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AP Chemistry > Ch 16 acid base equilibria > Flashcards

Ch 16 acid base equilibria Flashcards

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1
Q

arrhenius concept of acids and bases

A

Acids are substances that, when dissolved in water, increase the concentration of H+ ions.
Bases are substances that, when dissolved in water, increase the concentration of OH- ions.

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2
Q

bronsted lowry concept of acids and bases

A

Bronsted-Lowry concept: acid-base reactions involve the transfer of H+ ions
To be a bronsted lowry acid molecule must have a H atom that it can lose as an H+ ion
To be a bronsted lowry base molecule must have nonbonding pair of electrons that it can use to bind H+ ion

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3
Q

amphiprotic

A

capable of acting as an acid when combined with something more basic than itself and as a base when combined with something more acidic than itself

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4
Q

conjugate acid-base pair2

A
  • every acid has a conjugate base formed by removing a proton from the acid. Every base has associated with it a conjugate acid, formed by adding a proton to the base
  • The stronger an acid, the weaker is its conjugate base; the stronger a base, the weaker its conjugate acid.
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5
Q

equilibrium in acid base reaction

A

position of the equilibrium favors the transfer of the proton to the stronger base. (favors reaction of stronger acid and stronger base to form weaker acid and weaker base) Equilibrium mixture contains more of the weaker acid and weaker base

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6
Q

autoionization of water

A

one water molecule can donate a proton to another water molecule. H2O+H2O→ H3O+ +OH-

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7
Q

ion product constant4

A
  • Equilibrium constant for autoionization of water= Kw=[H3O+][OH-]
  • Autoionization of water can also be written as H2O→ H+ +OH-, so Kw also equals [H+][OH-].
  • At 25ºC, Kw equals 1.0*10-14
  • The Kw is useful because it can be applied to any aqueous solution. You can use it to find value of [H+] or [OH-] if the other is known
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8
Q

identify acidity based on Kw

A

Natural solution: [H+]=[OH-]. Acidic solution: [H+]>[OH-]. Basic solution:[H+]

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9
Q

pH 6

A
  • [H+] is expressed in terms of pH. pH= -log[H+]
  • The pH of neutral solution= -log( 1.0*10-14) =7.00
  • The pH decreases as [H+] increases. The pH of an acidic solution is less than 7 and of a basic solution is greater than 7.
  • A solution of pH 6 has 10 times the concentration of H+ as a solution of pH 7.
  • The pH of a solution of a weak acid is higher than the pH of a solution of a strong acid of same molarity
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10
Q

pOH 3

A

pOH= -log[OH-].
-log[H+] + ( -log[OH-]) = -logKw
pH+pOH=14.00 (at 25º)

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11
Q

most common strong acids

A

HCl, HBr, HI, HNO3, HClO3, HClO4 (

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12
Q

strong acids3

A
  • strong electrolytes: extist in solution entirely as ions
  • monoprotic acids in equilibrium because strong acid is entirely ionized and reaction lies entirely to the right
  • pH of the solution of a strong monoprotic acid is [H+]= original concentration of the acid
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13
Q

most common strong bases

A

ionic hydroxides of group 1A and heavy 2A: such as NaOH, KOH, Ca(OH)2
Ionic oxides, hydrides, and nitrides

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14
Q

acid dissociation constant

A
  • equilibrium constant of ionization of acid
  • Larger value of Ka=stronger acid (more tendency to ionize)
  • Ka for weak acid are usually between 10^-3 and 10^-10
  • to calculate pH from Ka: Express concentrations in terms of x (eq value of H): (x)(x)/(initial conc-x)=Ka
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15
Q

percent ionization

A

Percent ionization =concentration ionized/original concentration *100%

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16
Q

polyprotic acids

A
  • have more than one ionizable H atom. Acid dissociation constants are labeled Ka1 and Ka2 (Ka2 always refers to the equilibrium involving the removal of the second proton of a polyprotic acid)
  • It is always easier to remove the first proton from a polyprotic acid than to remove the second (Ka value become smaller as successive protons are removed)
  • As long as successive Ka values differ by a factor of 10^3 or more, you can basically just consider the Ka1 value when determining pH and stuff (treat as a monoprotic acid)
17
Q

special case H2SO4

A

H2SO4 is strong acid for first proton, so first ionization step lies completely to right. But it is polyprotic acid so second step is in equilibrium

