Chapter 10: Acids and Bases

if pH = 3 what is pOH

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Approximating x in acid base problems to determine pH or concentration of H3O+ also for pOH and OH-

A rule of thumb is that the approximation is valid as long as x is less than 5 percent of the initial concentration. This typically occurs when Ka is at least 100 times smaller than the concentration of the starting solution. For example, if Ka is 10−4 and the concentration of the starting solution is 0.01 M (10−2 M), then the ratio between the values is 102 or 100.

What species are considered the equivalents for acids and bases, respectively?

Acids use moles of H+ (H3O+) as an equivalent. Bases use moles of OH− as an equivalent.

conjugate acid base pairs can have a Ka and a Kb if one adds the reaction and its reverse together we get a net reaction of the autoionization (dissociation) of water we can use Ka of the conjugate acid to find Kb of the conjugate base and vice versa using Kw

Ka (acid) x Kb (CB)= Kw= 10^-14 Kb (base) x Ka (CA)= Kw= 10^-14

What is the mathematical relationship between Ka, Kb, and Kw?

Ka x Kb = Kw

Normality

M x n

Water dissociation constant (Kw)

concentration of each of the ions in pure water at equilibrium at 298 K is 10^-7 even if concentration is not equal Kw will always be 10^-14 at 298K Kw is an equilibrium constant. Unless the temp of the water is changed the value for Kw cannot be changed isolated changes in concentration, pressure or volume will not affect Kw

Conjugate acid base pairs

conjugate acid: the acid formed when a base gains a proton conjugate base: the base formed when an acid loses a proton

Arrhenius acids and bases

definitions are generally limited to aqueous acids and bases easily identified -acids contain H at the beginning of their formula (HCl, HNO3, H2SO4...) -bases contain OH at the end of their formula (NaOH, Ca(OH)2, Fe(OH)3...) most restrictive definition water is not considered an acid

Arrhenius Acid

dissociates to form an excess of H+ in solution

Arrhenius Base

dissociates to form and excess of OH- in solution

Lewis base

electron pair donor (lone pair)

Acid equivalent

equal to one mole of H+ (or H3O+) ions

Base equivalent

equal to one mole of OH- ions

Acid-Base Equivalence Points

equivalence point is reached when the number of acid equivalents present in the original solution equals the number of base equivalents added or vice versa while a strong acid/strong base titration will have its equivalence point at a pH of 7, the equivalence point does not always occur at pH 7 when titrating polyprotic acids or bases there are multiple equivalence points

partially dissociated conjugate vase of polyvalent acid is usually amphoteric

ex. HSO4- in acid will go to H2SO4 in base will go to SO42-

Oxyacid nomenclature

if the anion ends in -ite then the acid will end with -ous acid if the anion ends in -ate then the acid will end with -ic acid prefixes in the name of the anions are retained

Key concept: indicators

indicators change color as they shift between their conjugate acid and base forms Because this is an equilibrium process, we can apply Le Châtelier's principle. Adding H+ shis the equilibrium to the le. Adding OH− removes H+ and therefore shis the equilibrium to the right.

pH and pOH

logarithmic scales for the concentrations of hydrogen and hydroxide ions reactivity of an acidic solution is not a function of H+ concentration but on the logarithm of H+ concentration

Gram equivalent weight

mass of a compound that produces one equivalent (one mole of charge) ex. H2SO4 (molar mass 98 g/mol) is a divalent acid so each mole of the acid compound yields two acid equivalents the gram equivalent is 98/2=49 grams the complete dissociation of 49 g H2SO4 will yield one acid equivalent

if pH and pOH are 7

neutral, equal H+ and OH-

Peptide bond formation

neutralization reactions are often condensation reaction in bio and biochem because they form bonds with a small molecule as a byproduct (usually) water The peptide bonds in proteins, for example, are created from the reaction of a carboxyl group (acid) and an amino group (base), while forming a water molecule The salt in this reaction is the polypeptide itself; breaking it apart requires hydrolysis.

Bronsted-Lowry Acids and Bases

not limited to aqueous solutions ex. OH-, NH3, F- are bases because they can accept a hydrogen. -these would not be Arrhenius bases because they do not dissociate to produce an excess of OH- water is considered an acid --- acid and bases always come in pairs because the def requires a proton transfer -conjugate acid base pairs

Weak acids and bases

only partially dissociate in aqueous solutions weak acids and bases will dissociate partially to achieve an equilibrium state weak= ka or kb is less than 1.0

if H+ is .001 (10^-3)

pH = 3 -log(10^-3)=3

Weak acid and strong base reactions

pH will be within the basic range because the salt hydrolyzes with concurrent formation of hydroxide ions increase in OH- concentration will cause the system to shift away from autoionization, Redding the concentration of the hydronium ion. thus, the concentration of the hydroxide ion will be greater than that of the hydronium ion at equilibrium and the pH will rise above 7

