LeChatelier’s Principlechemistry.bd.psu.edu/jircitano/LeChatelier.pdf1 Henri LeChatelier French...

1

Henri LeChatelier

French

1884

if a system at equilibrium is disturbed,

reaction shifts to oppose disturbance

(+)

(–)

heat +

+ heat

R P

LeChatelier’s Principle

ΔH + :

ΔH – : A + B C + D + heat

heat + E + F G + H

T

P N2O4(g) 2 NO2(g)

R P

gas moles

heat as P or R

1.

3.

2.

N2(g) + 3 H2(g) 2 NH3(g)

H2(g) + I2(g) 2 HI(g)

no shift from s or ℓ

P(g) vs R(g)

K changes:

Co(H2O)62+ + 4 Cl– CoCl4

2– + 6 H2O

K2Cr2O7 + K2O 2 K2CrO4

Pgas

a. To increase the yield of H2, should the pressure be high or low?

P side (H2) more moles gas (2:1); low P favors more moles

H2 increase

b. To increase the yield of H2, should the temperature be high or low?

T reaction goes in direction to consume heat () (heat as R)

H2 increase

c. What will adding carbon do to the yield of H2?

carbon is a solid no change;

d. What effect will adding Ni metal have on the yield of H2, knowing

that Ni reacts with CO to give Ni(CO)4(ℓ)?

[CO] reaction shifts to make more ()

DHo = +131.3 kJ

C(s) + H2O(g) CO(g) + H2(g) heat +

; H2 increase

2

Chapter 16

Acids and Bases

Base Base

Acid Acid

+ NH3(g)

substance that [OH–] in water

, react with acids

ancient acid:

base: litmus blue , bitter , slippery

, react with bases litmus red , sour , corrosive

Arrhenius acid:

base:

H2O(ℓ) H+(aq) + OH–(aq) connected by:

substance that [H+] (protons) in water

salts

CH3O

N

H

HO

N

quinine

acid:

HCl(g) Cl– + NH4+

+ – +

Brønsted - Lowry

+ –

proton donor (to base)

proton acceptor (acid) base:

Johannes Thomas

Svante

1923

1884

X– M+

Acid – Base Definitions

MX salt

+

–

Gilbert N. acid:

base:

e– pair acceptor (from a base)

e– pair donor (to an acid) 1923

Lewis • • H N H

H

+

– F B F

F

• •

• •

• •

• • • •

• •

• • • •

• •

3

+

Arrhenius:

HCl(aq) H+(aq) + Cl–(aq)

NaOH(aq) OH–(aq) + Na+(aq)

HCl(aq) H3O+(aq) + H2O(ℓ) + Cl–(aq)

acid hydronium ion

NH3(aq) ?

+ –

Brønsted – Lowry:

HCl, NaOH, NH3

base

base – OH–(aq) H2O(ℓ) + H2O(ℓ) + OH–(aq)

acid

NH3(aq) NH4+(aq) + H2O(ℓ) + OH–(aq)

base

+ – acid

+

–

+

hydroxide ion

HA A– + B HB+

H+ transferred both ways:

acid base

base + acid

base acid

acid + base forms

+

General Brønsted-Lowry Acid-Base Reaction

conjugate

base

conjugate

acid

– base

+ acid

conjugate acid-base pairs

+ – +

4

+

HA(aq) + H3O+(aq) + H2O(ℓ) A–(aq)

acid base

B(aq) HB+(aq) + H2O(ℓ) + OH–(aq)

acid base basic

Aqueous Solutions of Pure Acids and Bases

acidic

– +

+

– +

HA(aq) A–(aq) + H+(aq)

not

B-L definition and reality

amphiprotic (acid and base)

K Kw

H2O(ℓ) + H3O+(aq) + H2O(ℓ) OH–(aq)

H2O H2O [OH–] [H3O+]

initial

Δ

equil

0 0

+x +x

– –

– –

– – x x

x = 1.0 x 10–7 M [H3O+] = [OH–]

