Acids and Bases 2

acid-base indicators

(HIn) an indicator is an acid or a base that undergoes dissociation in a known pH range a valuable tool for measuring pH because its acid form and base form have different colors in solution HIn (acid form) <-> (H+) + (In-) (base form) low pH = high [H+] high pH = high [OH-] limited by temperature changes, distorted, dissolved salts in solution pH meter

ethanoic acid

CH3COOH

carbonic acid

H2CO3

sulfuric acid

H2SO4 strong

phosphoric acid

H3PO4

hydrochloric acid

HCl strong

nitric acid

HNO3 strong

acidic solution

[H+] is greater than [OH-] [H+] is greater than 1 x 10^-7 M

basic solution

[H+] is less than [OH-] [H+] is less than 1 x 10^-7 M also known as alkaline solutions

buffer

a solution in which the pH remains relatively constant when small amounts of acid or base are added a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts

Lewis acid

a substance that can accept a pair of electrons to form a covalent bond

amphoteric

a substance that can act as both an acid and a base ex. water

Lewis base

a substance that donate a pair of electrons to form a covalent bond

hydrogen ion (H3O+)

a water molecule that gains a hydrogen ion becomes positively charged

Arrhenius theory

acids are hydrogen containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution. bases are compounds that ionize to yield hydroxide ions (OH-) in aqueous solution acids: only the hydrogens in very polar bonds are ionizable-joins to a very electronegative element-dissolve in water-releases hydrogen ions because they are stabilized bases: the alkali metals react with water to produce solutions that are basic

monoprotic acids

acids that contain one ionizable hydrogen ex. nitric acid (HNO3)

triprotic acids

acids that contain three ionizable hydrogens ex. phosphoric acids (H3PO4)

diprotic acids

acids that contain two ionizable hydrogens ex. sulfuric acid (H2SO4)

Lewis theory

an acid accepts a pair of electrons during a reaction while a base donates a pair of electrons

Bronsted-Lowry theory

an acid is a hydrogen ion donor a base is a hydrogen ion acceptor

neutral solution

any aqueous solution in which [H+] and [OH-] are equal in pure water at 25˚C the equilibrium concentration of [H+] and [OH-] is 1 x 10^-7 M so they are equal in pure water if [H+] increases, [OH-] decreases for aqueous solutions the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0 x 10^-14

bases

bitter taste, slippery feel, will change color of an acid-base indicator, can be strong or weak electrolytes in aqueous solution ex. soap

strong acids

completely ionized in aqueous solution [H3O+] is high ex. hydrochloric and sulfuric acid

salts

compounds consisting of an anion from an acid and a cation from a base

conjugate acid-base pair

consists of two substances related by the loss or gain of a single hydrogen ion

strong bases

dissociate completely into metal ions and hydroxide ions in aqueous solution ex. calcium hydroxide, magnesium hydroxide - not very soluble in water

strong/weak

extent of ionization or dissociation how many particles ionize or dissociate into ions

concentrated/dilute

how much dissolved in solution number of moles in a given volume

weak acids

ionize only slightly in aqueous solution [H3O+] is low ex. ethanoic - not complete

pH scale

ranges from 0-14 neutral solutions pH=7 ; [H+]=1x10^-7 0 = strongly acidic 14 = strongly basic pH of a solution is the negative log of the hydrogen-ion concentration pH = -log[H+] a solution in which [H+] is greater than 1 x 10^-7 M has a pH less than 7 and is acidic. the pH of pure water or a neutral aqueous solution is 7. a solution with a pH greater than 7 is basic and has a [H+] less than 1 x 10^-7 M PAGE 597 : if the [H+] is written in scientific notation and has the coefficient of 1, then the pH of the solution equals the exponent (+).

weak bases

react with water to form the hydroxide ion and the conjugate acid of the base ex. ammonia

neutralization reactions

reactions in which an acid and a base react in an aqueous solution to produce a salt and water a way to prepare pure samples of salts the reaction of an acid with a base produces water and one of a class of compounds called salts if a strong acid and strong base react > mixed with mole ratios of a balanced equation = neutral solution with properties of neither an acid or base reaction of weak acid and weak base do not usually produce a neutral solution

acids

tart, sour taste ex. lemons or vinegar aqueous solutions of acids are electrolytes-conduct electricity (strong/weak) they change the color of an acid-base indicator

buffer capacity

the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs too much base = no more acid molecules are present to donate hydrogen ions human blood (pH 7.35-7.45)

salt hydrolysis

the cations or anions of a dissociated salt remove hydrogen ions from or donate hydrogen ions to water in solutions they may be acidic or basic usually derived from a strong acid and weak base or a weak acid and a strong base ex. soap (strong base/weak acid) salts that produce acidic solutions contain positive ions that release protons to water. salts that produce basic solutions contain negative ions that attract protons from water hydrolysis-splits a hydrogen ion off a water molecule, the resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration-the solution is basic also, if strong enough, acid can donate a hydrogen ion to a water molecule [H3O+] greater than [OH-] therefore acidic

pOH

the negative log of the hydroxide ion concentration pOH = -log[OH-] neutral solution has a pOH of 7 pOH less than 7 is basic pOH greater than 7 is acidic pH + pOH = 14

conjugate acid

the particle formed when a base gains a hydrogen ion

conjugate base

the particle that remains when an acid has donated a hydrogen ion

end point

the point at which the indicator changes color in titration the point of neutralization - pH = 7

titration

the process of adding a known amount of solution of known concentration to determine the concentration of another solution 1. measure volume of acid 2. several drops of indicator 3. measured volumes of base added till slight color change titration with strong acid and base > pH changes, pH of acid starts low and increases as base is added b/c becomes neutralized

ion product constant for water (Kw)

the product of the concentrations of the hydrogen ions and hydroxide ions in water Kw = [H+] x [OH-] = 1.0 x 10^-7

acid dissociation constant (Ka)

the ratio of the concentration of the dissociated (or ionized) form of an acid to the concentration of the undissociated (non ionized) form Keq x H2O = Ka also known as ionization constants weak acids have small Ka values the stronger an acid is, the larger the Ka value - more complete each ionization reaction has its own separate dissociation constant (diprotic and triprotic)

base dissociation constant (Kb)

the ratio of the concentration to the conjugate acid times the concentration of the hydroxide ion to the concentration of the base Kb = (conjugate acid x [OH-]) / (base) the smaller the Kb value, the weaker the base

self ionization

the reaction in which water molecules produces ions if a water molecule loses a hydrogen ion-becomes negatively charged (OH-) if a water molecule gains a hydrogen ion- becomes positively charged (H3O+)

standard solution

the solution of known concentration

calculating the dissociation constants

to find the Ka of a weak acid or the Kb of a weak base, substitute the measured concentrations of all the substances present at equilibrium into the expression for Ka or Kb Ka = [H+][A-]/[HA]

equivalence point

when the number of moles of hydrogen ions equals the number of moles of the hydroxide ions when an acid and base are mixed