10: Acids & Bases Flashcards Preview

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Flashcards in 10: Acids & Bases Deck (46):

Arrhenius definition

acid: dissociates to form excess H+ in solution (HCl, HNO3, H2SO4, etc.)

base: dissociates to form excess OH- in solution (NaOH, Fe(OH)3, etc.)


Bronsted-Lowry definition

acid: species that donates H

base: species that accepts H+ (OH-, NH3, F-)

*not limited to aqueous solutions

*acids and bases always occur in conjugate acid-base pairs


Lewis definition

acid: e- pair acceptor (BF3, AlCl3, etc.)

base: e- pair donor 

*idea is that one species pushes a l.p. to form a bond with another

  • coordinate covalent bond formation
  • complex ion formation 
  • nucleophile-electrophile interactions


amphoteric species

one that reacts like an acid in a basic enviornment and like a base in an acidic enviornment 

H2O + B⇔ HB + OH-

H2O + HA ⇔ H3O+ + A-

  • hydroxides of certain metals (Al, Zn, Pb, Cr) are amphoteric 
  • species that can act as both oxidizing and reducing agents are also considered amphoteric



amphoteric species that can either gain or lose a proton (Bronsted-Lowry)

ex. HSO4- can gain or lose a proton to become SO42- or H2SO4

  • water, amino acids, bicarbonate and bisulfate are common examples



acid-base nomenclature

anions: -ide → hyro__ic

ex. F-: Flouride --> HF: hydroflouric acid

anion: -ite → ous acid

ex. ClO-: hypochlorite --> HClO: hypochlorous acid

anion: -ate → ic acid

ex. CO32-: carbonate --> HCO3-: carbonic acid

ex. PO43-: phosphate --> H3PO4 phosphoric acid


acid-base behavior of water... autoionization

H2O(l)+ H2O(l) ⇔ H3O+(aq) + OH- (aq)


water dissociation constant 

Kw= [H3O+][OH-] = 10-14 at 298 K

*only changed by temperature, like all other equilibrium constants

[H3O+] = [OH-] = 10-7 for pure water



p scale

p scale = negative log of value 

pH = -log[H+] = 1/log[H+]

pOH = -log[OH-] = 1/log[OH-]

because [H3O+][OH-]=10-14...

pH + pOH = 14 for water at 298 K

pH=7 is neutral at 298 K




strongs acids & bases

completely dissociate into their component ions in aqueous solutions

ex. NaOH (s) --> Na+ (aq) + OH- (aq)

pH = 14 because 1 M [OH-] from 1 M NaOH and

pH = 14-pOH = 14 - -log[OH-] = 14+ log[1] = 14

*when calculated concentration of OH- or H+ ions from dissociation of acid and base, must take into consideration autionization of water... unless concentration of acid/base if significantly greater than 10-7 M


strong acids

HCl, HBr, HI, H2SO4, HNO3, HClO4


strong bases

NaOH, KOH, other soluble hydroxides of group 1A metals


weak acids and bases

acids and bases that only partially dissociate in aqueous solutions

 HA (aq) + H2O (l) ⇔ H3O+ (aq) + A- (aq)

BOH (aq) ⇔ B+ (aq) + OH- (aq)


acid dissociation constant 

smaller Ka means weaker acid and consequentially the less it will dissociates 

  • weak acid if Ka < 1 M 

A image thumb

base dissociation constant 

smaller Kb means weak base and consequentially, the less it will dissociate

  • weak base if Kb<1
  • Kb = [B+][OH-]/[BOH]


conjugate acid-base pairs

HCO3- (aq) + H2O (l) ⇔ CO32- (aq) + H3O(aq)

  • CO32- is conjugate base of HCO3-
  • H3Ois conjugate acid of H2O

Ka = [CO32-][H3O+]/[HCO3-]

reverse reaction is...

CO32- (aq) + H2O (l) ⇔ HCO3- (aq) + OH(aq)

Kb for CO32- is [HCO3-][OH-]/ [CO32-]


acid-base reactivity in water ultimately reduces to acid-base behavior of water (amphoteric)

Ka, acid x Kb, conjugate base = Kw = 10-14

Kb, base x Ka, conjugate acid = Kw = 10-14

  • Ka and Kb are inversely related so conjugate of a strong acid/base is inert because it is almost completely unreactive
  • for weak acids and bases, use Ka and Kb to calculate concentration of ions at equilibrium using the x is small approximation is Ka and Kare sufficiently small



electronegative elements positioned near an acidic proton increase acid strength by pulling electron desnity out of bond holding the acidic proton


neutralization reaction

when acids and bases react with each other to form a salt in water... usually goes to completion

  • reverse reaction is hydrolysis



strong acid + strong base

HCl + NaOH → NaCl + H2O

products of a reaction between equal ocncentrations of strong and strong base are equimolar amounts of salt and water... solution is neutral (pH=7) and ions formed in the reaction will not react with water because they are inert conjugates



strong acid + weak base

HCl + NH3 → NH4Cl

  • no water formed becuause weak bases are not hydroxides
  • cation of salt is weak acid and will react with water solvent and reform some of weak base through hydrolysis

