Acids and bases
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Acids and Bases. The Secret Life of the Proton. Properties of acids and bases. Properties that you probably shouldn’t test Taste: acids are sour, bases bitter Feel: Bases feel slippery, acids don’t Properties you can test Reaction with metals:

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Acids and bases

Acids and Bases

The Secret Life of the Proton


Properties of acids and bases

Properties of acids and bases

  • Properties that you probably shouldn’t test

    • Taste: acids are sour, bases bitter

    • Feel: Bases feel slippery, acids don’t

  • Properties you can test

    • Reaction with metals:

    • Acids react with active metals to release H2, bases don’t (corrosiveness)


Properties of acids and bases1

Properties of acids and bases

  • Bases turn organic materials into organic salts (soap)

  • Acids and bases conduct electricity in water solution (ionization)

  • Indicators change color in presence of acid or base

  • Acids and bases react together to form a salt and water (neutralization)


Definitions

Definitions

  • Self-ionization of water:

    H2O + H2O  H3O+ + OH–

  • Svante Arrhenius (1859-1927)

  • Acids dissociate in water to give H+ ion, bases give OH-

    H+ + H2O  H3O+


Problems with arrhenius definition

Problems with Arrhenius definition

  • Did not take into account solvent effects

  • Did not explain why NH3 is a base

  • Did not explain basic salts

  • H+ does not exist in water


Br nsted lowry definition

Brønsted/Lowry definition

  • Acids are proton donors, and bases are proton acceptors

    HCl + H2O  H3O+ + Cl-

    acid base

    CH3COOH + OH- CH3COO- + H2O

    acid base


Br nsted lowry definition1

Brønsted/Lowry definition

  • Some species can act as an acid or a base – amphiprotic (usually a negative ion with H in front)

    HCO3- + H3O+ H2CO3 + H2O

    base acid

    HCO3- + OH- CO3-2 + H2O

    acid base


Br nsted lowry definition2

Brønsted/Lowry definition

  • Water is also amphiprotic

    H2O + NH2- NH3 + OH-

    acid base

    H2O + C7H7SO3H  C7H7SO3- + H3O+

    base acid

  • Amphoteric substances also can act as either acids or bases, but not necessarily involving a proton


Br nsted lowry definition3

Brønsted/Lowry definition

  • Definition does not depend on water

    H- + NH3 NH2- + H2

    base acid

    C7H7SO3H + C2H5OH  C7H7SO3- + C2H5OH2+

    acid base


Br nsted lowry definition4

Brønsted/Lowry definition

  • Conjugate pairs

    • Any acid has a conjugate base; any base has a conjugate acid

      HF + H2O  H3O+ + F-

      acid base acid base

    • This is an ionization reaction


Br nsted lowry definition5

Brønsted/Lowry definition

C2H3O2- + H2O  HC2H3O2 + OH-

base acid acid base

  • This is a hydrolysis reaction, responsible for basicity of some salts.

  • Mono- and polyprotic acids

    • Monoprotic acids have one acidic proton, i.e. HCl, HNO3, HC2H3O2 (or CH3COOH).


Br nsted lowry definition6

Brønsted/Lowry definition

  • Polyprotic acids have two or more acidic protons, i.e. H2SO4, H3PO4

    • In neutralization reactions all of the acidic protons react

  • Anhydrides

    • Nonmetal oxides are acid anhydrides

      SO3 + H2O  H2SO4

      CO2 + H2O  H2CO3


Br nsted lowry definition7

Brønsted/Lowry definition

  • Organic acid anhydrides are condensation dimers of the parent acids (missing water)

    acetic anhydride

acetic acid


Br nsted lowry definition8

Brønsted/Lowry definition

  • Metal oxides are basic anhydrides

    Na2O + H2O  2NaOH

    CaO + H2O  Ca(OH)2

  • Metal oxides react with water to produce hydroxides


Lewis acids and bases

Lewis acids and bases

  • Lewis acids are electron acceptors, and Lewis bases are electron donors

  • Examples: NH3 – Lewis base because of lone pair of electrons

  • BF3 – Lewis acid because of electron deficit on boron, and electron w/d characteristics of fluorine

  • NH3 + BH3 NH3BF3

    base acid

  • Strongly associated complex formed


Lewis acids and bases1

Lewis Acids and Bases

  • Water can act as either because of its polarity

  • Hydrogen bonds are results of a Lewis acid/base reaction

  • Oxidation and reduction are Lewis processes

    Na +H2ONa+ + OH- +H2(g)

    base acid


Acid base strength

Acid/base strength

  • A strong acid or base dissociates completely in water

  • Examples: HCl, H2SO4, HClO4, HNO3, most mineral acids, soluble hydroxides such as NaOH

