Chemistry 100

Aqueous Reactions

Solutions

A solution is a homogenous mixture of two or more substances

One substance (generally the one present in the greatest amount) is called the solvent

The other substances - those that are dissolved - are called the solutes

The Solution Process

Favourable interactions between the solute and the solvent drive the formation of a solution

Example: NaCl (an ionic solid) dissolving in water

Water is a polar fluid (i.e., possesses a permanent dipole)

Electrolytes

Salt is an ionic compound. NaCl is dissolved in water - the ions

separate. The resulting solution conducts electricity .

A solute with this property is called an electrolyte

Strong Electrolytes

Strong electrolytes - completely dissociated

Some molecular compounds dissolve in water to form ions. Dissolve HCl (g) in water. All the molecules dissociate. So it is also a

strong electrolyte.

Weak and Nonelectrolytes

Weak electrolytes - only some of the molecules dissociate, i.e., acetic acid

Compounds that do not dissociate - nonelectrolytesSugarsUreasAlcohols

Acids

Acid - a substance that ionizes in water to form hydrogen ions H+.

HCl (aq) H+ (aq) + Cl(aq) What is H+? A hydrogen atom without its

electron - a bare proton.

Monoprotic, diprotic, triprotic One molecule of HCl gives one H+ ion:

HCl H+ + Cl

We say that HCl is monoprotic - one proton

One molecule of sulphuric acid, H2SO4, has two hydrogens to give away. It is said to be diprotic.

Phosphoric acid, H3PO4 is triprotic.

OO

O

S

O H

H

C

H

H

H C

O

OH only this H ionizes

both H's ionize

Some Chemical Structures

Acetic Acid

Generally write as CH3COOH, not HC2H3O2. Weak acid - doesn’t dissociate completely

CH3COOH (aq) ⇄ CH3COO- (aq) + H+ (aq)

The double arrow - the system is in chemical equilibrium!!!!

Bases Bases are substances that accept (react with)

H+ ions. Hydroxide ions, OH, are basic. They react with H+ ions to form water:

H+ (aq) + OH (aq) H2O (l)

Ionic hydroxides like NaOH, KOH, Ca(OH)2 are basic. When dissolved in water they form hydroxide ions.

Ammonia solution When ammonia gas dissolves in water,

some NH3 molecules react with water:

NH3(aq) + H2O(l) ⇄ NH4+

(aq) + OH– (aq)

NOTE - only some NH3 molecules

react with water. Ammonia is a

weak electrolyte.

Strong and Weak Acids and Bases Acids and bases that are strong

electrolytes are called strong acids and strong bases.

Strong acids are more reactive than weak acids. Likewise for bases.

Note exception - HF, a weak acid, is very reactive

Acids you should knowChloric acid HClO3

Hydrobromic acid HBr

Hydrochloric acid HCl

Hydroiodic acid HI

Nitric acid HNO3

Perchloric acid HClO4

Sulphuric acid H2SO4

Acetic acid CH3COOH (weak)

Bases you should know Know the following bases:

Strong bases

a) Hydroxides of alkaline metals: LiOH, NaOH, KOH

b) Hydroxides of the heavy alkaline earth metals: Ca(OH)2, Sr(OH)2, Ba(OH)2

Weak base: ammonia solution NH3

Metathesis reactions A metathesis reaction is an aqueous

solution in which cations and anions appear to exchange partners.

AX + BY AY + BX

AgNO3 (aq)+ NaCl (aq) AgCl (s) + NaNO3

(aq)

Metathesis reactions (cont.)

Three driving forces Precipitate formation (insoluble

compound)

AgNO3(aq)+ NaCl(aq) AgCl(s) + NaNO3(aq)

Metathesis Reactions (Cont’d)

Weak electrolyte or nonelectrolyte formation

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)

Gas formation 2HCl(aq) + Na2S(aq) NaCl(aq) + H2S(g)

Neutralization Mix solutions of acids and bases - a

neutralization reactions occurs.

acid + base salt + water Salt does not necessarily mean sodium

chloride!!!! Salt - an ionic compound whose cation

(positive ion) comes from a base and whose anion (negative ion) comes from an acid

Precipitation Reactions Some ionic compounds are insoluble in

water. If an insoluble compound is formed by

mixing two electrolyte solutions, a precipitate results.

Precipitation (Cont’d)

Solubility - maximum amount of substance that will dissolve in a specified amount of solvent.

