chapter 4: types of chemical reactions - sherrill...

CHAPTER 4: Types of Chemical Reactions

•Dissolution •Precipitation•Acids and Bases and their reactions•Oxidation-Reduction Reactions

CHEM 1310 A/B Fall 2006

Dissolution of Ionic Compounds


Dissolution of Molecules


Electrolytes• Substances which

increase the conductivity of water when they dissolve

• Usually ionic compounds (e.g., NaCl), but some molecules can also dissolve into ions (e.g., acids like HCl, HNO3). Strong bases can also be electrolytes (NH3 can split H2O into ions).


Solubility

• Tells how much of something (the “solute”) will dissolve into a given solvent

• Example: incomplete dissolution of sugar in iced tea

• Observation: hot tea dissolves sugar much easier than iced tea. This implies…

• Another observation: Coke goes flat faster at higher temperatures. Why? Solubility of gases change …


Rules about solubility• Most general rule: “Like

dissolves like”• Polar solvents (water,

methanol, etc) usually dissolve polar molecules and ionic substances; nonpolar solutes not very soluble

• Nonpolar solvents (CCl4, benzene, etc) usually dissolve nonpolarmolecules; don’t dissolve polar molecules well

Polar ethylene glycol dissolves in water

Oil and water do not mixCHEM 1310 A/B Fall 2006

More rules about solubility• No gases or solids are infinitely soluble (“miscible in all

proportions”) with water; many liquids are• Common inorganic acids soluble in water; organic acids

which are small molecules soluble in water• Small organic compounds with –OH and –NH2 groups

usually soluble in water• Gases are less soluble as temperature increases.

Liquids and soluds usually more soluble as temperature increases

• Solubilities of ionic compounds varies in water; see Table 4-1


Precipitation Reactions• Opposite of dissolution --- a

solid comes out of solution Na+ (aq) + Cl- (aq) → NaCl(s)

• Can cause precipitation by (a) lowering temperature (e.g., “rock candy”), (b) evaporating solvent, (c ) mixing 2 or more solutions, e.g,

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

all of these are quite soluble in water except BaSO4

Precipitation of lead iodide, PbI2,from KI (aq) + Pb(NO3)2 (aq)


Ionic equations and net ionic equations

• The previous example, BaCl2 (aq) + K2SO4 (aq) → BaSO4 (s) + 2 KCl (aq), can be written in terms of ionic equations asBa2+ (aq) + 2 Cl- (aq) + 2 K+ (aq) + SO4

2- (aq) →BaSO4 (s) + 2 K+ (aq) + 2 Cl- (aq)

• Notice that 2 Cl- (aq) and 2 K+ (aq) appear on both sides of the equation. They “cancel” and we are left withBa2+ (aq) + SO4

2- (aq) → BaSO4 (s)

the net ionic equation. The Cl- and K+ ions are spectator ions


Predicting precipitation• In the previous example, how did we know that BaSO4

would precipitate out, and KCl would not?• Need to consider all possible precipitates, and consult

solubility tables to see if they remain dissolved or if they precipitate

• Table 4-1 in book gives rules of thumb• Could look up “solubility constants” to do this more

quantitatively --- more on this later in the course!• Example: what happens when an aqueous solution of

beryllium nitrate is mixes with an aqueous solution of sodium acetate?


Example


Example, cont’d


Acids & Bases• Acids

– Citric acid (lemons)– Acetic acid (CH3COOH, in vinegar)– Hydrochloric (HCl), sulfuric (H2SO4), nitric (HNO3)– Taste sour, can be very reactive

• Bases– In soaps (KOH, NaOH)– Drain openers (NaClO, NaOH)– Bleach (NaClO; ClO- is hypochlorite)– Taste bitter, are corrosive, degrade organic matter


Chemistry of acids & bases• All our example acids contained hydrogen• Several of our example bases contained hydroxide• An “Arrhenius acid” gives H+ ions when dissolved in

water• An “Arrhenius base” gives OH- ions when dissolved in

water• Water itself is both an Arrhenius acid and Arrhenius

base. Small amounts undergo this reaction:H2O (l) → H+ (aq) + OH- (aq)

• Note: the H+ ions stay closely surrounded by at least a few H2O molecules


Strong acids and bases• Dissociate completely in water (and are

therefore also strong electrolytes)• HCl (g) → H+ (aq) + Cl- (aq) [demo]• NaOH (s) → Na+ (aq) + OH- (aq)• Strong acids: HCl, HNO3, H2SO4, HClO3, HClO4,

HBr, HI• Strong bases: NaOH, KOH• Neutralization reaction: H+ from acid combines

with OH- from base, e.g.,HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)

What’s the “net ionic equation” for this reaction?