    HCl + H2O  H3O+ + Cl-

    NaOH  Na+ + OH-


Acid base strength1

Acid/base strength

  • A weak acid or base does not dissociate very much

  • Examples: HF, most organic acids, NH3

  • Has very little to do with corrosiveness or reactivity


Acid base strength2

Acid/base strength

  • Conjugate pairs

    • The stronger the acid, the weaker the conjugate base

      HCl + H2O  H3O+ + Cl-

      strong acid weak base strong acid weak base

      HF + H2O  H3O+ + F-

      weak acid weak base strong acid strong base


Acid ionization constants

Acid ionization constants

  • Ka is an equilibrium constant

  • H2O + HF  H3O++F-

  • Ka=[H3O+][F-]/[HF]

  • Water is left out because it is constant

  • Higher Ka means stronger acid

  • Typical Ka values: 10-4 to10-11 for weak acids.


Acid ionization constants1

Acid ionization constants

  • Degree of dissociation is related to Ka

  • Example: Find value of Ka for HCN if a 0.1M sample is 0.0063% ionized.

    Ka=[H3O+][CN-]/[HCN]

    [H3O+]=[CN-]=0.0063% of 0.1M=0.0000063M

    [HCN]=0.1-.00000630.1

    Ka=0.00000632/0.1=4x10-10


Acid ionization constants2

Acid ionization constants

  • Find % ionization of 0.15M carbonic acid, Ka=4.4x10-7

    Assume [H2CO3]0.15M in ionized acid

    Ka=4.4x10-7=

    [HCO3-][H3O+]/[H2CO3]=y2/0.15

    y=2.6x10-4M

    % ionization = 2.6x10-4M/0.15M =0.17%


Acid base nomenclature

Acid base nomenclature

  • Bases are usually metal hydroxides and are named as salts [i.e. KOH – potassium hydroxide, Fe(OH)3 – iron (III) hydroxide]

  • Nonmetallic bases are usually derivatives of ammonia (NH3) – example C6H5NH2 – aniline

  • Hydrides are basic (NaH, LiAlH4) and are named as salts


Acid base nomenclature1

Acid-base nomenclature

  • Binary acids (H + another element) are named with prefix “hydro-”, suffix

  • “-ic” followed by “acid”. Example – HCl is hydrochloric acid

  • Ternary acids (oxyacids) are named after the central polyatomic anion.

  • “-ate” “-ic” “-ite”  “-ous”

  • H2SO4 – sulfuric acid

  • HClO – hypochlorous acid


P h and p oh scale

pH and pOH scale

  • Measures acidity, which depends on [H3O+]

  • pH=-log[H3O+]

  • Scale runs from 0-14

  • pH 7=neutral

  • pH>7=base

  • pH<7=acid


P h and p oh scale1

pH and pOH scale

  • Range of possible [H3O+] is 14 orders of magnitude – too large to use anything but a log scale

  • Derived from equilibrium constant for the self ionization of water

    2H2O  H3O+ + OH-

    Keq=[H3O+][OH-]=Kw=1x10-14


P h and p oh scale2

pH and pOH scale

  • Neutral water

    [H3O+]=[OH-]=10-7

    so pH=7

  • pOH

    pOH=-log[OH-]

    • for any solution,

      pH + pOH=14

      (same as Kw=10-14)


Indicators

Indicators

  • Indicators are compounds containing acidic protons that have colored ion(s). They are used to identify endpoints in titrations or determine pH.


Ph of weak acid solutions

pH of Weak Acid Solutions

  • Tend to be higher than strong acids

  • Find pH of 0.25M HF (Ka=6.6x10-4)

  • Solution: Ka=6.6x10-4= [H3O+][F-]/[HF]

  • [H3O+]=y=[F-], [HF]=.25-y

  • 6.6x10-4=y2/(.25-y)


Ph of weak acid solutions1

pH of Weak Acid Solutions

  • Assume y<<0.25, so 0.25-y0.25

  • Then 6.6x10-4=y2/0.25

  • y=0.0128=[H3O+]

  • pH=1.9

  • 0.25M HCl – pH=0.6

  • 20x as much hydronium in the strong acid


Neutralization reactions and titration

Neutralization reactions and titration

  • Definition

    acid + base  salt + water

  • Only holds for Brønsted acids and bases

    HNO3+LiOH  LiNO3+H2O

    H2SO4+2KOHK2SO4+2H2O

  • Neutralization reaction is exothermic


Titrations

Titrations

  • Titration – use of neutralization reaction to measure amounts of acids and bases

  • Endpoint or equivalence point – point at which equal # of equivalents of acid and base are present