= x2 = [H3O+][OH–]

pure water

pure water

Autoprotolysis of Water

autoprotolysis constant

(25 oC)

B + H2O HB+ + OH–

HA + H2O A– + H3O+

–

+

+

=

= 1.0 x 10–14

HA(aq) + H3O+(aq) + H2O(ℓ) A–(aq)

B(aq) HB+(aq) + H2O(ℓ) + OH–(aq)

5

pure water

neutral

[H3O+] > [OH–] [OH–]

acidic

[H3O+] < [OH–] [H3O

+]

basic

[H3O+] = 1.0 x 10–6 M

[OH–] = 1.0 x 10–8 M

acidic

Arrhenius/Brønsted-Lowry Connection

acid

base

(1.0 x 10–6)[OH–] = 1.0 x 10–14

< 1.0 x 10–6 M [H3O

+] = [OH–]

= 1.0 x 10–14 Kw = [H3O+][OH–]

B + H2O HB+ + OH–

HA + H2O A– + H3O+

true for any aqueous solution

= 1.0 x 10–14 Kw = [H3O+][OH–]

= 14.00

= 1.0 x 10–14

= 1.0 x 10–14

6.00

7.00

8.00

1.0 x 10–8

1.0 x 10–7

1.0 x 10–6

1.0 x 10–6

1.0 x 10–7

1.0 x 10–8

0.00 1.0 x 10–14

1.00

2.00

3.00

1.0 x 10–13

1.0 x 10–12

1.0 x 10–11

1.0 x 10–0

1.0 x 10–1

1.0 x 10–2

1.0 x 10–3

4.00 1.0 x 10–10 1.0 x 10–4

5.00 1.0 x 10–9 1.0 x 10–5

9.00 1.0 x 10–5 1.0 x 10–9

10.00 1.0 x 10–4 1.0 x 10–10

11.00

12.00

13.00

14.00

1.0 x 10–3

1.0 x 10–2

1.0 x 10–1

1.0 x 10–0

1.0 x 10–11

1.0 x 10–12

1.0 x 10–13

1.0 x 10–14

pH = 7.00

pH < 7.00

pH > 7.00

neutral

acidic

basic

x

= 1.0 x 10–14

x

x

[H3O+] pH [OH–]

pH

[H3O+] = 2.2 x 10–9 M

pH = 8.65 basic

pH = 5.35

[H3O+] = 10–5.35 = 4.5 x 10–6 M

Søren Sørensen

Danish

(1868-1939)

pH = –log[H3O+]

[H3O+] = 10–pH –log H+

+ pOH = pH pKw

–log(1.0 x 10–8) = 8.00

–log(1.0 x 10–7) = 7.00

–log(1.0 x 10–6) = 6.00

6

pH 0 1 2 3 4 5 6 7 8 9 10 11 12

alizarian yellow R

thymolphthalein

phenolphthalein

thymol blue (base)

phenol red

bromothymol blue

chlorophenol red

bromocresol green

methyl orange

bromophenol blue

thymol blue (acid)

methyl violet

pH can be measured with indicators

Measuring pH

Universal indicator

thin glass membrane

solution of known pH

Ag/AgCl reference

Pt conductor

voltage pH difference

and pH electrodes/meter

H+ H+

H+

H+

H+

H+

H+ H+

H+

H+ H+ H+

H+ H+ H+

H+

H+

H+

H+

H+

H+

H+

H+

H+

H+

H+ H+

H+

H+

H+

H+ H+

H+

H+

H+

H+

H+

H+ H+ H+

H+ H+

H+

H+

H+

H+

H+

H+ H+

H+

NH3 + H2O NH4+ + OH– N2H4 + H2O N2H5+ + OH– C5H5N + H2O C5H5NH+ + OH–

NH4+ + H2O NH3 + H3O

+ H3PO3 + H2O H2PO3– + H3O

+ HC2H3O2 + H2O C2H3O2– + H3O

+ HCl + H2O Cl– + H3O+ H2SO4 + H2O HSO4

– + H3O+ HC7H5O2 + H2O C7H5O2

– + H3O+

HCl hydrochloric acid H2SO4 sulfuric acid

H3PO3 phosphorous acid HC2H3O2 acetic acid

NH4+ ammonium ion HC7H5O2 benzoic acid

(CH3)3N trimethylamine CH3O– methoxide ion

≥ one H+, often first

neutral or + (and H)