Saturated solution of PbI2 contains 1 x 10-3

mol/L. A compound with a solubility of less than

0.01 mol/L - insoluble. More accurately - sparingly soluble.

Solubility Fact 1

All the common ionic compounds of the alkali metals are soluble in water. The same is true of the compounds containing the ammonium ion, NH4

+.

NaCl, K2CO3, (NH4)2S are all soluble

Solubility Fact 2

Salts containing the following anions are soluble

Anion exception, salts ofNO3

nitrate noneCH3COO acetate noneCl chloride Ag+, Hg2

2+,Pb2+

Br bromide Ag+, Hg22+,Pb2+

I iodide Ag+, Hg22+,Pb2+

SO42 sulphate Ca2+, Sr2+, Ba2+,

Hg22+, Pb2+

Solubility Fact 3

Salts containing the following anions are insolubleAnion exception, salts of

S2 sulphide alkaline metal cations, NH4

+, Ca2+, Sr2+, Ba2+, CO3

2 carbonate alkaline metal cations, NH4

+ PO4

3 phosphate alkaline metal cations, NH4

+

OH hydroxide alkaline metal cations, Ca2+, Sr2+, Ba2+,

Reaction forming gases

A metathesis reaction can occur due to the formation of a gas which is not very soluble in water.

Examples involving hydrogen sulphide and carbon dioxide

Reactions forming H2S

A metathesis reaction occurs when hydrochloric acid is added to a sodium sulphide solution.

2HCl(aq) + Na2S(aq) H2S(g) + 2NaCl(aq) Net ionic reaction:

2H+(aq) + S2(aq) H2S (g)

Reactions involving CO2

Carbonates and bicarbonates may be thought of as the salts of carbonic acid H2CO3 – unstable!!

H2CO3(aq) CO2(g) + H2O(l)

Ionic Equations

Consider the reaction

HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l) The above is known as the molecular

equation Note: the compounds are ionic (except

water)!!

Ionic Equations #2

Let’s show ionic compounds as ions

H+(aq) + Cl–(aq) + Na+(aq) + OH– (aq)

Na+(aq) + Cl–(aq) + H2O(l) Some ions appear on both sides of the

equation.

Out with the spectators!

Remove ions that appear on both sides

H+ (aq) + Cl– (aq) + Na+ (aq) + OH– (aq)

Na+ (aq) + Cl– (aq) + H2O (l) The unchanged ions are called spectators

The Net Ionic Equation

We are left with is the net ionic equation:

H+(aq) + OH– (aq) H2O(l)

Note that the equation is balanced for both mass and charge!!!

Another ionic reaction

Place zinc metal in a hydrochloric acid solution – hydrogen is evolved!!

Zn (s) + 2HCl (aq) ZnCl2 (aq) + H2 (g)

Why use ionic reactions?

They summarize many reactions. neutralization of any strong acid by a strong base

is given by H+(aq) + OH– (aq) H2O(l)

The chemical behaviour of a strong electrolyte behaviour of its constituent ions.

Ionic equations can be written only for strong electrolytes which are soluble.

Concentrations

How do we express the concentration of a solution?

Percentage is one way. 2% milk 35% cream. (These are not true solutions)\Some beer is 5% alcohol

Note: % measurements can be %w/w, %w/v, %v/v

Molarity

Must work in moles to do chemical arithmetic.

Chemists - molarity as their unit of solution concentration

(L)solution of volume

solute of molesMolarity

Dilution

Dilute a solutionmore solvent is added but the amount (mass

or moles) of solute is unchanged.

M1V1 = M2V2

The volumes can be either millilitres (mL) or litres (L).

Ionic Concentration

NaCl in water - totally ionized into Na+ and Cl ions.

A 2.0 M NaCl solution Na+ concentration will be 2.0 M Cl concentration also 2.0 M

A 2.0 M solution of K2CO3, K+ concentration will be 4.0 M The concentration of CO3

2 2.0 M.

Oxidation and reduction

A piece of calcium metal exposed to the air will react with the oxygen in the air

2Ca(s) + O2(g) 2 CaO(s) Ca has been converted to an ion Ca2+ by

losing two 2 electrons. Dissolve Ca in acid

Ca(s) + 2H+(aq) Ca2+(aq) + H2(g) Again the Ca has lost 2 electrons — oxidation

Redox reactions

In the last two reactions, the Ca atom lost two electrons. Where did they go?

