Chapter 4; Reactions in Aqueous Solutions

I. Electrolytes vs. NonElectrolytesII. Precipitation Reaction

a) Solubility Rules

III. Reactions of Acidsa) Neutralizationb) Acid and Carbonatec) Acid and Metal

IV. Metal and Water

Chapter 4; Reactions in Aqueous Solutions

V. Redox ReactionsVI. MolarityVII. Acid-Base TitrationsVIII. Dilution of Solutions

4.1

A solution is a homogenous mixture of 2 or more substances

The solute is(are) the substance(s) present in the smaller amount(s)

The solvent is the substance present in the larger amount

In aqueous solutions (aq)*solvent is water*solute can be ionic compounds, aqueous acids, bases, or molecular compounds

An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.

A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.

nonelectrolyte weak electrolyte strong electrolyte4.1

4 Types of Inorganic Compounds

1. Molecular; Made of 2 or more Nonmetals2. Ionic; Made of + and – ion. Generally +

ion is from metal and – ion from nonmetal.

3. Bases; + ion; - ion is hydroxide (OH)-

4. Aqueous Acid; H+ and – ion dissolved in water. Generally – ion is nonmetal

Inorganic Compounds Dissolved in WaterALL BUT MOLECULAR PRODUCE IONS

Na+ Cl-

1. Ionic Compounds Na Cl (aq)

Na+ (OH)-

2. Bases Na (OH) (aq)

3. Aqueous AcidsHCl (aq)

H+ Cl-

4. Molecular ICl (aq)

ClI

NO Ions!

4.1

Strong Electrolyte

Weak Electrolyte Nonelectrolyte

Strong Acids Weak Acids Molecular Compounds

Strong Bases Weak Bases

Ionic Compounds

Electrolytic Solutions Contain Mobile Cations (+) and Anions(-)

• Ionic Compounds, Aqueous Acids, and Base Dissociate Into the Ions They are Made of When Dissolved in Water.

NaCl (s) Na+ (aq) + Cl-(aq)

• More Ions in Solutions; Stronger Electrolyte

H2O

Strong Electrolyte – 100% dissociation

HCl (g) H+ (aq) + Cl- (aq)H2O

Weak Electrolyte – not completely dissociated

Weak vs. Strong ElectrolyteA strong electrolyte will produce more ions when same amount of solid is dissolved in solvent.

HNO2 NO2- (aq) + H+ (aq)

H2O

H+ Cl-

H+Cl-

H+ NO2-

H+NO2-

Precipitation Reactions

• Mix two aqueous solutions made by dissolving ionic compounds in water.

• If a reaction happens, a precipitate (solid) is formed.

Double ReplacementMETATHESIS Reactions

AX (aq) + BY(aq) AY (s) + BX (aq)

Ions are removed in one of three ways:1. *Formation of an insoluble product (precipitate)2. Formation of either a weak electrolyte or a

nonelectrolyte3. Formation of a gas that escapes from solution

Predicting Products of Precipitation Reactions

1) Ionic Compounds are Strong Electrolytes –Determine charge on all ions of reactants

2) Using Ion Charges; Predict formula of products. ( + ion of one reactant forms compound with – ion of other reactant)

3) Balance Equation (Molecular Equation)4) Determine is product is solid or aqueous

solution

Predicting Products of Precipitation Reactions (Cont)

5) Determine spectator ions (Ions that are still dissolved in water in the product) -Ionic Equations

6) Write net ionic equation (Only shows ions involved in forming solid)

Solubility Rules for Common Ionic CompoundsIn water at 250C

Soluble Compounds ExceptionsCompounds containing alkali metal ions and NH4

+

NO3-, HCO3

-, ClO3-

Cl-, Br-, I- Halides of Ag+, Hg22+, Pb2+

SO42- Sulfates of Ag+, Ca2+, Sr2+, Ba2+,

Hg2+, Pb2+

Insoluble Compounds Exceptions

CO32-, PO4

3-, CrO42-, S2- Compounds containing alkali

metal ions and NH4+

OH- Compounds containing alkali metal ions and Ba2+

4.2

S A P C A NI. 100% (aq) with any partnerSodium Na+ Chlorate (ClO3)-

Ammonium NH4+ Acetate (C2H3O2)-

Potassium K+ Nitrate (NO3)-

II. Aq only without these partnersHalides ( F-, Cl-, Br-, I-)

NO Ag+, Hg2+ or Pb2+

Sulfate (SO4)-2

NO Ag+, Hg2+ or Pb2+

NO Ca2+ , Sr2+ or Ba2+

Precipitation ReactionsPrecipitate – insoluble solid that separates from solution

molecular equation

ionic equation

net ionic equation

Pb2+ + 2NO3- + 2Na+ + 2I- PbI2 (s) + 2Na+ + 2NO3

-

Na+ and NO3- are spectator ions

PbI2

Pb(NO3)2 (aq) + 2NaI (aq) PbI2 (s) + 2NaNO3 (aq)

precipitate

Pb2+ + 2I- PbI2 (s)

4.2

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

Ag+ + NO3- + Na+ + Cl- AgCl (s) + Na+ + NO3

-

Ag+ + Cl- AgCl (s)4.2

Write the net ionic equation for the reaction of silver nitrate with sodium chloride.

