chemistry 102(01) fall 2010

16-1CHEM 102, Fall 2010, LA TECH

Instructor: Dr. Upali Siriwardane

e-mail: [email protected]

Office: CTH 311 Phone 257-4941

Office Hours: M,W 8:00-9:00 & 11:00-12:00 am; Tu,Th,F 8:00 - 10:00 am.

Test Dates: September 23, October 21, and November 16; Comprehensive Final Exam: November 18, 2010

Exam: 10:0-10:15 am, CTH 328.

September 23, 2010 (Test 1): Chapter 13

October 21, 2010 (Test 2): Chapters 14 & 15

November 16, 2010 (Test 3): Chapters 16, 17 & 18

Comprehensive Final Exam: November 18, 2010 :Chapters 13, 14, 15, 16, 17 and 18

Chemistry 102(01) Fall 2010


Chapter 16. Acids and Bases

16.1 The Brønsted-Lowry Concept of Acids and Bases

16.2 Types of acids/bases:Organic Acids and Amines

16.3 The Autoionization of Water

16.4 The pH Scale

16.5 Ionization Constants of Acids and Bases

16.6 Problem Solving Using Ka and Kb

16.7 Molecular Structure and Acid Strength

16.8 Acid-Base Reactions of Salts

16.9 Practical Acid-Base Chemistry

16.10 Lewis Acid and Bases


Types of Reactionsa) Precipitation Reactions. Reactions of ionic compounds or

saltsb) Acid/base Reactions. Reactions of acids and basesc) Redox Reactions. reactions of oxidizing & reducing

agents


What are Acids &Bases?

Definition?

a) Arrhenius

b) Bronsted-Lowry

c) Lewis


Arrhenius, Svante August (1859-1927), Swedish chemist, 1903 Nobel Prize in chemistry

•Acid Anything that produces hydrogen ions in a water solution.

HCl (aq) H+

( aq) + Cl-

( aq)

•Base Anything that producs hydroxide ions in a water solution.

NaOH (aq) Na+

( aq) + OH

- ( aq)

•Arrhenius definitions are limited proton acids and hydroxide bases

to aqueous solutions.

Arrhenius Definitions


Expands the Arrhenius definitions to include many bases other than hydroxides and gas phase reactions

Acid Proton donor

Base Proton acceptor

This definition explains how substances like ammonia can act as bases.

Eg. HCl(g) + NH3(g) ------> NH4Cl(s)

HCl (acid), NH3 (base).

NH3(g) + H2O(l) NH4+

+ OH-

Brønsted-Lowry definitions


Lewis DefinitionG.N. Lewis was successful in including acid and bases

without proton or hydroxyl ions.

Lewis Acid: A substance that accepts an electron pair.

Lewis base: A substance that donates an electron pair.

E.g. BF3(g) + :NH3(g) F3B:NH3(s)

the base donates a pair of electrons to the acid forming a

coordinate covalent bond common to coordination

compounds. Lewis acids/bases will be discussed later in

detail


DissociationStrong Acids:

HCl(aq) + H2O(l) H3+O(aq) + Cl-(aq)

H2SO4(aq) + H2O(l) H3+O(aq) + HSO4

-(aq)

Dissociation Equilibrium Weak Acid/base:

H2O(l) + H2O(l) H3+O(aq) + OH-(aq)

This dissociation is called autoionization of water.

HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2

-(aq)

NH3 (aq) + H2O(l) NH4+ + OH-(aq)

Equilibrium constants: Ka, Kb and Kw


Conjugate acid-base pairs.

Acids and bases that are related by loss or gain of H+ as H3O+

and H2O.

Examples. Acid Base

H3O +

H2O

HC2H3O2 C2H3O2-

NH4 +

NH3

H2SO4 HSO4-

HSO4- SO4

2-

Brønsted-Lowry Definitions


Bronsted acid/conjugate base and base/conjugate acid pairs inacid/base equilibria