Weak acid and weak base reactions

pH will depend on relative strengths of the reactants If Ka is less than Kb = basic and vice versa

p Scales

pOH and pH are examples p scales are defined as the negative logarithm of the number of items -log(____)

calculating the pH from a ___ M solution of acid/base when concentration is close to 10^-7 autoionization of water is not negligible

plug in to Kw=[H+][OH-]=14 if acid: (x+___)(x)

Strong acid and weak base reactions

product is also a salt but often no water will be formed because weak bases are often not hydroxides the cation of the salt is a weak acid and will react with the water solent reforming some of the weak base through hydrolysis ex. HCl + NH3 -) NH4+ + Cl- hydrolysis that occurs NH4+ + H2O -) NH3 + H3O+ The increase in the concentration of the hydronium ion causes the system to shi away from autoionization, thereby reducing the concentration of hydroxide ion. Consequently, the concentration of the hydronium ion will be greater than that of the hydroxide ion at equilibrium, and as a result, the pH of the solution will fall below 7. this reaction leads to a slightly acidic solution

Strong acid and strong base reactions

products of equal concentrations of each are equimolar amounts of salt and water acid and base neutrally each other so the resulting pH is 7, ions formed will not react with water as they are inert conjugates ex. HCl + NaOH -) NaCl + H2O

Acid dissociation constant (Ka)

because weak acids exist in an equilibrium state was can write the dissociation equation to determine Ka smaller Ka = weaker acid, less will dissociate weak= ka is less than 1.0

Strong acids and bases

completely dissociate into their component ions in aq solutions ex. NaOH will completely dissociate into 1M Na+ and 1M OH- pH of this would be pH=14-pOH=14-(-log[OH-])=14 + log(1M)= 14+0=14

What is an amphoteric species?

An amphoteric species can act as an acid or a base.

Compare and contrast the three definitions for acids and bases:

Arrhenius Acid: dissociates to from excess H+ in solution Arrhenius Base: dissociates to form excess OH- in solution ---- Bronsted-Lowry Acid: H+ donor Bronsted-Lowry Base: H+ acceptor ---- Lewis acid: electron pair acceptor Lewis base: electron pair donor

Lewis acid-Base chemistry

Boron trifluoride serves as the lewis acid, accepting a lone pair ammonia serves as a Lewis base donating a lone pair

Acetic Acid

CH3COOH

bicarbonate buffer system and conjugates

CO3-2 is a weak base HCO3- is its conjugate acid which is also weak the reaction of CO3-2 with water to produce HCO-3 and OH- occurs to a greater extent-is more thermodynamically factor- then HCO3- and water to get H3O+ and CO3-2 this makes this equilibrium ideal for buffering solutions a part of this system thermodynamic preference for the bicarbonate ion intermediate is a major reason why the bicarbonate buffer system in the body is ideal for maintaining a stable pH.

Acid strength and induction

EN elements positioned near an acidic proton increase acid strength by polling e- density out of the bond holding the acidic proton this weakens the proton bonding and facilitates dissociation acids that have EN elements nearer to acidic hydrogens are stronger than those that do not

Carbonic Acid

H2CO3

Chromic Acid

H2CrO4

Sulfuric Acid

H2SO4

Boric Acid

H3BO3

Phosphoric Acid

H3PO4

Strong acid examples

HCl. HBr, HI, H2SO4, HNO3 (nitric acid), HCLO4 (perchloric acid)

Hypochlorous acid

HClO

Chlorous Acid

HClO2

Chloric Acid

HClO3

Perchloric Acid

HClO4

Nitrous Acid

HNO2

Nitric Acid

HNO3

If a compound has a Ka value » water, what does it mean about its behavior in solution? How does this compare with a solution that has only a slightly higher Ka than water?

High Ka indicates a strong acid, which will dissociate completely in solution. Having a Ka slightly greater than water means the acid is a weak acid with minimal dissociation.

If a compound has a Kb value » water, what does it mean about its behavior in solution? How does this compare with a solution that has only a slightly higher Kb than water?

High Kb indicates a strong base, which will dissociate completely in solution. Having a Kb slightly greater than water means the base is a weak base with minimal dissociation.

Molecular (nonionic) weak bases

almost exclusively amines

Utilizing Arrhenius acid naming trends, predict the acid formula and name for the following anions:

MnO4-: -Acid formula: HMnO4 -Permanganic Acid Titanate (H2TiO3): -acid formula: H2TiO3 -titanic acid I-: -acid formula: HI -Hydroiodic acid IO4-: -acid formula: HIO4 -periodic acid

Acid-Base Nomenclature

Mose acids are named from their parent anions ( the anion that combines with H+ to form the acid). Acids formed from anions with names in -ide have the prefix hydro- and the ending -ic.

Strong base examples

NaOH, KOH, soluble hydroxides of group IA metals

approximating p scale values

One can obtain a relatively close approximation of a p scale value using the following shortcut: if the nonlogarithmic value is written in proper scientific notation, it will be in the form n × 10−m, where n is a number between 1 and 10. Taking the negative logarithm and simplifying, the p value will be: Because n is a number between 1 and 10, its logarithm will be a decimal between 0 and 1 (log1=0 and log10=1).The closer n is to 1, the closer logn will be to 0;the closer n Is to 10, the closer logn will be to 1. As a reasonable approximation, one can say that: pvalue=m-0.n where 0.n represents sliding the decimal point of n one position to the le (dividing n by ten).

