Chem6A_F11 Hol - Chapter 4A (Part 1)

Handout #4A (Fall 2011) Chem 6A - Prof. John E. Crowell

Chem 6A - Chapter 4 (Part 1) 1

Three Major Classes of Chemical Reactions

4.6 Elements in Redox Reactions

4.1 The Role of Water as a Solvent

4.2 Writing Equations for Aqueous Ionic Reactions

4.3 Precipitation Reactions

4.4 Acid-Base Reactions

4.5 Oxidation-Reduction (Redox) Reactions

Electron distribution in molecules of H2 and H2O As a result, water is a polar solvent and is able to solvate polar molecules and ionic species!



The dissolution of an ionic compound

Determining Moles of Ions in Aqueous Ionic Solutions How many moles of each ion are in the following solutions?

(a) 5.0 mol of ammonium sulfate dissolved in water (b) 78.5 g of cesium bromide dissolved in water (c) 7.42 x 1022 formula units of copper(II) nitrate dissolved in water (d) 35 mL of 0.84 M zinc chloride

(a) (NH4)2SO4(s) 2 NH4+(aq) + SO4

2-(aq)

5.0 mol (NH4)2SO4 x 2 mol NH4

+

1 mol (NH4)2SO4

= 10. mol NH4+

and 5.0 mol SO42-

7.42 x 1022 formula units x Cu(NO3)2

mol Cu(NO3)2

6.022 x 1023 formula units = 0.123 mol Cu(NO3)2

and 0.123 mol Cu2+

= 0.246 mol NO3-

(c) Cu(NO3)2(s) Cu2+(aq) + 2 NO3-(aq)



Writing Equations for Aqueous Ionic Reactions The molecular equation

shows all of the reactants & products as intact, undissociated compounds:

The total ionic equation shows all of the soluble ionic substances dissociated into ions:

The net ionic equation omits the spectator ions & shows the actual chemical change taking place:

The reaction of Pb(NO3)2 and NaI

Double-displacement reaction (metathesis)

NaI(aq) + Pb(NO3)2(aq)

PbI2(s) + NaNO3(aq)

2 NaI(aq) + Pb(NO3)2(aq)

PbI2(s) + 2 NaNO3(aq)

2 Na+(aq) + 2 I-(aq) + Pb2+(aq) + 2 NO3-(aq)

PbI2(s) + 2 Na+(aq) + 2 NO3-(aq)

2 NaI(aq) + Pb(NO3)2(aq)

PbI2(s) + 2 NaNO3(aq)



An aqueous ionic reaction

and its equations:

Molecular,

Total Ionic,

And

Net Ionic.

Predicting Whether a Precipitate Will Form

1. Note the ions present in the reactants.

2. Consider the possible cation-anion combinations (Double Displacement).

3. Decide if any of the ion combinations are insoluble.

To do this we need to know the solubility rules.



Solubility Rules for Ionic Compounds in Water Soluble Ionic Compounds: 1.  All common compounds of Group 1A ions (Li+, Na+, K+, etc.) and the

ammonium ion (NH4+) are soluble.

2.  All common nitrates (NO3-), acetates (C2H3O2

-), and most perchlorates (ClO4

-) are soluble. 3.  All common chlorides (Cl-), bromides (Br-), and iodides (I-) are

soluble, except those of Ag+, Pb2+, Cu+, and Hg22+. All common

fluorides (F-) are soluble, except those of Pb2+ and Group 2A. 4.  All common sulfates (SO4

2-) are soluble, except those of Ca2+, Sr2+, Ba2+, Ag+, and Pb2+.

Insoluble Ionic Compounds: 1.  All common metal hydroxides (Mn+ + OH-) are insoluble, except

those of Group 1A (Li+, Na+, K+, etc.) and the larger members of Group 2A (i.e. Ca2+ and larger).

2.  All common carbonates (CO32-) and phosphates (PO4

3-) are insoluble, except those of Group 1A and NH4

+. 3.  All common sulfides (S2-) are insoluble, except those of Group 1A,

Group 2A, and NH4+.