HCl(aq): acid

H2O(l): base

H3+O(aq): conjugate acid

Cl-(aq): conjugate base

H2O/ H3+O: base/conjugate acid pair

HCl/Cl-: acid/conjugate base pair


Select acid, base, acid/conjugate base pair,base/conjugate acid pair

H2SO4(aq) + H2O(l) H 3+O(aq) + HSO4

-(aq)acid

base

conjugate acid

conjugate base

base/conjugate acid pair

acid/conjugate base pair


Types of Acids and BasesBinary acids: HCl, HBr, HI, H2S

More than two elements: HCN

Oxyacid: HNO3, H2SO4, H3PO4

Polyprotic acids: H2SO4, H3PO4

Organic acids: R-COOH, R= CH3-, CH3CH2-

Acidic oxides: SO3, NO2, CO2,

Basic oxides: Na2O, CaO

Amine: NH3. R-NH2, R= CH3-, CH3CH2- : primary

R2-NH : secondary, R3-N: tertiary

Lewis acids & bases: BF3 and NH3


Strong Acid vs. Weak AcidsStrong acidcompletely ionizedHydrioidic HI Ka ~ 1011 pKa = -11Hydrobromic HBr Ka ~ 109 pKa = -9Perchloric HClO4 Ka ~ 107 pKa = -7Hyrdrochloric HCl Ka ~ 107 pKa = -7Chloric HClO3 Ka ~ 103 pKa = -3Sulfuric H2SO4 Ka ~ 102 pKa = -2Nitric HNO3 Ka ~ 20 pKa = -1.3Weak acidpartially ionizedHydrofluoric acid HF Ka = 6.6x10-4 pKa = 3.18Formic acid HCOOH Ka = 1.77x10-4 pKa = 3.75Acetic acid CH3COOH Ka = 1.76x10-5 pKa = 4.75Nitrous acid HNO2 Ka = 4.6x10-4 pKa = 3.34Acetyl Salicylic acid C9H8O4 Ka = 3x10-4 pKa = 3.52Hydrocyanic acid HCN Ka = 6.17x10-10 pKa = 9.21


Strong Base vs. Weak BaseStrong Basecompletely ionizedLithium hydroxide LiOHSodium hydroxide NaOH

Potassium hydroxide KOH Kb~ 102-103

Rubidium hydroxide RbOHCesium hydroxide CsOHBoarder-line Bases

Magnesium hydroxide Mg(OH)2

Calcium hydroxide Ca(OH)2

Strotium hydroxide Sr(OH)2 Kb~ 0.01 to0.1

Barium hydroxide Ba(OH)2

Weak Basepartially ionized

Ammonia NH3 Kb=1.79x10-5 pKb = 4.74

Ethyl amine CH3CH2NH2 Kb=5.6x10-4 pKb = 3.25


•Strong acids Ionize completely in water. HCl, HBr, HI,

HClO3,

HNO3, HClO4, H2SO4.

•Weak acids Partially ionize in water.

Most acids are weak.

•Strong bases Ionize completely in water.

Strong bases are metal hydroxides - NaOH, KOH

•Weak bases Partially ionize in water.

Acid and Base Strength


Common Acids and BasesAcids Formula Molarity*

nitric HNO3 16

hydrochloric HCl 12

sulfuric H2SO4 18

acetic HC2H3O2 18

Bases

ammonia NH3(aq) 15

sodium hydroxide NaOH solid

*undiluted.


Autoionization When water molecules react with one another to form ions.

Acids and bases alter the dissociation equilibrium of water based on Le

Chaterlier’s principle

Kw = [ H3O+

] [ OH-

]

= 1.0 x 10-14

at 25oC

Note: [H2O] is constant and is included in Kw.

ion product

of water

ion product

of water

H2O(l) + H2O(l) H3O+

(aq) + OH-(aq)

(10-7

M) (10-7

M)

Autoionization of Water


We need to measure and use acids and bases over a very large

concentration range.

pH and pOH are systems to keep track of these very large ranges.

pH = -log[H3O+

]

pOH = -log[OH-]

pH + pOH = 14

Substance pH

1 M HCl 0.0

Gastric juices 1.0 - 3.0

Lemon juice 2.2 - 2.4

Classic Coke 2.5

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45

Milk of Magnesia 10.5

Household ammonia 12.0

1M NaOH 14.0

Substance pH

1 M HCl 0.0



Classic Coke 2.5

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45



1M NaOH 14.0

pH and other “p” scales


A logarithmic scale used to keep track of the large changes in [H+

].

0 7

1410

-14 M 10

-7 M 10-14 M

Very Neutral Very

acidic Basic

When you add an acid to, the pH gets smaller.

When you add a base to, the pH gets larger.

pH scale


Substance pH

1 M HCl 0.0



Classic Coke 2.5

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45



1M NaOH 14.0

Substance pH

1 M HCl 0.0



Classic Coke 2.5

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45



1M NaOH 14.0

pH of some common materials


pH of Aqueous Solutions


What is pH?