  • 1 equivalent = 1 mole of donated or accepted protons

  • HCl: 1 mole = 1 equivalent


Titrations1

Titrations

  • H2SO4 1 mole=2 equivalents

  • H3PO4 1 mole=3 equivalents

  • NaOH 1 mole=1 equivalent

  • Ca(OH)2 1 mole=2 eq

  • Normality – equivalents/liter

  • For monoprotic acids and bases with one hydroxide, molarity = normality


Titrations2

Titrations

  • For polyprotic acids and bases with more than one hydroxide (i.e. Ca(OH)2): normality = molarity x acidic protons (or hydroxides)

  • Titration equation:

    NaVa = NbVb

  • Equivalence point is determined by indicator color change or pH meter


Titrations3

Titrations

  • Standard solution – a prepared solution of known concentration used as a standard in titrations

  • Sample problem – a 0.125 M standard solution of HCl is used to titrate 65 mL of an unknown solution of calcium hydroxide. The titration requires 22.3 mL HCl. What is the concentration of the base solution?


Titrations4

Titrations

  • Solution:

    • At the endpoint, eq acid = eq base.

      eq acid = 0.0223L(0.125mol/L)1eq/mol =eq base

      = .00279eq

      concentration base =

      0.00279eq(mol/2eq)(1/.065L) = 0.022M


Titrations5

Titrations

  • Using the titration equation:

  • NaVa = NbVb

  • 0.125N(22.3) = x(65)

  • x = 0.043N

  • 0.043eq/L(mol/2eq) = 0.022M


Characteristics of titrations

Characteristics of titrations

  • Strong acid/strong base – product solution is neutral, reaction goes to completion

  • Strong acid/weak base – result is slightly acidic

    NH4OH + HCl  NH4Cl + H2O

  • NH4+ + H2O  H3O+ + NH3

  • NH4+is a good acid; Cl- is a lousy base.


Characteristics of titrations1

Characteristics of titrations

  • Weak acid/strong base – result is slightly basic

    HC2H3O2 + NaOH  NaC2H3O2 + H2O

    C2H3O2- + H2O  HC2H3O2 + OH-

  • Acetate is a decent base, but Na+ is a terrible acid.

  • Na+ + 2H2O  NaOH + H3O+


Characteristics of titrations2

Characteristics of titrations

  • Weak acid/weak base – reaction does not complete – equilibrium established. pH depends on the acid and base used.

  • Titration curves

    • Strong acid/strong base

    • pH stays relatively unchanged until endpoint is near, then a rapid change occurs


Titration curves

Titration curves

Titration of 50 ml of a strong acid HA 0.1000 M by NaOH 0.1000 M :


Titration curves1

Titration curves


Titration curves2

Titration curves

  • Titration of weak acids with strong bases

    • Endpoint tends to occur at pH>7 because of hydrolysis


Titration curves3

Titration curves

  • pH after equivalence point depends on concentration of base


Titration curves4

Titration curves

  • Shape of curve depends on Ka of acid – lower the Ka, the smaller the pH jump at equivalence


Titration curves5

Titration curves

  • Titration of weak bases with strong acids


Buffers

Buffers

  • Buffers are made from a weak acid and the salt of a weak acid – ex. HC2H3O2 + NaC2H3O2

    H2O + HC2H3O2 H3O+ + C2H3O2-

    high low high

  • Since concentrations of both the acid and conjugate base are high, addition of small amounts of strong acid or base have little effect on pH


Buffers1

Buffers

  • Addition of acid:

    H3O+ + C2H3O2- HC2H3O2 + H2O

  • Addition of base:

    OH- + HC2H3O2 H2O + C2H3O2-

  • pH of buffers is governed by the acid equilibrium equation for the weak acid.


Buffers2

Buffers

  • Buffers with equimolar amounts of weak acid and its salt: pH = pKa

  • Example: Find pH of a buffer that is 0.1M in acetic acid and 0.1M in sodium acetate. Ka for acetic acid is 1.8x10-5.

  • Solution:

    1.8x10-5 = [H3O+][C2H3O2-]/[HC2H3O2]


Buffers3

Buffers

  • Since acid and salt are equimolar,

    [C2H3O2-]=[HC2H3O2]

    and the quantities cancel, leaving

    Ka = [H3O+],

    so pH = pKa

    = -log(1.8x10-5) = 4.7


Buffers4

Buffers

  • Buffers with unequal amounts of salt and acid – the Henderson-Hasselbalch equation (for acid HA and anion A-)

  • pH = pKa + log[A-]/[HA]

  • Example: Find the pH of a buffer that is 0.1M in acetic acid and 0.5M in acetate.


Buffers5

Buffers

  • Solution:

  • pH = log1.8x10-5 + log(0.5/0.1) = 5.4


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