≥ one lone-pair e– (Lewis)

neutral or –, often N or O

acid (HA):

base (B):

NH3 ammonia CO32– carbonate ion

C5H5N pyridine N2H4 hydrazine

Recognizing Acids and Bases

HA + H2O A– + H3O+

(CH3)3N + H2O (CH3)3NH+ + OH– CO32– + H2O HCO3

– + OH– CH3O– + H2O CH3OH + OH– B + H2O HB+ + OH–

7

Cl– chloride

NO2– nitrite

C2H3O2– acetate

hydro_ic

–ous

–ic

–ide

–ite

–ate

acid: CB:

C5H5N pyridine

N2H4 hydrazine

CH3NH2 methylamine CH3NH3+ methylammonium

C5H5NH+ pyridinium

HN2H4+ hydrazinium

CA:

HCl hydrochloric acid

HNO2 nitrous acid

HC2H3O2 acetic acid

base–ium

NH3 ammonia NH4+ ammonium

Recognizing CB and CA From Name

Kb

Ka

pH acid/base reaction [H3O+] K

Acid-Base Equilibrium and pH

HA + H2O A– + H3O+

B + H2O HB+ + OH–

>> 1

<< 1

strong

weak

SA SB

WB WA

K [H3O

+][A–]

[HA] =

[OH–][HB+]

[B] = K

([OH–])

8

complete ionization:

7 common SA: HCl, HBr, HI hydrohalic acids (not HF)

: one H+ strong H2SO4 sulfuric

HNO3 nitric

HClO3, HClO4 chloric, perchloric

Ka = [H3O

+][X–]

[HX] = large

HX + H2O H3O+ + X–

stoichiometry

SA

6 common SA:

0 WA

not SA WA HC2H3O2 acetic HNO2 nitrous HF hydrofluoric

stronger weaker

Ka = [H3O

+][A–]

[HA] = 10–2 to 10–12

Ka = [H3O

+][B]

[HB+] HB+ + H2O B + H3O

+ CA of WB:

H3C CH C

OH O

OH lactic HO C CH2

O

C CH2 C

O

OH

CH2

OH

C O

OH

citric

2 Na+(aq) + O2–(aq) Na2O(aq) 2 Na+(aq) + O2–(aq)

OH– + OH– O2– + H2O 2 OH–

Na2O(aq)

O2– + H2O

NaOH(aq) Na+(aq) + OH–(aq)

2 – –

+

SB

OH–

O2– 2 types SB: ionic

stoichiometry

WB

H H

H

R’ R

R’’

CH3NH2 methylamine

(CH3)2NH dimethylamine

(CH3)3N trimethylamine

NH3

amines

NH2(CH2)4NH2 putracine

N

Kb = [OH–][HB+]

[B] = 10–2 to 10–12

Kb = [OH–][HA]

[A–] HA + OH– A– + H2O CB of WA:

Anthony Ulinski

Na+ (Mn+) not A/B

pyridine C

C N

C

C C

H

H

H H

H

9

6 SA: HCl HBr HI HClO3 HClO4 HNO3

Strong Acids

Weak Acids

if acid, but not SA WA:

Strong Bases

2 SB: ionic OH– ionic O2– (= 2 OH–)

Weak Bases

CA of WB (-ium)

CB of WA (-ide, -ite, -ate)

H-something

-NH2, -NH, -N- if base, but not OH–/O2– WB:

2. write acid-base reaction: CA of SB? pH = OH–

only SB “CA”

H2O 7.00 ?