Kw = [H3+O][OH-] = 1 x 10-14

[H3+O][OH-] = 10-7 x 10-7

Extreme cases:

Basic medium

[H3+O][OH-] = 10-14 x 100

Acidic medium

[H3+O][OH-] = 100 x 10-14

pH value is -log[H+]

spans only 0-14 in water.


pH, pKw and pOHThe relation of pH, Kw and pOH Kw = [H+][OH-] log Kw = log [H+] + log [OH-] -log Kw= -log [H+] -log [OH-] ; previous equation multiplied by -1 pKw = pH + pOH; pKw = 14 since Kw =1 x 10-14

14 = pH + pOH

pH = 14 - pOH

pOH = 14 - pH

14 = pH + pOH

pH = 14 - pOH

pOH = 14 - pH


pH and pOH calculations of acid and base solutionsa) Strong acids/bases

dissociation is complete for strong acid such as HNO3 or base NaOH

[H+] is calculated from molarity (M) of the solution

b) weak acids/bases

needs Ka , Kb or percent(%)dissociation


pH of Strong Acid/bases

HNO3(aq) + H2O(l) H3+O(aq) + NO3-(aq)

Therefore, the moles of H+ ions in the solution is equal to moles of HNO3 at the beginning.

[HNO3] = [H+] = 0.2 mole/L

pH = -log [H+]

= -log(0.2)

pH = 0.699

Substance pH

1 M HCl 0.0



Classic Coke 2.5

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45



1M NaOH 14.0

Substance pH

1 M HCl 0.0



Classic Coke 2.5

Coffee 5.0

Pure Water 7.0

Blood 7.35 - 7.45



1M NaOH 14.0


pH of 0.5 M H2SO4 Solution


-(aq)

HSO4-(aq) + H2O(l) H3

+O(aq) + SO42-(aq)

[H3+O][HSO4

-]

H2SO4 ; Ka1 = -------------------

[H2SO4]

[H3+O][SO4

2-]

H2SO4 ; Ka2 = ------------------- ; Ka2 ignored

[HSO4-]


H2SO4(aq) + H2O(l) H3+

O(aq) + HSO4-(aq)

the moles of H+

ions in the solution is equal to moles of H2SO4 at the

beginning.

[H2SO4] = [H+

] = 0.5 mole/L

pH = -log [H+

]

pH = -log(0.5)

pH = 0.30

pH of 0.5 M H2SO4 Solution


1.5 x 10-2 M NaOH.1.5 x 10-2 M NaOH.

NaOH is also a strong base dissociates completely in water.

[NaOH] = [HO- ] = 1.5 x 10-2 mole/L

pOH = -log[HO-]= -log(1.5 x 10-2)

pOH = 1.82

As defined and derived previously:

pKw= pH + pOH; pKw= 14

pH = pKw + pOH

pH = 14 - pOH

pH = 14 - 1.82 ; pH = 12.18


Mixtures of Strong and Weak Acids

the presence of the strong acid retards the dissociation of the weak acid


Measuring pH

Arnold Beckman

inventor of the pH meter

father of electronic instrumentation


Equilibrium, Constant, Ka & Kb

Ka: Acid dissociation constant for a equilibrium reaction.

Kb: Base dissociation constant for a equilibrium reaction.

Acid: HA + H2O H3+O + A-

Base: BOH + H2O B+ + OH-

[H3+O][ A-] [B+ ][OH-]

Ka = --------------- ; Kb = -----------------

[HA] [BOH]



[H3+O][Cl-]

Ka= ----------------- [HCl]

[H+][Cl-] Ka= ----------------- [HCl]

Acid Dissociation Constant


Base Dissociation Constant

NH3 + H2O NH4+ + OH-

[NH4+

][OH-]

K =

[NH3]


Hydrated Metal Ions as Acids

[Fe(H2O)6]3+ (aq) + H2O () [Fe(H2O)5(OH)]2+ (aq) + H3O+ (aq)

Ka [Fe(H2O)5 (OH)2 ][H3O ]

Fe(H2O)63 6.310 3


Ionization Constants for Acids


Comparing Kw and Ka & Kb

Any compound with a Ka value greater than Kw of water will be a an acid in water.

Any compound with a Kb value greater than Kw of water will be a base in water.