Weak acids• Do not completely dissociate in water; weak electrolytes• E.g., HF, acetic acid, CH3COOH, formic acid HCOOH,

phosphoric acid H3PO4, … [demo]• Although they don’t normally completely dissociate in

water, a strong base can still “force” them to fully dissociate, e.g.,CH3COOH (aq) + NaOH (aq) →H2O (l) + NaCH3COO (aq)most of the acetic acid remains undissociated before the reaction, so the net ionic equation isCH3COOH (aq) + OH- (aq) →H2O (l) + CH3COO- (aq)

• See book about naming acids


Weak bases

• The only strong bases are alkali metal hydroxides (e.g., NaOH, KOH) and Ba(OH)2

• Everything else is a weak base (incomplete dissociation in water unless forced by an acid)


Generalization of Arrhenius Model• Ammonia (NH3) tastes bitter, feels soapy,

neutralizes acids, just like a base --- but it does not contain OH- ions!

• Revised definition: An Arrhenius base, when dissolved in water, increases the amount of hydroxide ion over that present in pure solvent

• Definition of Arrhenius acid expanded similarly for H+ ion

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

[demo]


Acid & base anhydrides

• Substances which become H+ or OH- donators when or after they react with water– Acid anhydride: SO3 (g) + H2O (l) → H2SO4 (aq)– Base anhydride: Na2O (s) + H2O (l) → 2 NaOH (aq)

• Usually oxides of nonmetals are acid anhydrides• Usually oxides of metals are base anhydrides• (More electronegative → more acidic,

more electropositive → more basic)


Other reactions of acids• Besides neutralization reactions, acids…• React with carbonates and hydrogen carbonates

to liberate gaseous CO2, producing a salt and water, e.g., CaCO3 (s) + HCl (aq) →CaCl2 (aq) + H2O (l) + CO2 (g) [demo]

• act with oxides of metals to form salts and water, e.g.,CuO (s) + H2SO4 (aq) → ???

• React with various metals to liberate H2 and form a salt, e.g.,Mg (s) + 2 HCl (aq) → ?? [demo]


Other reactions of bases

• Bases (except NH3) act on ammonium salts to generate gaseous ammonia, a salt, and water. E.g.,NH4Cl (s) + NaOH (aq) →NaCl (aq) + NH3 (g) + H2O (l)Note: gaseous ammonia can be dangerous!

• Bases react with the oxides of non-metals to produce salts and water, e.g.,SO3 (g) + 2 NaOH (aq) → Na2SO4 (aq) + H2O (l)


Oxidation-Reduction (“Redox”) Reactions

• Example: Cu wire in AgNO3 (aq) [demo]. Why does this Rx happen? Why does the solution turn blue?

• Goals: to understand the twin prcesses of oxidation and reduction; assign “oxidation numbers”; become familiar with the basic types of redox reactions


A simple redox reaction

• Recall periodic table, electronegativity, octets

• 2 Na (s) + Cl2 (g) → 2 Na+ Cl- (s)each Na atom each Cl atom Na+ and Cl- bothloses an e- gains an e- have filled shells“oxidation” “reduction” like a noble gas…(“LEO”) (“GERtrude”) happy

• Transfer of electrons. To do bookkeeping on electrons transferred, use “oxidation numbers”


Oxidation numbersA bookkeeping device, like formal charges (but not the same)Book has a set of 5 rules…

1. Sum of oxidation numbers must equal total charge (=0 if neutral molecule)

2. Group I has oxidation number of +1 alwaysGroup II has +2 alwaysGroup III has +3 always

3. Halogens have o.n. -1 usually (always for F) except in compounds with O or other halogens

4. H has +1 except for metal hydrides like LiH, where it has -15. O has -2 except (a) in compounds with F, use rule 3; (b) in O-O

bonds, use rules 2, 4


Examples

• What are the oxidation numbers for H2O2, NaO2, ClO- ?