Predicting Precipitation Reactions Predict whether a reaction occurs when each of the following pairs of solutions are mixed. If a reaction does occur, write balanced molecular, total ionic, and net ionic equations, and identify the spectator ions.

(a) Potassium fluoride (aq) + strontium nitrate (aq)

(b) Ammonium perchlorate (aq) + sodium bromide (aq)

(a) KF (aq) + Sr(NO3)2 (aq) 2 KNO3 (aq) + SrF2 (s)

2 K+ (aq) + 2 F- (aq) + Sr2+ (aq) + 2 NO3- (aq) 2 K+ (aq) + 2 NO3

- (aq) + SrF2 (s)

(b) NH4ClO4(aq) + NaBr (aq) NH4Br (aq) + NaClO4(aq)

All reactants and products are soluble so no reaction occurs.

2 F- (aq) + Sr2+ (aq) SrF2 (s) Net Ionic Equation

Molecular Equation

Total Ionic Equation



Acids and Bases: Definitions

All definitions are operationally equivalent; however the Lewis definition is by far the most general

Acids and Bases: Definitions PROTON is commonly used to describe a hydrogen ion (“H+”); in acid/base chemistry it is still used to denote an acidic hydrogen and should NOT be confused with the more traditional definition of a proton (a positively charged nuclear particle).

PROTON TRANSFER: in acid/base chemistry this is used to denote the transfer of a hydrogen ion from an acid (H-donor) to an acceptor (H-acceptor).

WEAK ACID or BASE VS. STRONG ACID or BASE: The strength of an acid or a base is a direct function of the degree of ionization of the acid or base (or, how well it does proton transfer reactions). A relative term.



Acids and Bases: Definitions “H+”, as a discrete species, does NOT exist in water; rather, it combines with water to give the hydronium ion:

It is the magnitude of the hydronium ion concentration ([H3O+]; in mol/L) that determines how acidic a solution is!

Since H3O+ is generated from H+, and since H+ comes directly from the acid under consideration, measurement of the [H3O+]

will have a direct correlation to the strength of the acid!

Acid-Base Reactions In the most general sense, an acid reacts with a base by transferring an ionizable hydrogen (an “acidic H”) to it:

The term “conjugate” is used exclusively to describe the materials on the product side of this reaction. Conjugates are easy to write:

If the base is a hydroxide base, a neutralization reaction occurs:

A-, is the conjugate base of the acid HA H3O+ is the conjugate acid of the base H2O

NOTE: Conjugate Pairs differ by a PROTON



Before Dissociation

After Dissociation R

elat

ive

Num

ber o

f Mol

es

Perchloric acid (HClO4) Chloric acid (HClO3) Sulfuric acid (H2SO4) Selenic acid (H2SeO4) Nitric acid (HNO3)

Hydrochloric acid (HCl) Hydrobromic acid (HBr) Hydroiodic acid (HI)

Ka >> 1

100% Dissociation

At Equilibrium

The 8 Strong Acids:

Before Dissociation After Dissociation

Rel

ativ

e N

umbe

r of M

oles

Ka << 1

Few % Dissociation

At Equilibrium

Chlorous acid (HClO2) Hypochlorous acid (HClO) Sulfurous acid (H2SO3) Nitrous acid (HNO2)

Hydrofluoric acid (HF) Acetic acid (CH3COOH) RCOOH, HSO4

-, NH4+

Examples:



Strong Acids:!HBr, HCl, HI, HNO3, H2SO4, "H2SeO4, HClO4, HClO3!

Weak Acids: all other acids:!HX, HOX, HOXOn!

Strong Bases:!NaOH, KOH, Ca(OH)2, Sr(OH)2, "Ba(OH)2, Group I and Group II "oxides (MO) and amides (MNH2)!

Weak Bases: all other bases!NH3, RNH2, R2NH, R3NH!

Start of titration Excess of acid

Point of neutralization

Slight excess of base

An acid-base titration

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Chem6A_F11 Hol - Chapter 4A (Part 1)