Acids

Have a sour taste. Vinegar owes its taste to acetic acid. Citrusfruits contain citric acid.

React with certain metals to produce hydrogen gas.

React with carbonates and bicarbonates to produce carbon dioxide gas

Have a bitter taste.

Feel slippery. Many soaps contain bases.

Bases

4.3

Acids Produce H+ (proton) or (H3O)+ when

dissolved in water Proton donor

HNO3 (aq) + H2O (l) H3O+ (aq) + (NO3)- (aq)

HNO3 (aq) H+ (aq) + (NO3)- (aq)H2O

Monoprotic acids; Produce one H+ when dissolved in water. i.e. one ionizable proton.

HNO3 H+ + NO3-

Diprotic acids; Produce two H+ when dissolved in water (polyprotic). i.e. two ionsizable protons

H2SO4 (aq) H+ + HSO4-

HSO4- H+ (aq) + SO4

2-(aq)

4.3

Triprotic acids; Produce three H+ when dissolved in water (polyprotic)

H3PO4 3 H+ + PO4-3

7 Strong AcidsSulphuric Acid H2SO4

Hydrochloric Acid HCl

Hydrobromic Acid HBr

Hydroiodic Acid HI

Nitric Acid HNO3Chloric Acid HClO3Perchloic Acid HClO4

Bases

•Produce (OH)- when dissolved in water•Proton (H+) acceptor

F- (aq) + H2O (l) <-> HF (aq) + (OH)- (aq)

Na(OH) (s) -----> Na+ (aq) + (OH)- (aq)H2O

Strong BasesLithium Hydroxide LiOH

Sodium Hydroxide NaOH

Potassium Hydroxide KOH

Calcium Hydroxide Ca(OH)2

Stontium Hydroxide Sr(OH)2

Barium Hydroxide Ba(OH)2

Cesium Hydroxide CsOHRubidium Hydroxide RbOH

Neutralization ReactionAcid + Base -> Salt + H2O

A salt is any compound that is formed from the anion of an acid

and a cation from a base. HCl (aq) + NaOH (aq) H2O(l) + NaCl (aq)

Cl- from acid and Na+ from Base

NaCl = salt

Strong Acid + Strong Base Neutral Salt

Strong Acid + Weak Base Acidic Salt

Weak Acid + Strong Base Basic Salt

“Weak Acid” + “Weak Base” Stronger Will Prevail!

Gas-Forming Reactions

• These metathesis reactions do not give the product expected.

• The expected product decomposes to give a gaseous product (CO2 or SO2).

CaCO3 (s) + HCl (aq) CaCl2 (aq) + CO2 (g) + H2O (l)NaHCO3 (aq) + HBr (aq) NaBr (aq) + CO2 (g) + H2O (l)

SrSO3 (s) + 2 HI (aq) SrI2 (aq) + SO2 (g) + H2O (l)

Gas-Forming Reactions

• This reaction gives the predicted product, but you had better carry it out in the hood, or you will be very unpopular!

• Just as in the previous examples, a gas is formed as a product of this reaction:

Na2S (aq) + H2SO4 (aq) Na2SO4 (aq) + H2S (g)

Common Gas Forming ReactionsH2S Any sulfide (S2-) + any acid form H2S(g) and

a salt

CO2 Any carbonate (CO32-) + any acid form

CO2(g), H2O and a salt

SO2 Any sulfite (SO32-) + any acid form SO2(g),

H2O, and a salt (Sulfur smell)

NH3 Any ammonium salt (NH4+) + a soluble strong hydroxide react upon heating to form NH3(g), H2O and a salt (Ammonia smell)

Acid-Base Reaction Identifiers

• When a base is added to an NH4- solution

– A distinctive ammonia odor will be detected• When an acid is added to a solution

containing S2-

– A rotten-egg odor will be detected (H2S)• When acid is added to CO3

2- or CO2– Gas is produced.

Oxidation-Reduction Reactions

• An oxidation occurs when an atom or ion loses electrons.

• A reduction occurs when an atom or ion gains electrons.

Oxidation-Reduction ReactionsOxidizing agent – A

substance that causes oxidation of another substance. (It will be reduced)

Reducing agent – A substance that causes the reduction of another substance. (It will be oxidized)

One cannot occur without the other.