1. HCl + NH3 → NH4(aq) + Cl- (aq)

2. NH4(aq) + H2O (l) → NH3 (aq) + H3O+

  • pH falls below 7



strong base + weak acid

salt hydrolyzes to form hydroxide ions so pH>7

ex.CH3COOH (weak acid) + NaOH 

1. CH3COOH + NaOH → Na+ + CH3COO- + H2O

2. CH3COO+ H2O → CH3COOH + OH-



weak acid + weak base

pH of solution depends on relative strenghts of reactants...

ex. HClO + NH3 → NH4ClO


acid & base equivalent

acid equivalent: 1 mole of H+

base equivalent: 1 mole of OH- ions



each mole of acid/base liberates more than 1 acid/base equivalent

(aka polyprotic under Bronsted-Lowry definition)

ex. H2SO4 + H2O (l) → HSO4- + H3O+

HSO4- + H2O (l) ⇔ SO42- + H3O+

1 mole H2SO4 can produce 2 acid equivalents (2 moles H3O+)



indicates quantity of acid or base... ex. each mole of H3PO4 yields 3 moles of H3Oso for a 2 M H3PO4 solution, the normality is 6 N


gram equivalent weight

mass of compound that produces 1 equivalent...

measurement useful for acid-base chemistry

ex. for H2SO4 each mole of acid yields 2 acid equivalents so gram weight equivalent is molar mass of H2SO4 /2 = 98/2 = 48 grams


common polyvalent acids and bases

acids: H2SO4, H3PO4, H2CO3

bases: Al(OH)3, Ca(OH)2, Mg(OH)2



procedure used to determine the concentration of a known reactant in a solution... acid-base and ox-red titrations

  • performed by adding small amounts of titrant (solution of known concentration) to titrand (solution of unknown concentration) until equivalence point


acid-base equivalent points

equivalence point is reached when the number of acid equivalents present in the original solution equals number of base equivalents added or vice-versa

  • equivalence point is pH 7 for strong acid/base titration, but not always pH 7 for other titrations


equation to calculate unknown value of titrand

NaVa = NbVb


pH meter

plotting pH of unknown solution as a function of added titrant... one way ways to determine equivalence point in an acid-base titration




weak organic acids/bases that have different colors in protonated/deprotonated states

H - indicator (color 1) ⇔ H+ + indicator- (color 2)

*indicator must always be weaker acid/base than the acid/base being titrated

point at which the indicator changes to its final color is enpoint, not equivalent point but difference is negligible... choose indicator that has closest pKa value to equivalence point 


strong acid-strong base titration

  • equivalence point of titration is pH 7 
  • ex. NaOH titrated into a solution of HCl
    • good indicator would have pKa 8


weak acid-strong base titration

ex. NaOH titrated into CH3COOH

  • pH is low but higher than when strong acid is titrand
  • equivalent point will be pH>7
  • pH curve rises earlier on and less sudden rise at equivalence point


strong acid-weak base titration

ex. HCl titrated into NH3

  • titration curve starts out pH 10-12 and then drops gradually with addition of strong acid.... sharp drop at equivalent point
  • pH>7

*identify type of titration by identifiying starting position and equivalence point 


weak acid-weak base

initial pH is betwen 3-11 and there's a very shallow drop at equivalence point

*least effective type of acid-base titration


polyvalent acids & bases

ex. Na2CO3 with HCl has H2CO3 as ultimate product

  • multiple equivalence points indicate polyvalent titration... 
  • flat part of curve is buffer region
    • center of buffer region is sometimes called half-equivalence point because it occurs when half of given species has been protonated/deprotonated


buffer soltuion

mixture of weak acid and its salt ( conjugate base and a cation) or a mixture of weak base and its salt (conjugate acid and anion)

  • resists changes in pH when small amounts of acid/base are added



acid buffer 

acetic acid (CH3COOH) and its salt, sodium acetate (CH3COO-Na+)

ex. CH3COOH (aq) + H2O (l) ⇔ H3O+ CH3COO-

  • when a small amount of strong base (NaOH) is added, OH- ions from NaOH react with H3O+ and more acetic acid dissociates to restore equilbirium... weak acid component is neutralizing strong base 
  • when a small amount of HCl is added, Hions from the HCl react with the acetate ions to form acetic acid, which affects pH less than HCl



base buffer

ammonia (NH3) and its salt, ammonium chloride (NH4+Cl-)


bicarbonate buffer system

H2CO3 (carbonic acid) & its conjugate base, HCO3- (bicarbonate) from a weak acid buffer for maintaining pH of blood 

  • CO2 is one of the waste products in cellular respiration and it travels through bicarbonate buffer system
    • CO2(g) + H2O(l) ⇔ H2CO3 ⇔ H+ HCO3-


Henderson-Hasselbach Equation

used to estimate pH of pOH of a buffer solution

  • changing concentrations of acid & its conjugate base would not change pH, but would change buffering capacity: ability to which system can resist changes in pH


weak acid buffer solution

pH = pKa + log [A-]/[HA]

  • [A-] for conjugate base
  • [HA] for weak acid
  • when [conjugate base] = [weak acid], pH=pKa
    • occurs at half-equivalence points


weak base buffer solution

pOH = pKb + log [B+]/[BOH]

  • [B+] for conjugate acid
  • [BOH] for weak base
  • pOH=pKb when [conjugate acid]=[weak base]