, SB , WA , WB CB of SA , CB of WA , CA of WB

1. X in one of 7 categories:

pH of X(aq):

SA

3. Ka (WA, CA of WB) or Kb (WB, CB of WA)

pH Calculations

WB WA

,

A H3O+ or B OH–

10

6. (CH2OH)3CNH2

tris(hydroxymethyl)aminomethane

3. NaC7H5O2

sodium benzoate

1. HCl

2. HC8H7O3

mandelic acid

4. NaOH

CaO

7. NH3OH+Cl–

hydroxylammonium chloride

pH Calculations, 0.10 M Solutions of:

5. NaCl

HCl + H2O H3O+ + Cl–

1 HCl 1 H3O+

[H3O+] = [HCl]i = 0.10 M

SA acidic

0.10 M HCl

pH = –log(0.10) = 1.00

11

1. HCl

2. HC8H7O3

mandelic acid

7. NH3OH+Cl–


SA pH = 01.00

3. NaC7H5O2

sodium benzoate

4. NaOH

CaO

6. (CH2OH)3CNH2


5. NaCl


Dissociation Constants for Acids at 25 oC.

Name Formula Ka

Acetic HC2H3O2 1.8 x 10–5

Arsenous H3AsO3 5.1 x 10–10

Benzoic HC7H5O2 6.5 x 10–5

Butanoic HC4H7O2 1.5 x 10–5

Chloroacetic HC2H2O2Cl 1.4 x 10–3

Chlorous HClO2 1.1 x 10–2

Cyanic HCNO 3.5 x 10–4

Dichloroacetic HC2HO2Cl2 5.0 x 10–2

Formic HCHO2 1.8 x 10–4

Hydroazoic HN3 1.9 x 10–5

Hydrocyanic HCN 6.2 x 10–10

Hydrofluoric HF 6.8 x 10–4

Hypobromous HBrO 2.3 x 10–9

Hypochlorous HClO 3.0 x 10–8

Hypoiodous HIO 2.3 x 10–11

Iodic HIO3 1.7 x 10–1

Lactic HC3H5O3 1.4 x 10–4

Nitrous HNO2 4.5 x 10–4

Phenol HC6H5O 1.3 x 10–10

Propionic HC3H5O2 1.3 x 10–5

Pyruvic HC3H3O3 2.8 x 10–3

Thiocyanic HSCN 1.3 x 10–1

Dissociation Constants for Bases at 25 oC.

Name Formula Kb

Ammonia NH3 1.8 x 10–5

Aniline C6H5NH2 4.3 x 10–10

Dimethylamine (CH3)2NH 5.4 x 10–4

Ethylamine C2H5NH2 6.4 x 10–4

Hydrazine NH2NH2 1.3 x 10–6

Hydroxylamine NH2OH 1.1 x 10–8

Methylamine CH3NH2 4.4 x 10–4

Pyridine C5H5N 1.7 x 10–9

Trimethylamine (CH3)3N 6.4 x 10–5

Tris (HOCH2)3CNH2 1.2 x 10–6

0.10 M HC8H7O3

x = [H3O+]

x2 + (1.4 x 10–4)x – (1.4 x 10–5) = 0

Ka = (x)(x)

(0.10 – x) =

[H3O+][C8H7O3

–]

[HC8H7O3]

pH = –log(3.7 x 10–3) = 2.43

x = 3.7 x 10–3 M, –7.6 x 10–3 M

(assumption good)

; x = 3.7 x 10–3 M assume x << 0.10: x2

0.10 – x

x2

0.10 ≈ = 1.4 x 10–4

HC8H7O3 + H2O H3O+ + C8H7O3

–

WA acidic

Mandelic HC8H7O3 1.4 x 10–4

–

–

–

H2O

x x 0.10 – x equil

+x +x –x D

0 ~0 0.10 initial

[C8H7O3–] [H3O

+] [HC8H7O3]

