Why do reactions occur at all?

Why do reactions occur at all?

• The driving force of all reactions is related to energy• The driving force is not “make everything achieve its

lowest possible energy,” although that often happens

• The driving force is not “increase the entropy of the system,” although that often happens.

• We use a bookkeeping technique called “free energy” to calculate the driving force of all reactions.

Wait until chapter eight

Types of Reactions (I)

• Precipitation Reactions: A process in which an insoluble solid precipitate drops out of the solution.

• Most precipitation reactions occur when the anions and cations of two ionic compounds change partners.

Pb(NO3)2(aq) + 2 KI(aq) 2 KNO3(aq) + PbI2(s)


• Acid–Base Neutralization: A process in which an acid reacts with a base to yield water plus an ionic compound called a salt.

• The driving force of this reaction is the formation of the stable water molecule.

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


• Oxidation–Reduction (Redox) Reaction: A process in which one or more electrons are transferred between reaction partners.

• The driving force of this reaction is the decrease in electrical potential.

Mg(s) + I2(g) MgI2(s)

Types of Reactions (II)

• Metathesis Reactions: These are reactions where two reactants

just exchange parts.

AX + BY AY + BX

HNO3(aq) + KOH(aq) KNO3(aq) + HOH(l)

BaCl2(aq) + K2SO4(aq) BaSO4(s) + 2 KCl(aq)

Double Exchange


• Single Exchange Reactions: These are reactions where one

reactants switches place with a partnered reactant.

A + BY AY + B

3 C(s) + 2 Fe2O3(s) 4 Fe(s) + 3 CO2(g)


• Combination Reactions: These are reactions where two parts

come together.

A + B AB

H2O(g) + SO3(g) H2SO4(aq)


• Decomposition Reactions: These are reactions where a large

molecule falls apart.

AB A + B

NH4NO2(s) N2(g) + 2 H2O(g)

Electrolytes

• Electrolytes: Dissolve in

water to produce ionic

solutions.

• Nonelectrolytes: Do not

form ions when they

dissolve in water.

• Why do ionic

compounds conduct

electricity when

molecular ones

generally do not?

Electrolytes

Electrolytes

• Strong Electrolyte: Total dissociation when dissolved in water.

• Weak Electrolyte: Partial dissociation when dissolved in water.

Electrolytes

Calculating Electrolyte Concentrations

• At this point, to calculate electrolyte concentration we must assume 100% of the ions dissociate

• Be sure to account for the number of ions in one formula unit when calculating ion concentrations

For example:

aqClOaqCasClOCa OH 42

24 22

3 ions from 1 formula unit

Chemical Equations

• Three basic equations are used for reactions:

– Molecular Equation: All reactants and products are written

in molecular form.

– Ionic Equation: All dissolved strong electrolytes are written

as the dissociated ions.

– Net Ionic Equation: All ions that are identical on both sides

are deleted.

Chemical Equations

• Molecular Equation: All reactants and products are written in

molecular (non-dissociated) form along with their phases.

– Pb(NO3)2(aq) + 2 KCl(aq) PbCl2(s) + 2 KNO3(aq)

– 2 HCl(aq) + Cu(OH)2(s) CuCl2(aq) + 2 HOH(l)

– C2H3O2H(aq) + KOH(aq) KC2H3O2(aq) + HOH(l)

Chemical Equations

• Ionic Equation: All dissolved strong electrolytes in the molecular

equation are broken into their ions.

– Pb2+ + 2 NO3– + 2 K+ + 2 Cl– PbCl2(s) + 2 K+ + 2 NO3

–

– 2 H+ + 2 Cl– + Cu(OH)2(s) Cu2+ + 2 Cl– + 2 HOH(l)

– C2H3O2H(aq) + K+ + OH– K+ + C2H3O2– + HOH(l)

Chemical Equations

• Net Ionic Equation: Spectator ions that occur on both sides are cancelled to give only those species which undergo change.

– Pb2+ + 2 Cl– PbCl2(s)

– 2 H+ + Cu(OH)2(s) Cu2+ + 2 HOH(l)

– C2H3O2H(aq) + OH– C2H3O2– + HOH(l)

• Solubility is not really an “on-off switch”, but rather a continuum

• A compound is soluble if it can make a 0.01 M solution.

• Many things can affect solubility: temperature, pressure, presence of other solutes, etc.