WEAKER/STRONGER Acids and Bases & Ka and Kb values

A larger value of Ka or Kb indicates an equilibrium favoring product side.

Acidity and basicity increase with increasing Ka or Kb.

pKa = - log Ka and pKb = - log Kb

Acidity and basicity decrease with increasing pKa or pKb.


Which is weaker?a. HNO2 ; Ka= 4.0 x 10-4.

b. HOCl2 ; Ka= 1.2 x 10-2.

c. HOCl ; Ka= 3.5 x 10-8.

d. HCN ; Ka= 4.9 x 10-10.


What is Ka1 and Ka2?H2SO4(aq) + H2O(l) H3

+O(aq) + HSO4-(aq)


+O(aq) + SO42-(aq)



-(aq)


+O(aq) + SO42-(aq)

[H3+O][HSO4

-]

H2SO4 ; Ka1 = -------------------

[H2SO4]

[H3+O][SO4

2-]

H2SO4 ; Ka2 = -------------------

[HSO4-]

Ka Examples


HC2H3O2(aq) + H2O(l) H3+O(aq) + C2H3O2

-(aq)

[H+][C2H3O2

-]

H C2H3O2; Ka= ------------------

[H C2H3O2]

NH3 (aq) + H2O(l) NH4+ + OH-(aq)

[NH4+][OH-]

NH3; Kb= --------------

[ NH3]

Ka Examples


How do you calculate pH of weak acids/basesFrom % dissociation

From Ka or Kb

What is % dissociation

Amount dissociated

% Dissoc. = ------------------------- x 100

Initial amount


How do you calculate % dissociation from Ka or Kb

1.00 M solution of HCN; Ka = 4.9 x 10-10

What is the % dissociation for the acid?



First write the dissociation equilibrium equation:

HCN(aq) + H 2O(l) <===> H 3+O(aq) + CN-(aq)

[HCN] [H+ ] [CN- ]

Ini. Con. 1.00 M 0.0 M 0.00 M

Cha. Con -x x x

Eq. Con. 1.0 - x x x

[H 3+O ][CN-] x2

Ka = ----------- = ----------------

[HCN] 1.0 - x



1.0 - x ~ 1.00 since x is small

x2

Ka = -----------; Ka = 4.9 x 10-10 = x2

1.0

x = 4.9 x 10-10 = 2.21 x 10 -5

Amount disso. 2.21 x 10 -5

----------------- x 100 =- ------------- x 100 Ini. amount 1.00

% Diss. =2.21 x 10 -5 x 100 = 0.00221 %



% Dissociation gives x (amount dissociated) need for pH calculation

Amount dissociated

% Dissoc. = ------------------------- x 100

Initial amount/con.

x

% Dissoc. = --------------------------- x 100

concentration


1 M HF, 2.7% dissociated

Notice the conversion of % dissociation to a fraction (x): 2.7/100=0.027) x=0.027

Calculate the pH of a weak acid from % dissociation


HF(aq) + H 2O(l) <===> H 3+O(aq) + F-(aq)

[H+][F-]

Ka = ----------- [HF] [HF] [H+ ] [F- ]Ini. Con. 1.00 M 0.0 M 0.00 MChg. Con -x x xEq.Con. 1.0-0.027 0.027 0.027 pH = -log [H+] pH = -log(0.027) pH = 1.57

Calculate the pH of a weak acid from % dissociation


Weak acid Equilibria

Example

Determine the pH of a 0.10 M benzoic acid solution at 25 oC if Ka = 6.28 x 10-5

HBz(aq) + H2O(l) H3O+(aq) + Bz-(aq)

The first step is to write the equilibrium expression

Ka = [H3O+

][Bz-]

[HBz]


Weak acid Equilibria

HBz H3O+ Bz-

Initial conc., M 0.10 0.00 0.00

Change, DM -x x x

Eq. Conc., M0.10 - x x x

[H3O+] = [Bz-] = x

We’ll assume that [Bz-] is negligible compared to [HBz]. The contribution of H3O+ from water is also negligible.


