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 nonelectrolytes    Sugars Ureas Alcohols

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 + HCl  H + + Cl  ion: We say that HCl is

monoprotic -

one proton One molecule of sulphuric acid, H 2 SO 4 , has two hydrogens to give away. It is said to be

diprotic

.

Phosphoric acid, H 3 PO 4 is

triprotic.

Some Chemical Structures

O S O O H O H both H's ionize H H C H C O OH only this H ionizes

Acetic Acid

 Generally write as CH 3 COOH, not HC 2 H 3 O 2 .  Weak acid doesn’t dissociate completely CH 3 COOH (aq) ⇄ CH 3 COO (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)  H 2 O (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 NH 3 water: NH 3

(aq)

molecules react with + H 2 O

(l)

⇄ NH 4 +

(aq) +

OH –

(aq)

NOTE - only some NH 3 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 know

Chloric acid Hydrobromic acid Hydrochloric acid Hydroiodic acid Nitric acid Perchloric acid Sulphuric acid Acetic acid HClO 3 HBr HCl HI HNO 3 HClO 4 H 2 SO 4 CH 3 COOH (weak)

Bases you should know

 Know the following bases: Strong bases a) Hydroxides of alkali 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 NH 3

Metathesis reactions

 A metathesis reaction is an aqueous solution in which cations and anions appear to exchange partners.

AX + BY  AY + BX AgNO 3 (aq)+ NaCl (aq)  NaNO 3 (aq) AgCl (s) +

Metathesis reactions (cont.)

  Three driving forces Precipitate formation (insoluble compound) Ag NO 3

(aq)+

Na Cl

(aq)

 NaNO 3

(aq)

AgCl

(s) +

Metathesis Reactions (Cont’d)

  Weak electrolyte or nonelectrolyte formation H Cl

(aq)

+ Na OH

(aq)

 H 2 O

(l)

NaCl

(aq)

Gas formation 2 H Cl

(aq)

+ Na 2 S

(aq)

 H 2 S

(g)

2 NaCl

(aq)

+ +

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 PbI 2 10 -3 mol/L.

contains 1 x 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, NH 4 + .

NaCl, K 2 CO 3 , (NH 4 ) 2 S are all soluble

Solubility Fact 2

Salts containing the following anions are soluble

Anion exception, salts of

NO 3  CH 3 COO  Cl  Br  I  SO 4 2  nitrate acetate chloride bromide iodide sulphate none none Ag + , Hg 2 2+ ,Pb 2+ Ag + , Hg 2 2+ ,Pb 2+ Ag + , Hg 2 2+ ,Pb 2+ Ca 2+ , Sr 2+ , Ba 2+ , Hg 2 2+ , Pb 2+

Solubility Fact 3

 Salts containing the following anions are insoluble S 2 

Anion

sulphide CO 3 2  PO 4 3  OH  carbonate phosphate hydroxide

exception, salts of

alkaline metal cations, NH 4 + , Ca 2+ , Sr 2+ , Ba 2+ , alkaline metal cations, NH 4 + alkaline metal cations, NH 4 + alkaline metal cations, Ca 2+ , Sr 2+ , Ba 2+ ,

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 H

2

S

  A metathesis reaction occurs when hydrochloric acid is added to a sodium sulphide solution. 2HCl(aq) + Na 2 S(aq)  2NaCl(aq) H 2 S(g) + Net ionic reaction: 2H + (aq) + S 2  (aq)  H 2 S (g)

Reactions involving CO

2  Carbonates and bicarbonates may be thought of as the salts of carbonic acid H 2 CO 3 – unstable!!

H 2 CO 3

(aq)

 CO 2

(g)

+ H 2 O

(l)

Ionic Equations

  Consider the reaction HCl (aq) + NaOH (aq)  H 2 O (l) NaCl (aq) + 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)

+ H 2 O

(l)

 Some ions appear on both sides of the equation.

Out with the spectators!