Solubility

Precipitation Reactions: Solubility Rules

• Rules are not a complete set, but still pretty good

• Lower numbered rules take precedence over higher-numbered rules

Precipitation Reactions: Solubility Rules

1. All compounds of the alkali metals and ammonium are soluble

2. All nitrates, nitrites, perchlorates, chlorates, acetates, and hydrogen carbonates are soluble.

3. All chlorides, bromides, and iodides are soluble except with silver(I), lead(II) and Hg2

2+.4. All sulfates are soluble except with lead(II), Hg2

2+, and alkali earths below magnesium.

5. All oxides and hydroxides are insoluble except with alkali earths below magnesium.

6. All phosphates, carbonates, sulfites, and sulfides are insoluble.

The ions chart is on p.57

Precipitation Reactions: Thinking synthetically

•A precipitation reaction demands that both reactants start as soluble compounds.•A precipitation reaction demands that at least one product ends as an insoluble compound.•Precipitation reactions are one way to synthesize insoluble salts.


Synthesizing an insoluble saltAB(aq) + CD(aq) AD(aq) + CB(s)

•C must be insoluble with B but soluble with D.•B must be soluble with A but insoluble with C.•A must be soluble with both B and D.•D must be soluble with both C and A.

Synthetic thinking is generally easier when you start with the names and then work to the formulas.

What would be good choices for A and D?


• Arrhenius Acid:

• A substance which

dissociates to form

hydrogen ions (H+)

in solution.

Acid-Base Reactions

• Arrhenius Base:

• A substance that

dissociates in, or

reacts with, water

to form hydroxide

ions (OH–).

Acid-Base Reactions

• Arrhenius Acid: • A substance which dissociates to

form hydrogen ions (H+) in solution.

• This H+ ion attaches to water and forms a hydronium ion (H3O+).

• In general, this means any anion with hydrogens attached as cations.

Acid-Base Reactions• Arrhenius Base: • A substance that dissociates in, or

reacts with, water to form hydroxide ions (OH–).

• In general, this means metal

hydroxides.

Acid-Base Reactions

• Brønsted Acid: Can donate protons (H+) to another substance in solution.

• Brønsted Base: Can accept protons (H+) from another substance in solution.

Acid-Base Reactions: Which hydrogens?

The only hydrogens which can behave acidically (that is, dissociate as a H+ ion) are hydrogens written

as if they were cations.

Acetic acid (or hydrogen acetate): HC2H3O2

AcidicNot acidic

Acid-Base Reactions: Polyprotic Acids

•Anions with a negative charge grater than -1 require more than one acidic hydrogen to form a neutral compound (acid).•These acids are called polyprotic (diprotic, triprotic, et cet.)•It is possible to remove only one of the multiple acidic hydrogens. In that case, the created anion is itself acidic.

Acid-Base Reactions

Acid-Base Reactions

• Neutralization Reaction: produces salt & water.

– HA(aq) + MOH(aq) H2O(l) + MA(aq)

• Write ionic and net ionic equations for the following:

– (a) Ca(OH)2(aq) + 2 HC2H3O2(aq)

– (b) HBr(aq) + Ba(OH)2(aq)

– (c) HCl(aq) + NH3(aq)

Acid-Base Reactions: Thinking synthetically

•Acid-base neutralization reactions are good methods to create solutions of soluble salts.•The cation of the salt is the metal in the metal hydroxide (base).•The anion of the salt is the anion of the acid.•The driving force of this reaction is so strong that the base can be insoluble and the reaction still works.

lOHaqClOFesOHFeaqHClO 22424 22

As before, it is easier to start with the name and then go to the formulas.

Acid-Base Reactions: Polyprotic Acids

•Anions with a negative charge grater than -1 require more than one acidic hydrogen to form a neutral compound (acid).•These acids are called polyprotic (diprotic, triprotic, et cet.)•It is possible to remove only one of the multiple acidic hydrogens. In that case, the created anion is itself acidic.

24442 SOHHSOaqHaqSOH

Sulfuric acid

Hydrogen sulfate

sulfate

Acid-Base Reactions

•Strong acids and strong bases are strong electrolytes. They completely dissociate in water.•Weak acids and bases only partially dissociate in water. But an acid-base neutralization reaction will drive them to completely dissociate during the reaction.

Oxidation NumbersAssigning Oxidation Numbers: All atoms have an “oxidation number”

regardless of whether it carries an ionic charge.1. An atom in its elemental state has an oxidation number of zero.

2. An atom in a monatomic ion has an oxidation number identical to its charge.

Oxidation Numbers

• Assigning Oxidation Numbers: All atoms have an “oxidation number” regardless of whether it carries an ionic charge.

