. Mainly BH3 and bases as (1) OH O + H (2) NH2 NH + H Increasing Size As a general picture, visualilze acids as Increasing Electronegativity Oxygen is more electronegative, that is why the alcohol is more acidic: pKa1 = 16 Acid Strength Large atoms have higher polarizability and handle the negative charge better as the charge is spread over. This equation shows whether the compound with a given pKa is in the acidic or Aromiatic - stable 6 P electrons, flat, fully conjugated Non-Aromatic, cyclic hydrocarbon (2) NH2 NH + + H H Given equal concentration, whichever produces more H+ is the stronger acid. O + H SH + H S 1) If the pH < pKa, OH O H + N H log + Acid-Base reactions are reversible, so another way of asking the question is: pKa = 13 Is the equilibrium of the reaction shifted to the right (products) or to the left (reactants)? To answer this question, you need to remember that the equilibrium of an acid-base The sulfur atom is larger and stabilizes the negative charge better. So, you need to compare the pKa of the acid (left side) and the conjugate acid (right side). The equilibrium will be shifted to the higher pKa (weaker acid) side. 2) When comparing the same atom, check which conjugate base is Resonance-stabilized. Ka = H A HA The stronger the acid, the larger the Ka, Notice that Ka = Keq H2O OH O OH O +H pKa1 = 16 O OH NH O +H Base pK a = 5 O Not resonance stabilized Resonance stabilized [HA] > [A-] 2) If the pH = pKa, log 3) If the pH > pKa, [HA] log [A-] < 0, [HA] = 0, [A-] [HA] = [A-] since log1 = 0 compound exists mainly in its acidic form [A-] > [HA] compound exists mainly in its basic form N H Acid Conjugate Acid must be deprotonated + (basic form) to be Conjugate Base basic acidic For example, Aspirin O H + pKa = 10 O O [HA] > 0, [A-] (Acetylsalicylic acid) A = Conjugate base HA = Acid [HA] [A-] reaction is shifted towards the weaker acid (and the base). Thiols are more acidic even though oxygen is more electronegative than sulfur. The dissociation of the acid is a reversible reaction that is described by equilibrium constant: solution. pKa = pH + log NH pKa1 = 16 basic (deprotonated) form in an aqueous solution depending on the pH of that The Position of Equilibrium for an Acid–Base Reaction A OH O Li enolate Key points and implications of the Henderson-Hasselbalch equation: Generally speaking, the more the acid dissociates into a proton, the stronger it is. OH pKa = 60 H Is the following acid-base reaction possible? * When comparing elements in different rows, the atomic size is the key. (1) O + CH3CH2CH2CH2H Henderson-Hasselbalch Equation + H H H pKa = 60 H H pKa = 16 < pKa2 = 38 What does it mean "strong acid" ? + O pKa = 19 + H * Within the same row, atomic sizes don't change much, so electronagaticity is the key: * None of the Acid-Base theories contradict to one another. H pKa = 60 CH3CH2CH2CH2Li + So, in the Lewis theory, acids and bases are described in terms of electrons. HA CH3CH2CH2CH2H + ALi acid must be smaller than 60 - stronger acid reacts, weaker acid forms. Aromaticity also stabilizes the conjugate base: its atomic Size (Polarizability) and Electronegativity. Lewis adduct Acid CH3CH2CH2CH2Li + AH pK a must be < 60 pKa is 19 so it is a lot stronger of an acid than butane: Lewsis Acids and Bases: BH3 which is a very reactive base and needs to be quenched before being disposed. To do that, write the chemical equation of nBuLi with an acid (AH): Acetone is often used as a proton donor (acid) for protonating the nBuLi because its H HA + Suppose you need a reagent to protonate (quench) nBuLi (CH3CH2CH2CH2Li) The pKa of butane (the conjugate acid of nBuLi) is 60. Therefore, the pKa of AH What if there is not proton involved? O Choosing an Acid to Protonate a Given Compound O O H pKa = 38 4) The fourth factor is the Orbital of the atom connected to the hydrogen. More % S - stroger acid NH O O OH O O O O Aspirin acidic form + H O basic form physiologically active. pKa = 38 The pKa of aspirin (carboxylic acid) is ~5. Therefore, it is mainly in the acidic O In this case the pKa of the conjugate acid is greatly higher the acid's pKa, so the form hence inactive in stomach (pH = ~2), while in the environment of the equilibrium is shifted to the right. Therefore the right-pointing arrow is drawn bigger. cell (pH 7.4), the drug will be in its active basic form.