Weak Acid Equilibria

Solve the equilibrium equation in terms of x

Ka = 6.28 x 10-5 =

x = (6.28 x 10-5 )(0.10)

H3O+ = 0.0025 M

pH = 2.60

x2

0.10


pH from Ka or Kb


First write the dissociation equilibrium equation:

HCN(aq) + H 2O(l) H 3+O(aq) + CN-(aq)

[HCN] [H+ ] [CN- ]

Ini. Con. 1.00 M 0.0 M 0.00 M

Chg. Con -x x x

Eq. Con. 1.0 - x x x


[H 3

+O ][CN-] x2

Ka = --------------- = ----------------

[HCN] 1.0 - x

1.0 - x ~ 1.00 since x is small x2 Ka = -----------; Ka = 4.9 x 10-10 = x2

1.0

x = 4.9 x 10-10 = 2.21 x 10 -5

pH = -log [H+]

pH = -log(2.21 x 10-5)

pH = 4.65

Weak Acid Equilibria


The Conjugate Partners of Strong Acids and BasesThe conjugate acid/base of a strong base/acid has

no net effect on the pH of a solution

The conjugate base of a weak acid hydrolyze in water and basic or

pH of a solution > 7.00 E.g. Na+C2H3O2- sodium

acetate

The conjugate acid of a weak base hydrolyze in water and acidic or

pH of a solution < 7.00 E.g NH4Cl


Reaction of a basic anion or acidic cation with water is an ordinary Brønsted-

Lowry acid-base reaction.

CH3COO-(aq) + H2O(l) CH3COOH(aq) + OH

-(aq)

NH4+

(aq) + H2O(l) NH3 (aq) + H3O+

(aq)

This type of reaction is given a special name.

Hydrolysis

The reaction of an anion with water to produce the conjugate acid and OH-.

The reaction of a cation with water to produce the conjugate base and H3O+

.

Hydrolysis


Acid-Base Properties of Typical Ions


What salt solutions would be acidic, basic and neutral?

1) strong acid + strong base = neutral

2) weak acid + strong base = basic

3) strong acid + weak base = acidic

4) weak acid + weak base = neutral,

basic or an acidic solution depending on the relative strengths of the acid and the base.


What pH? Neutral, basic or acidic?

•a)NaCl • neutral•b) NaC2H3O2

• basic•c) NaHSO4 • acidic•d) NH4Cl• acidic


How do you calculate pH of a salt solution?

Find out the pH, acidic or basic?

If acidic it should be a salt of weak base

If basic it should be a salt of weak acid

if acidic calculate Ka from Ka= Kw/Kb

if basic calculate Kb from Kb= Kw/Ka

Do a calculation similar to pH of a weak acid or base


What is the pH of 0.5 M NH4Cl salt solution? (NH 3; Kb = 1.8 x 10-5)

Find out the pH, acidic

if acidic calculate Ka from Ka= Kw/Kb

Ka= Kw/Kb = 1 x 10-14 /1.8 x 10-5)

Ka= 5.56. X 10-10

Do a calculation similar to pH of a weak acid


Continued

NH4+ + H2O H 3

+O + NH3

[NH4+] [H3

+O ] [NH3 ]Ini. Con. 0.5 M 0.0 M 0.00 MChange -x x xEq. Con. 0.5 - x x x [H 3+O ] [NH3 ]

Ka(NH4+) = -------------------- =

[NH 4+] x2

---------------- ; appro.:0.5 - x . 0.5 (0.5 - x)


x2 Ka(NH4

+) = ----------- = 5.56 x 10 -10

0. 5 x2

= 5.56 x 10 -10 x 0.5 = 2.78 x 10 -10

x= 2.78 x 10 -10 = 1.66 x 10-5

[H+ ] = x = 1.66 x 10-5 MpH = -log [H+ ] = - log 1.66 x 10-5

pH = 4.77

pH of 0.5 M NH4Cl solution is 4.77 (acidic)

Continued


Types of Acids and Bases

Binary acids

Oxyacid

Organic acids

Acidic oxides

Basic oxides

Amine

Polyprotic acids


Influence of Molecular Structure on Acid Strength

Binary Hydrides• hydrogen & one other element

Bond Strengths• weaker the bond, the stronger the acid

Stability of Anion• higher the electronegativity, stronger the acid


Binary Acids

Compounds containing acidic protons bonded to a more electronegative atom.

e.g. HF, HCl, HBr, HI, H2S

The acidity of the haloacid (HX; X = Cl, Br, I, F)Series increase in the following order: HF < HCl < HBr < HI


Oxyacids

Compounds containing acidic - OH groups in the molecule.

Acidity of H2SO4 is greater than H2SO3 because of the extra O (oxygens)

The order of acidity of oxyacids from the a halogen (Cl, Br, or I) shows a similar trend.