REDOX

O = OxidationI = IsL = Loss of e-

R = ReductionI = IsG = Gain of e-

L = LosingE = Electrons O = is OxidationSaysG = GainingE = ElectronsR = is Reduction

Make a Soda Can Disappear orWhy doesn’t soda “eat” away the can?

A vigorous reaction is observed - heat is released, the blue color due to copper(II) ions fades, the aluminum can disintegrates, and a new, reddish brown solid appears in the reaction mixture. The reaction is summarized in Equation 1. Copper(II) ions from copper(II) chloride are converted to copper metal, and aluminum metal is converted to aluminum cations in aluminum chloride, which is soluble in water. In contrast, when a piece of copper metal is placed in a solution of aluminum chloride, no reaction takes place (Equation 2).

2Al(s) + 3CuCl2(aq) → 2AlCl3(aq) + 3Cu(s) Equation 1Cu(s) + AlCl3(aq) → No Reaction Equation 2

A more active metal always replaces the ion of a less active metal in a compound.

In general, the activity of a metal may be defined as follows: An active metal will react with a compound of a less active metal, which is converted to its “free element” form.

The more active metal forms a new compound containing metal cations. Based on Equation 1, aluminum is more active than copper.

Nitric Acid & Copper DemoCopper is oxidized by concentrated nitric acid, HNO3, to produce Cu2+ ions.

Nitric acid is reduced to nitrogen dioxide, a poisonous brown gas with an irritating odor.

Cu(s) + 4HNO3(aq) ——> Cu(NO3)2(aq) + 2NO2(g) + 2H2O(l)

When the copper is first oxidized, the solution is very concentrated, and the Cu2+ product is initially coordinated to nitrate ions from the nitric acid, giving the solution first a green, and then a greenish-brownish color. When the solution is diluted with water, water molecules displace the nitrate ions in the coordinate sites around the copper ions, causing the solution to change to a blue color.

Electronegativity

A property of an atom which increases with its tendency to attract the electrons of a bond.

Examples: The chlorine atom has a higher electronegativity than the hydrogen atom, so the bonding electrons will be closer to the Cl than to the H in the HCl molecule.

Electrochemical reactions involve the transfer of electrons.

Mass and charge are conserved when balancing these reactions, but you need to know which atoms are oxidized and which atoms are reduced during the reaction.

Oxidation numbers are used to keep track of how many electrons are lost or gained by each atom. These oxidation numbers are assigned using the following rules:

The convention is that the cation is written first in a formula, followed by the anion.

For example, in NaH, the H is H-; in HCl, the H is H+.

The oxidation number of a free element is always 0.

The atoms in He and N2, for example, have oxidation numbers of 0.

The oxidation number of a monatomic ion equals the charge of the ion.

For example, the oxidation number of Na+ is +1; the oxidation number of N3- is -3.

The usual oxidation number of hydrogen is +1.**The oxidation number of hydrogen is -1 in compounds containing elements that are less electronegative than hydrogen, as in CaH2.

The oxidation number of oxygen in compounds is usually -2.

**Exceptions include OF2, since F is more electronegative than O, and BaO2, due to the structure of the peroxide ion, which is [O-O]2-.

The oxidation number of a Group IA element in a compound is +1.

The oxidation number of a Group IIA element in a compound is +2.

The oxidation number of a Group VIIA element in a compound is -1, except when that element is combined with one having a higher electronegativity.

The oxidation number of Cl is -1 in HCl, but the oxidation number of Cl is +1 in HOCl.

The sum of the oxidation numbers of all of the atoms in a neutral compound is 0.

The sum of the oxidation numbers in a polyatomic ion is equal to the charge of the ion.

For example, the sum of the oxidation numbers for SO4

2- is -2.

Which are Redox Reactions?

Redox = Transfers of electrons

Single Replacement & Combustion = REDOX (*metals are based on activity series)

Combo & Decomp = Sometimes (?)

Acid/Base Reactions = Most of the times

Elements in natural state = ALWAYS REDOX

Balancing Redox1. Assign oxidation numbers to all atoms in the equation

2. Identify the substances being oxidized and determine the # of e- lost

3. Identify the substances being reduced and determine the # of e-

gained

4. Use coefficients to balance the atoms in the substances oxidized and reduced

5. Use coefficients to balance the electrons gained and lost

6. Use coefficients to balance the non redox substances.

Ex: CdS + I2 + HCl CdCl2 + HI + S1. Ox #’s: Cd2+ S2- + H+ Cl- Cd2+ Cl2- + H+ I- + S0

2. Oxidation: Sulfur changes froma 2- to a 0 (2 lost e-)

3. Reduction: Iodine changes form 0 to a -1 (1 gained e-)

4. Balance atoms: I is diatomicCdS + I2 +HCl CdCl2 + 2 HI + S

5. Balance e’s: S loses 2 e- and each I gains 1e- (all balanced)

6. Balance atoms: Balance the H and Cl atoms

CdS + I2 + 2 HCl CdCl2 + 2 HI + S

Displacement Reactions – Metal Displaces H from acid or waterBasic Reaction: A + BX AX + B Metal + Acid -> Salt + H2 (g)

Metal + Water -> Base + H2(g)

Metal + Salt **Use Activity Series to Know if a Reaction Will Happen

Displacement Reactions

• In displacement reactions, ions oxidize an element.