1. An atom in its elemental state has an oxidation number of zero.

2. An atom in a monatomic ion has an oxidation number identical to its charge.

Oxidation Numbers

3. An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion.

–A. Hydrogen can be either +1 or –1.

–B. Oxygen usually has an oxidation number of –2.•In peroxides, oxygen is –1.

–C. Halogens usually have an oxidation number of –1.•When bonded to oxygen, (chlorine, bromine, and iodine) have positive oxidation numbers.

Oxidation Numbers4. The sum of the oxidation numbers must be zero for a neutral

compound and must be equal to the net charge for a

polyatomic ion.

–A. H2SO4 2(+1) + (?) + 4(–2) = 0 net charge

– ? = 0 – 2(+1) – 4(–2) = +6

–B. ClO4– (?) + 4(–2) = –1 net charge

– ? = –1 – 4(–2) = +7That is, there is often one element in a compound that does not have an oxidation number determined from rules 1-3. You use this rule to figure out that element’s oxidation number.

Identifying Redox Reactions

Identifying redox reactions systematically is simple, but time-consuming:

1. Assign oxidation numbers for all atoms in a reaction

2. Determine if any element’s oxidation number has changed.

Remember that there must be both oxidation and reduction in the same reaction. If you find only one, you are wrongety-wrong-wrong.

Activity Series•The more active the metal, the higher up it appears on this list.•A more active metal is able to force its electrons onto a less active metal’s cation, thus forcing it out of solution.•That is, an active metal is a good reducing agent.

•The two half-reactions “happen clockwise.”

Activity Series

• Activity series looks at the relative reactivity of a free metal with an aqueous cation.

– Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

– Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

– Cu(s) + 2 Ag+(aq) 2 Ag(s) + Cu2+(aq)

– Mg(s) + 2 H+(aq) Mg2+(aq) + H2(g)

Activity Series

Given the following three reactions,

determine the activity series for Cu, Zn, &

Fe.

Fe(s) + Cu2+(aq) Fe2+(aq) + Cu(s)

Zn(s) + Cu2+(aq) Zn2+(aq) + Cu(s)

Fe(s) + Zn2+(aq) NR

Activity Series – Thinking Synthetically

• The activity series provides a new route towards making solutions of soluble salts, and a route toward making pure elemental metals.

Example: How could we synthesize a solution of aluminum chlorate? Start with a solution of metal chlorate (where the metal is less active than aluminum) and then add an aluminum ingot.

Al(s) + Ni(ClO3)2(aq) Al(ClO3)3(aq) + Ni(s)

And then balance the reaction.

In actual practice, the pressure, temperature, and concentration conditions matter a great deal.

Balancing Redox Equations: The Challenge

OHMnFeHMnOFe 223

42 4585

OHMnFeHMnOFe 223

42 4282

The correctly balanced redox equation is:

This redox equation is incorrectly balanced:

Both versions are balanced for mass, but only one has the charge balanced on both sides. The rules we are about to develop exist so that we have a technique that balances both mass and charge.

Balancing Redox: Oxidation Number Method


33

272

IOCr

IOCr


33

272

2 IOCr

IOCr


33

272

2 IOCr

IOCr

-2+6 -1

+3 -2+5


33

272

2 IOCr

IOCr

-2+6 -1

+3 -2+5

Reduction: -6Oxidation: +6


33

272

2 IOCr

IOCr

-2+6 -1

+3 -2+5

So we leave it untouched

1623


OHIOCr

HIOCr

233

272

42

8


OHIOCr

HIOCr

233

272

42

8

2 Cr 2

7 O 7

1 I 1

8 H 8

+5 +/- +5

Balancing Redox: ½ reaction method


33

272

IOCr

IOCr

Balancing Redox: ½ reaction method 32

72 CrOCr

3IOI

Balancing Redox: ½ reaction method 32

72 2 CrOCr

3IOI


OHCrHOCr 232

72 7214

HIOOHI 63 32


OHCr

eHOCr

23

272

72

614

eHIO

OHI

66

3

3

2


Matt will write out this step; ain’t enough space here to do it justice.

Balancing Redox: ConditionsThe rules you just learned assume that the redox reaction is taking place under acidic conditions. (You are, after all, either producing or consuming H+ ions.) There are slightly different rules for basic conditions:

1. Balance the reaction (using your method of choice) as if it were under acidic conditions.

2. Add one hydroxide ion per H+ ion to both sides of the reaction.

3. On the side with H+ and hydroxide, they become water. The other side simply gains hydroxide.

4. Be sure to check if you now have water on both sides of the reaction which could be cancelled.

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Why do reactions occur at all?