HClO4 > HClO3 > HClO2 > HClO

perchloric chloric chlorus hyphochlorus


Influence of Molecular Structure on Acid Strength

Oxyacids• hydrogen, oxygen, & one other element

H-O-E• higher the electronegativity on E, stronger the acid as

this weakens the bond between the O and H


< <<

<

Oxo Acid


Acidic Oxides

These are usually oxides of non-metallic elements such as P, S and N.

E.g. NO2, SO2, SO3, CO2

They produce oxyacids when dissolved in waterSO3 + H2O ---> H2SO4

CO2 + H2O ---> H2CO3

NO2 + H2O ---> HNO3


Basic Oxides

Oxides oxides of metallic elements such as Na, K, Ca. They produce hydroxyl bases when dissolved in water.

e.g. CaO + H2O ---> Ca(OH)2

Na2O + H2O ---> 2 NaOH


Protic Acids

Monoprotic Acids: The form protic refers to acidity due to protons. Monoprotic acids have only one acidic proton. e.g. HCl.

Polyprotic Acids: They have more than one acidic proton.

e.g. H2SO4 - diprotic acid

H3PO4 - triprotic acid.


Polyprotic Acidsacids where more than one hydrogen per molecule

is released


Polyprotic Acids


Organic or Carboxylic Acids

H C

H

H

C

H

H

C

H

H

C

O

O H

nonacidic hydrogens

butanoic acid

acidic hydrogen

CH 3 C

O

acetic acid

OH

electron-attractingoxygen atom

acidic hydrogen

CH 3 C

O

OH

-

CH 3 C

O

O-

acetate ion


FCH2CO2H (strongest acid) > ClCH2CO2H > BrCH2CO2H (weakest acid).

Acid Ka pKa

HCOOH (formic acid) 1.78 X 10-43 0.75

CH3COOH (acetic acid) 1.74 X 10-54 0.76

CH3CH2COOH (propanoic acid)1.38 x 10-5 4.86

Organic or Carboxylic Acids


Amines

Class of organic bases derived from ammonia NH3 by replacing hydrogen by organic groups. They are defined as bases similar to NH3 by Bronsted-Lowery or Lewis acid/base definitions.


Amines


Acid-Base Chemistryof Some Antacids


Acid-Base in the Kitchenvinegar - acetic acid

lemon juice (citrus juice) - citric acid

baking soda - NaHCO3

milk - lactic acid

baking powder - H2PO4- & HCO3

-


Household Cleaners

CH 3CH2CH 2CH2CH 2CH2CH 2CH 2CH 2CH 2CH2CH 2CH2CH 2SO3

-Na+

Oil-soluble part(hydrophobic)

Water-soluble part(hydrophilic)

A Typical Synthetic Detergent Molecule

CH 3(CH 2)4COO(CH 2)2O( CH2CH 2O) 2CH 2CH 2OH

esterlink

(hydro-philic)

etherlink

etherlink

(hydrophilic)

hydrocarbonchain

(hydro-phobic)

alcohol group(hydrophilic)

A nonionic detergent


Dishwashing Detergent


Lewis DefinitionG.N. Lewis was successful in including acid and bases

without proton or hydroxyl ions.

Lewis Acid: A substance that accepts an electron pair.

Lewis base: A substance that donates an electron pair.

E.g. BF3(g) + :NH3(g) F3B:NH3(s)

the base donates a pair of electrons to the acid forming a

coordinate covalent bond common to coordination

compounds. Lewis acids/bases will be discussed later in

detail


Lewis Acids and Bases ReactionsH+ + NH3 NH4

+

acid base

Cu+2 + 4 NH3 [Cu(NH3)4+2]

acid base


What acid base concepts (Arrhenius/Bronsted/Lewis) would best describe the following reactions:

a) HCl(aq) + NaOH(aq) ---> NaCl(aq) + H2O(l)

b)HCl(g) + NH3(g) ---> NH4Cl(s)

c)BF3(g) + NH3(g) ---> F3B:NH3(s)

d)Zn(OH)2(s) + 2OH-(aq) ---> [Zn(OH)4]2- (aq)

chemistry 102(01) fall 2010

Documents

aq hclaq

reactions of acids

hcl aq h aq

acids bases

h2ol h2ol h3 oaq ohaq

aq baseanything

acidbase reactions of

lewis base