= 1.4 x 10–4

12

1. HCl

2. HC8H7O3

mandelic acid

7. NH3OH+Cl–


SA pH = 01.00

WA pH = 02.43

3. NaC7H5O2

sodium benzoate

4. NaOH

CaO

6. (CH2OH)3CNH2


5. NaCl


HC7H5O2


Name Formula Ka











Name Formula Kb











0.10 M NaC7H5O2

Na+(aq) + C7H5O2–(aq)

+ OH– + H2O HC7H5O2

NaC7H5O2(aq)

basic CB of WA

CB/CA not in table + H3O+ + H2O C7H5O2

–

Kb = ?

Ka = 6.5 x 10–5

[H3O+][C7H5O2

–]

[HC7H5O2]

[OH–][HC7H5O2]

[C7H5O2–]

x = [H3O+][OH–] = Kw

KaKb = 1.0 x 10–14

C A/B pair

here, A stronger than CB: inverse relationship

= 1.5 x 10–10 Kb C7H5O2–

1.0 x 10–14

6.5 x 10–5 =

Kw

Ka HC7H5O2 =

0.10 M NaC7H5O2

C7H5O2–


+

13

–

–

–

H2O

x x 0.10 – x equil

+x +x –x D

0 ~0 0.10 initial

[HC7H5O2] [OH–] [C7H5O2–]

x = [OH–]

C7H5O2– + H2O OH– + HC7H5O2

x = 3.9 x 10–6 M = [OH–]

pH = 14.00 – 5.41 = 8.59

; pOH = 5.41

assume x small: x2

0.10 – x

x2

0.10 ≈ = 1.5 x 10–10

Kb (x)(x)

(0.10 – x) =

[OH–][HC7H5O2]

[C7H5O2–]

=

0.10 M NaC7H5O2

= 1.5 x 10–10

= 1.5 x 10–10

1. HCl

2. HC8H7O3

mandelic acid

7. NH3OH+Cl–


SA pH = 01.00

WA pH = 02.43

CB of WA pH = 08.59 3. NaC7H5O2

sodium benzoate

4. NaOH

CaO

6. (CH2OH)3CNH2


5. NaCl


14

[OH– ] = [NaOH]i = 0.10 M

pOH = 1.00 ; pH = 13.00

SB basic

NaOH Na+ + OH–

O2– + H2O 2 OH–

CaO Ca2+ + O2–

[OH– ] = 2 x [CaO]i = 0.20 M

pOH = 0.70 ; pH = 13.30

0.10 M NaOH; 0.10 M CaO

1. HCl

6. (CH2OH)3CNH2


7. NH3OH+Cl–


SA pH = 01.00

SB pH = 13.30

SB pH = 13.00

2. HC8H7O3

mandelic acid

WA pH = 02.43

5. NaCl


sodium benzoate

4. NaOH

CaO


15

Cl– + H2O HCl + OH–

KaKb = Kw

Ka large , Kb very small ; not basic

NaCl Na+ + Cl–

pH = 7.00

(pure water)

0.10 M NaCl

as Na+ and K+ salts

6 CB of SA: Cl–, Br–, I–, ClO3–, ClO4

–, NO3–

CB of SA (HCl)

1. HCl

6. (CH2OH)3CNH2


7. NH3OH+Cl–


SA pH = 01.00

2. HC8H7O3

mandelic acid

WA pH = 02.43

5. NaCl


sodium benzoate

4. NaOH

CaO SB pH = 13.30

SB pH = 13.00

CB of SA pH = 07.00


16


Name Formula Ka












Name Formula Kb










(CH2OH)3CNH2 + H2O OH– + (CH2OH)3CNH3+

WB basic

0.10 M (CH2OH)3CNH2

–

–

–

x x 0.10 – x

+x +x –x

0 ~0 0.10

x = [OH–]

; pOH = 3.46; pH = 10.54 x = 3.5 x 10–4 M = [OH–]