18
Q

base dissociation constant

A

equilibrium constant in which a base reacts with H2o to form a conjugate acid and OH-

19
Q

types of weak bases2

A
  • Amines: contain N and look like ammonia. One or more of the N-H bonds in NH3 is replaced by a N-C bond. (NH2CH3)
  • Anions of weak acids: when NaClO dissociates, ClO- becomes weak base in water
20
Q

relationship between Ka and Kb

A

KaKb=Kw, equals 110^-14 at 25ºC

pKa +pKb =pKw =14 (pKa = -logKa)

21
Q

hydrolysis

A

ions able to react with water to generate H+ or OH-

22
Q

how do salts affect pH of solution 6

A
  • An anion that is the conjugate base of a strong acid (Br- from HBr) will not affect pH of solution
  • An anion that is conjugate base of a weak acid (CN- from HCN) will become a weak base and increase the pH
  • A cation that is a conjugate acid of a weak base (NH3CH3+) will become a weak acid and decrease the pH
  • The cations of group 1A and heavy cations from group 2A (Ca, Sr, Ba) (strong bases) will not affect pH
  • Other metal ions will become weak acids and decrease the pH
  • When a solution contains both conjugate base of a weak acid and conjugate acid of a weak base, the ion with the larger equilibrium constant Ka or Kb will have the greater influence on the pH
23
Q

how does chemical structure affect acid base behavior

A
  • polarity: A molecule containing H will transfer a proton only if the H-X bond is partially positive in H and partially negative in X
  • bond strength: Strong bonds are less easily dissociated than weaker ones (strong bond=weak acid)
  • stability:The greater the stability of the conjugate base X-, the stronger the acid
24
Q

binary acids3

A
  • contain H and one other element X
  • When X goes down a group/column of PT: bond strength decreases down a group and acidity increases.
  • Across a row: electronegativity increases and acidity increases
25
Q

oxyacids4

A
  • acids in which OH groups and possibly additional oxygen atoms are bound to a central atom
  • If the central atom is a metal, it will be an ionic compound and behave as a base
  • If central atom is a nonmetal, compound is either acidic or neutral.
  • As electronegativity of central atom increases, acidity increases
  • Strength of the acid increases as additional electronegative atoms bond to the central atom Y
26
Q

oxyacid rules2

A
  • For oxyacids that have the same number of OH groups and the same number of O atoms, acid strength increases with increasing electronegativity of the central atom Y.
  • For oxyacids that have the same central atom Y, acid strength increases as the number of oxygen atoms attached to Y increases.
27
Q

carboxylic acids

A
  • contain a carboxyl group, COOH
  • Additional oxygen atoms added to carboxyl group increase polarity and stabilize conjugate base
  • Conjugate base of a carboxylic acid (carboxylate anion) can exhibit resonance which stabilizes it further.
  • Acid strength also increases as number of electronegative atoms increases
28
Q

lewis concept of acids and bases

A
  • A Lewis acid is an electron pair acceptor and a Lewis base is an electron pair donor
  • Anything that is a bronsted lowry base is also a lewis base. Not all lewis acids are bronsted lowry acids because it can apply to compounds that do not have H+
29
Q

hydration

A
  • metal ions attract unshared electron pairs of water molecules which causes salts to dissolve in water. Hydration is a lewis acid-base interaction where the metal ion is a lewis acid and water is a lewis base.
  • Acid dissociation constants for hydrolysis reactions increase with increasing charge and decreasing radius of the ion

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