Oxidation and Reduction

• Use oxidation numbers to tell if an atom is oxidized or reduced by a reaction

• “Oxidizing agents” promote the oxidation of other compounds. The oxidizing agent itself must be reduced (e.g., MnO4

-)• Redox reactions can change a lot with

conditions (acidity, basicity, etc)• Some strong oxidizing/reducing agents can

react slowly, some weak ones can react rapidly


Example of reaction conditions

In acidic solution…5 Fe2+ (aq) + MnO4

- (aq) + 8 H+ (aq) →5 Fe3+ (aq) + Mn2+ (aq) + 4 H2O (l)Which is oxidized and which is reduced?

In neutral solution…3 Fe3+ (aq) + MnO4

- (aq) + 4 H+ (aq) →3 Fe2+ (aq) + MnO2 (s) + 2 H2O (l)


Types of redox reactions• Combination reactions: 2 elements combine to

form a compound. Metals usually lose electrons (are oxidized)2 Fe (s) + 3 Cl2 (g) → 2 FeCl3 (s) [demo]

• Decomposition reactions: Compound breaks up into elements or simpler compounds2 Ag2O → 4 Ag (s) + O2 (g)

• Oxygenation: Reaction of element or compound with oxygen (either O2 or O3, both oxidants)4 Li (s) + O2 (g) → 2 Li2O (s)


Redox reaction types, cont’d• Hydrogenation: H2 is a good reducing agent

2 Na (l) + H2 (g) → 2 NaH (s)In organic chem, H2 often adds to double or triple bonds, e.g.,H2C=O + H2 → H3C-OH

• Displacement:2 AgNO3 (aq) + Cu(s) →Cu(NO3)2 + 2 Ag (s) [from the demo!]

Can predict ability to displace using “activity series”(Table 4-2). Ag is a weaker reducing agent than Cu (not as good at giving up electrons)


More on displacement reactions

• The famous sodium in water Rx is also a “displacement” Rx..2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)Na “displaces” hydrogen

• Halogens displace “backwards”; they are oxidizing agents. Increasing oxidizing power: I2 < Br2 < Cl2 < F2

Cl2 (g) + 2 KI (aq) → I2 (g) + 2 KCl (aq)


Redox reaction types, cont’d 2

• Disproportionation: A single ion undergoes both oxidation and reduction2 H2O2 (l) → 2 H2O (l) + O2 (g)

O reduced O oxidized


Stoichiometry of Reactions in Solution

• Quantitative problems like in chapter 2, but now with the solution phase

• Usually easiest to measure volumes of solutions. Need to know how much of a compound for a given volume. “Concentration”Csolute = nsolute / Volumemolarity = moles of solute / liters of solution1.0 M = 1.0 mol L-1 = “1 molar”


Example of molarity

• What molarity is a solution made by dissolving 10.0 g of Al(NO3)3 in enough water to make 250. mL of solution?


Solution stoichiometry, cont’d

• Note the wording of the last example. Don’t dissolve solute in exactly 250 mL of solvent. Dissolve in less solvent, then add solvent until 250 mL of solution.

• Dilution: See book. Basically, to get solution 1/3 as concentrated, need to increase volume by a factor of 3.


Stoichiometry example

• When treated with acid, lead (IV) oxide is reduced to a lead (II) salt, liberating oxygen2 PbO2 (s) + 4 HNO3 (aq) →2 Pb(NO3)2 (aq) [reduced Pb!] + 2 H2O (l) + O2 (g)

• What volume of a 7.91 M solution of HNO3is needed to react with 15.9 g of PbO2?


Example, cont’d

• 2 PbO2 (s) + 4 HNO3 (aq) →2 Pb(NO3)2 (aq) + 2 H2O (l) + O2 (g)


Titrations• A quantitative way to

perform a reaction with solutions by measuring volume of one solution added to another using a buret

• Frequently add base from buret to an acidic solution until the indicator in the flask just barely turns color when it turns basic

• “End point” – when base has just neutralized the acid


Titrations, cont’d

• Can know the amount of base added (from accurate concentration plus the known volume added) and get acid concentration (from volume of acid in solution, plus the knowledge that at endpoint, moles of acid = moles of base)

• Or, sometimes “standardize” base concentration by titrating against known amount of acid


Example of titrationA 100. mL solution of hydrochloric acid is just neutralized (indicator turns just pink) when 20. mL of a 1.5 M solution of NaOH is added. What is the concentration of the HCl solution?


chapter 4: types of chemical reactions - sherrill...

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