Kb (x)(x)

(0.10 – x) =

[OH–][RNH3+]

[RNH2] =

x2

0.10 ≈

H2O

equil

D

initial

[RNH3+] [OH–] [RNH2]


= 1.2 x 10–6

R

=

(tris(hydroxymethyl)aminomethane)

1. HCl

6. (CH2OH)3CNH2


7. NH3OH+Cl–


SA pH = 01.00

CB of SA pH = 07.00

WB pH = 10.54

2. HC8H7O3

mandelic acid

WA pH = 02.43

5. NaCl

SB pH = 13.30

SB pH = 13.00


sodium benzoate

4. NaOH

CaO

• • H N OH

H

+

H

H N OH

H


17


Name Formula Ka












Name Formula Kb










0.10 M NH3OH+Cl–

NH3OH+Cl– NH3OH+ + Cl–

Kw

Kb NH2OH Ka =

CB of SA not basic CA of WB

+ H2O H3O+ + NH2OH

–

–

–

x x 0.10 – x

+x +x –x

0 ~0 0.10

x = [H3O+]

; pH = 3.52 x = 3.0 x 10–4 M = [H3O+]

x2

0.10 ≈

(x)(x)

(0.10 – x) = Ka =

[H3O+][NH2OH]

[NH3OH+]

= 9.1 x 10–7 1.0 x 10–14

1.1 x 10–8 =

H2O

equil

D

initial

[NH2OH] [H3O+] [NH3OH+]


= 9.1 x 10–7

NH3OH+

WB pH = 10.54

1. HCl

7. NH3OH+Cl–


SA pH = 01.00

CA of WB pH = 03.52

CB of SA pH = 07.00

2. HC8H7O3

mandelic acid

WA pH = 02.43

SB pH = 13.30

SB pH = 13.00


sodium benzoate

4. NaOH

CaO

6. (CH2OH)3CNH2


5. NaCl


18


Name Formula Ka (or Ka1) Ka2 Ka3


Arsenic H3AsO4 5.8 x 10–3 1.1 x 10–7 3.2 x 10–12



Boric H3BO3 5.8 x 10–10 1.8 x 10–13 1.6 x 10–14


Carbonic H2CO3 4.3 x 10–7 5.6 x 10–11



Citric H3C6H5O7 7.4 x 10–4 1.7 x 10–5 4.0 x 10–7


Ka1 > Ka2 H2A + H2O HA– + H3O+ Ka1

HA– + H2O A2– + H3O+ Ka2

H+ donated stepwise: H2A, diprotic acid

Polyprotic Acids

Kb1 > Kb2

HA– + H2O H2A + OH– Kb2

A2– + H2O HA– + OH– Kb1

Ka1Kb2 = Kw

Ka2Kb1 = Kw

A2–

HA–

H2C6H6O6 + H2O H3O+ + HC6H6O6

–

; Ka2 = 1.6 x 10–12 Ka1 = 8.0 x 10–5

; pH = 2.55 ; x = 2.8 x 10–3 M H3O+ Ka1 = 8.0 x 10–5

x2

0.10 ≈

0.10 M H2C6H6O6, ascorbic acid

+ H2O HCO3– + OH–

H2CO3, carbonic acid: ; Ka2 = 5.6 x 10–11 Ka1 = 4.3 x 10–7

; pH = 8.63 ; pOH = 5.37 x = 4.3 x 10–6 M = [OH–]

= 1.8 x 10–4 ; Kb1 = 1.0 x 10–14

5.6 x 10–11 Ka2Kb1 = Kw

x2

0.10 ≈

0.10 M Na2CO3, sodium carbonate

Na2CO3 2 Na+ + CO32–

Ascorbic H2C6H6O6 8.0 x 10–5 1.6 x 10–12

LeChatelier’s Principlechemistry.bd.psu.edu/jircitano/LeChatelier.pdf1 Henri LeChatelier French...

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