Amino acid zwitterions are complex amphoteric species

amino group can release a proton (acid) and the carboxylate group can accept a proton (base)

Lewis acid

an electron pair acceptor (lone pair)

pH+pOH=14 for all solutions at 298 K

as pH increases, pOH decreases by the same amount... the opposite is true

NaVa=NbVb

at the equivalence point the number of equivalents of acid and base are equal this allows us to calculate the unknown concentration of the titrand N= normality V=volume (same units)

Four combos of strong and weak acids and bases are possible

Strong and Strong Strong and Weak Weak and Strong Weak and Weak

if pH is greater than 7 or pOH is less than 7

basic, more OH-

Key concepts

Water, amino acids, and partially deprotonated polyprotic acids such as bicarbonate and bisulfate are common examples of amphoteric and amphiprotic substances. Metal oxides and hydroxides are also considered amphoteric but not necessarily amphiprotic because they do not give off protons.

Titration

a procedure used to determine the concentration of a known reactant in a solution different kinds -acid-base -oxidation-reduction

Bronsted-Lowry Base

a species that accepts H+ ions

Bronsted-Lowry Acid

a species that donates H+ ions

Normality and acids/bases

acidic or basic capacity is directly indicated by the solutions normality ex mole of H3PO4 yields 3 moles of H3O+, therefore a 2M H3PO4 solution would be 6 N

if pH is less than 7 or pOH is greater than 7 at 298K

acidic, more hydrogen ions

Neutralization Reaction

acids and bases may react with each other to form a salt and often (but not always) water salt may precipitate out based on solubility reverse reaction where salt ions react with water to give back the acid or base is "hydrolysis"

Example of polyvalent acids and bases

acids: H2SO4, H3PO4, H2CO3 bases: Al(OH)3, Ca(OH)2, Mg(OH)2

Hydrolysis

reverse of neutralization reaction where salt ions react with water to give back the acid or base

Base dissociation constant (Kb)

smaller the Kb the weaker the base, the less it will dissociate weak= kb is less than 1.0

Polyvalent (polyprotic)

some acids and bases are these which means that each mole of acid or base liberates more than one acid or base equivalent ---- ex. diprotic acid H2SO4 one mole of H2SO4 produces two acid equivalents, 2 moles of H3O+ first dissociation goes to completion but second does not

Amphoteric species -amphiprotic

species that reacts like an acid in a basic environment and like a base in acidic environment in bronsted Lowry sense, an amphoteric species can either gain or lose a proton making it amphiprotic as well species that can act as both oxidizing and reducing agents are often considered amphoteric as they can accept or donate electron pairs

when Ka is large, Kb is small and vice versa

strong acid will have a Ka approaching and its conjugate base will be weak

Inert

the conjugate of a strong acid or base that is almost completely unreactive

Titrants and titrands

titrations are perfumed by adding small volumes of known concentration (titrant) to a known volume of a solution with unknown concentration (titrand) until completion of the reaction is achieved at the equivalence point

Determining the equivalence point

two common ways 1. graphical method plotting the pH of the unknown solution as a function of added titrant by using a pH meter 2. estimated by watching for a color change of an added indicator.

Lewis acids and Bases

underlying idea is that one species pushes a lone pair to form a bond with another (coordinate covalent bond formation) most inclusive definition every Arrhenius acid is also a Brønsted-Lowry acid, and every Brønsted-Lowry acid is also a Lewis acid (and likewise for bases). However, the converse is not necessarily true. The Lewis definition encompasses some species not included within the Brønsted-Lowry definition; for example, BF3 and AlCl3 are species that can each accept an electron pair, which qualifies them as Lewis acids, but they lack a hydrogen ion to donate, disqualifying them as both Arrhenius and Brønsted-Lowry acids. --- Lewis acids are often used as catalysts

Acid-Base behavior of water -amphoteric -autoionization

water is amphoteric water can react with itself in a process called auto ionization One water molecule donates a hydrogen ion to another water molecule to produce the hydronium ion (H3O+) and the hydroxide ion (OH−). autoionization of water is reversible for pure water at 298K the water dissociation constant Kw is Kw=[H3O+][OH-]=10^-14 at 25C (298 K)

Indicators and endpoint

weak organic acids or bases that have different colors in their protonated and deprotonated states This small structural change— the binding or release of a proton—leads to a change in the absorption spectrum of the molecule, which we perceive as a color change. Indicators are generally vibrant and can be used in low concentrations without significantly altering the equivalence point. The indicator must always be a weaker acid or base than the acid or base being titrated; otherwise, the indicator would be titrated first! The point at which the indicator changes to its final color is not the equivalence point but rather the endpoint. If the indicator is chosen correctly and the titration is performed well, the volume difference between the endpoint and the equivalence point is negligible and may be corrected for or simply ignored.