When one substance is oxidized, another is reduced. An oxidation-reduction reaction occurs. Or a redox reaction occurs.

Oxidation: loss of electrons (more positive) Reduction: gain of electrons (less positive)

Oxidation of Metals - by air

Many metals react with oxygen in the air. Na and K do so explosively!

Fe rusts - at a cost of $billions each year! Aluminum oxidizes

oxide layer forms a skin which prevent further oxidation. Al hides its reactivity.

Gold and platinum do not react with oxygen. Silver tarnishes mainly because of H2S in the

air. What does copper do?

Oxidation of Metals - by acids

Many metals react with acids:

metal + acid salt + hydrogen gas

Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g) Metals may also be oxidized by the salts of

other metals. Recall your lab experiment

Fe(s) + CuSO4(aq) Cu(s) + FeSO4(aq) In the above, which metal is oxidized, which is

reduced?

Activity Series

We has seen that some metals react with air, some also react with acids to give hydrogen. We has also seen that some metals can be oxidized by ions of other metals.

All this is summarized in the activity series.

Activity Series

Li Li+ + eK K+ + eBa Ba2+ + 2eCa Ca2+ + 2eNa Na+ + eMg Mg2+ + 2eZn Zn2+ + 2eFe Fe2+ + 2ePb Pb2+ + 2eH H+ + eCu Cu2+ + 2eAg Ag+ + eAu Au3+ + 3e

A metal can be oxidized by any ion below it

Metals above H, react with acids to give H2

The further up the series, the more readily the metal is oxidized

See your textbook (p124) for more elements

Some observations on the series Lead (Pb) is above H, so is Al. But these

metals are not attacked by 6M HCl. They form very protective oxides.

Cu reacts with nitric acid (HNO3) because that acid is a strong oxidizing agent in addition to being an acid.

Gold (Au) and platinum (Pt) are valuable because they are (a) rare and (b) unreactive - they do not tarnish

Oxidation Numbers

Oxidation number - a fictitious charge assigned to atoms either by themselves or when combined in compounds as an electron bookkeeping device.

There are a number of simple rules that chemists use to assign oxidation numbers.

Assigning Oxidation Numbers

In any elemental form (atom or molecule), an atom is assigned a 0 oxidation number e.g. He, Cu, N in N2, S in S8

For a monatomic ion, the oxidation number equals the chargee.g., -1 for Cl in Cl-, +2 for Ca+2, -2 for S-2

Fluorine’s oxidation number is -1 in any compound.e.g. -1 for F in CF4, but 0 for F in F2

Oxygen’s oxidation number is -2 except when combined with fluorine or in peroxides.e.g. -2 for O in H2O and OH-, +2 for O in OF2, -1

for O in H2O2

Assigning Ox. Numbers (#2)

For elements in Groups IA, IIA & most of IIIA, oxidation numbers are positive and equal to the group number.e.g. +3 for Al in AlCl3, +1 for Na in NaCl, +2 for

Mg in Mg SO4

Hydrogen has a +1 oxidation number. Exceptions to this rule are the metallic hydrides, in which it is -1.e.g., +1 for H in H2O and CH3OH, -1 for H in NaH

Assigning Ox. Numbers (#3)

The sum of the oxidation numbers of the atoms in a neutral compound is zero; in a polyatomic ion, the sum equals the charge.e.g. see OH- and H2O above, +6 for S in SO4

-2

Assigning Ox. Numbers (#4)

Balancing Oxidation-Reduction (Redox) Equations (#1) Assign oxidation numbers to all atoms in

the equation. Note - polyatomic ion that is unchanged in the

reaction may be treated as a single unit with an oxidation number equal to its charge.

Balancing Redox Equations (#2)

Isolate the ATOMS that have undergone a change of oxidation numberA reduction in number indicates a reduction

An increase in number, an oxidation

Balancing Redox Equations (#3)

Isolate the chemical species undergoing oxidation/reduction (note: separate into an oxidation and a reduction half-reaction).

Add the appropriate number of electrons to the half-reactionsOxidation – electrons on products side

Reduction – electrons on reactants side

Balancing Redox Equations (#4)

Remaining steps refer to the individual half reactionsBalance for charges

Add H+ in acidic solution

Add OH- in basic solution

Balance the H and the O atoms by adding water

Balance the number of electrons in the half-reactions Note: electrons lost = electrons gained

Add the half-reactions, eliminating the electrons and obtaining the complete REDOX equation

Balancing Redox Equations (#5)