• The ions, then, are reduced.

Displacement Reactions

In this reaction, silver ions oxidize copper metal.

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

Displacement Reactions

The reverse reaction, however, does not occur.

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

Single Replacement / Displacement Reactions

Single Replacement/ Displacement Reactions A more active element replaces a less active element in a compound.

Activity series for metals:

Most act. Li Ca Na Mg Al Zn Fe Pb [H2] Cu Ag Pt Least act.

Activity series for non-metals Most active F2 Cl2 Br2 I2 Least active

*Note that the activity series can also be determined by looking at the standard electrode potentials

1. Active metal replaces hydrogen in water or an acid (Hydrogen Displacement)

a. metal + acid salt and hydrogen gas • e.g. Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g) (Net : Mg + 2H+ Mg2+ + H2)

b. Group 1A or 2A metal and waterbase and hydrogen gas • e.g. 2Li(s) + 2H2O(l) 2LiOH(aq) + H2(g) (Net : Li + 2H2O Li+ + OH- + H2)

2. a more active metal replaces a less reactive metal ion from solution(see activity series) (metal displacement)

• e.g. ZnCl2(aq) + Mg(s) MgCl2(aq) + Zn(s) ( Net : Zn2+ + MgMg2+ + Zn )

3. a halogen replaces another halogen

• e.g. MgBr2(aq) + F2(g) MgF2(aq) + Br2(l) ( Net : 2 Br - + F2 2 F - + Br2 )

4. Hydrides of alkali metals react with water to form hydroxides :

• e.g. LiH(s) + H2O(l) LiOH + H2(g) (Net : LiH + H2O Li+ + OH- + H2)

Activity Series

Displace H2 from H2O, Steam & Acid

Displace H2 fromSteam & Acid

Displace H2 fromAcids only

Can’t Displace H2

Activity Series

ElectrochemistryVoltaic Cell – Makes a Battery

• Positive Electrode = CathodeReduction Takes place here(RED CAT)

• Negative Electrode = AnodeOxidation takes place here (AN OX)

• Salt Bridge/Porous Bride allows for the flow of ions between solutions

Solution StoichiometryThe concentration of a solution is the amount of solute present in a given quantity of solvent or solution.

M = molarity =moles of solute

liters of solution

What mass of KI is required to make 500. mL ofa 2.80 M KI solution?

volume KI moles KI grams KIM KI M KI

500. mL = 232 g KI166 g KI1 mol KI

x2.80 mol KI

1 L solnx

1 L1000 mL

x

4.5

4.5

Acid/Base Titrations

• Experimental technique that determines the concentration (in Molarity) of an acid (or base)

• This is based upon an acid/base neutralization reaction.– ACID +BASE -> SALT + H2O

• Base (or acid) is added until there is the same amount (same # moles) of base and acid.

TitrationsIn a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.

Equivalence point – the point at which the reaction is complete

Indicator – substance that changes color at (or near) theequivalence point

Slowly add baseto unknown acid

UNTILthe indicator

changes color4.7

Fig. 4.17a,b

Acid-Base Titrations

Introductory Chemistry 2/e by N Tro,Prentice Hall, 2006, pg 480

Base; (OH)-

Acid; H+

Acid + Base -> Salt + H2O

At the endpoint of an acid/base titration….

• Moles acid = Moles base• (MV)acid = (MV)base

• Note– If solid; moles = mass/ MM– If aqueous solution; moles = MV

What volume of a 1.420 M NaOH solution isRequired to titrate 25.00 mL of a 4.50 M H2SO4solution?

4.7

WRITE THE CHEMICAL EQUATION!

volume acid moles acid moles base volume base

H2SO4 + 2NaOH 2H2O + Na2SO4

M

acid

rx

coef.

M

base

4.50 mol H2SO4

1000 mL solnx

2 mol NaOH1 mol H2SO4

x1000 ml soln

1.420 mol NaOHx25.00 mL = 158 mL

Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.

DilutionAdd Solvent

Moles of solutebefore dilution (i)

Moles of soluteafter dilution (f)=

MiVi MfVf=4.5

How would you prepare 60.0 mL of 0.2 MHNO3 from a stock solution of 4.00 M HNO3?

MiVi = MfVf

4.5

Vi =MfVf

Mi= 0.200 x 0.06

4.00= 0.003 L = 3 mL

3 mL of acid + 57 mL of water = 60 mL of solution