 H + (aq) + Cl –  (aq) + Na + (aq) + OH – (aq) Na + (aq) + Cl – (aq) + H 2 O (l)  Remove ions that appear on both sides The unchanged ions are called spectators

The Net Ionic Equation

 We are left with is the

net ionic equation:

H +

(aq) +

OH –

(aq)

 H 2 O

(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)  ZnCl 2 (aq) + H 2 (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)

 H 2 O

(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 Molarity  moles volume of of solute solution (L)

Dilution

 Dilute a solution  more solvent is added but the amount (mass or moles) of solute is unchanged.

M 1 V 1 = M 2 V 2 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 + Cl  concentration will be 2.0 M concentration also 2.0 M A 2.0 M solution of K 2 CO 3 ,   K + concentration will be 4.0 M The concentration of CO 3 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)

+ O 2

(g)

 2 CaO

(s)

Ca has been converted to an ion Ca 2+ by losing two 2 electrons. Dissolve Ca in acid Ca

(s)

+ 2H +

(aq)

 Ca 2+

(aq)

+ H 2

(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 H 2 S in the air.

What does copper do?

Oxidation of Metals - by acids

  Many metals react with acids: metal + acid Mg

(s)

 salt + hydrogen gas + 2HCl

(aq)

 MgCl 2

(aq)

+ H 2

(g)

Metals may also be oxidized by the salts of other metals. Recall your lab experiment Fe

(s)

+ CuSO 4

(aq)

 Cu

(s)

+ FeSO 4

(aq)

Activity Series

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

All this is summarized in the

activity series.

Activity Series

Li K Ba    Ca  Na  Mg Zn Fe Pb     H Cu Ag Au     Li + K + Ba 2+ Ca 2+ Na + Mg 2+ Zn 2+ Fe 2+ Pb 2+ H + Cu 2+ Ag + Au 3+ + e + e + 2e + 2e + e + 2e + 2e + 2e + 2e + e + 2e + e + 3e     A metal can be oxidized by any

ion below

it Metals above H, react with acids to give H 2 The further up the series, the more readily the metal is oxidized See your textbook (p 136) 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 (HNO 3 ) 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 N 2 , S in S 8 For a monatomic ion, the oxidation number equals the charge  e.g., -1 for Cl in Cl , +2 for Ca +2 , -2 for S -2

Assigning Ox. Numbers (#2)

  Fluorine’s oxidation number is -1 in any compound.

 e.g. -1 for F in CF 4 , but 0 for F in F 2 Oxygen’s oxidation number is -2 except when combined with fluorine or in peroxides.

 e.g. -2 for O in H 2 O and OH , +2 for O in OF 2 , -1 for O in H 2 O 2

Assigning Ox. Numbers (#3)

  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 AlCl 3 , +1 for Na in NaCl, +2 for Mg in Mg SO 4 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 H 2 O and CH 3 OH, -1 for H in NaH

Assigning Ox. Numbers (#4)

 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 SO 4 -2 and H 2 O above, +6 for S in

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 number  A 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-reactions   Oxidation – electrons on products side Reduction – electrons on reactants side

Balancing Redox Equations (#4)

 Remaining steps refer to the individual half reactions  Balance for charges  Add H + in acidic solution  Add OH in basic solution  Balance the H and the O atoms by adding water

Balancing Redox Equations (#5)

 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

Titrations

  Volumetric analysis  technique based on volume measurements  used to determine the quantity of a substance in solution.

Titration  a solution of an accurately known concentration is added gradually to a solution of an unknown concentration  Reaction goes to completion.

Other Definitions

  Standard solution  solution of accurately known concentration.

Equivalence point  point at which unknown substance has completely reacted with standard solution.

 At the equivalence point reagents are present in stoichiometric amounts.

Gravimetric Analysis

 Determine concentration of an unknown by reacting it with a second substance to form a ppt.

AgNO 3

(aq)+

NaCl

(aq)

 NaNO 3

(aq)

AgCl

(s) +