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