Solution Stoichiometry

The 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 KI

1 mol KIx

2.80 mol KI

1 L solnx

1 L

1000 mLx

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

Dilution

Add Solvent

Moles of solutebefore dilution (i)

Moles of soluteafter dilution (f)=

MiVi MfVf=

How would you prepare 60.0 mL of 0.200 MHNO3 from a stock solution of 4.00 M HNO3?

MiVi = MfVf

Mi = 4.00 Mf = 0.200 Vf = 0.0600 L Vi = ? L

Vi =MfVf

Mi

=0.200 x 0.0600

4.00= 0.00300 L = 3.00 mL

3.00 mL of acid + 57.0 mL of water = 60.0 mL of solution

TitrationsIn a titration a solution of accurately known concentration is 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) the equivalence point

Slowly add baseto unknown acid

UNTIL

the indicatorchanges color

What volume of a 1.420 M NaOH solution isRequired to titrate 25.00 mL of a 4.50 M H2SO4 solution?

WRITE THE CHEMICAL EQUATION!

volume acid moles acid moles base volume base

H2SO4 + 2NaOH 2H2O + Na2SO4

4.50 mol H2SO4

1000 mL solnx

2 mol NaOH

1 mol H2SO4

x1000 ml soln

1.420 mol NaOHx25.00 mL = 158 mL

M

acid

rx

coef.

M

base

What volume of a 0.800 M HCl solution isrequired to titrate 40.00 mL of a 1.600 M NH3

solution?

HCl + NH3 NH4Cl

1.600 mol NH3

1000 mL solnx

1 mol HCl

1 mol NH3

x1000 ml soln

0.800 mol HClx40.00 mL = 80.0 mL

What volume of a 1.200 M Ba(OH)2 solution isrequired to titrate 90.00 mL of a 3.600 M H3PO4 solution?

3Ba(OH)2 + 2H3PO4 6H2O + Ba3(PO4)2

3.600 mol H3PO4

1000 mL solnx

3 mol Ba2(OH)

2 mol H3PO4

x1000 ml soln

1.200 mol Ba2(OH)x90.00 mL = 405.0 mL

Reactions in Aqueous SolutionChapter 4

•71% of Earth’s surface is water•Another 3% is covered with ice•66% of human body is water

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

Solution Solvent Solute

Soft drink (l)

Air (g)

Soft Solder (s)

H2O

N2

Pb

Sugar, CO2

O2, Ar, CH4

Sn

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 electrolyte

Strong Electrolyte – 100% dissociation

NaCl (s) Na+ (aq) + Cl- (aq)H2O

Weak Electrolyte – not completely dissociated

CH3COOH CH3COO- (aq) + H+ (aq)

Electrolyte Conducts electricity in solution?

Cations (+) and Anions (-)

Nonelectrolyte does not conduct electricity?

No cations (+) and anions (-) in solution

C6H12O6 (s) C6H12O6 (aq)H2O

http://www.chem.ufl.edu/~chm2040/cgi-bin/image_eval.cgi?Op=display&Chapter=4&File=Coolstuff/ionic&Image=c4f8.gif

Table 1 Electrolyte Classification of Some Common SubstancesStrong Electrolytes Weak Electrolytes Nonelectrolytes

HCl, HBr, HI CH3COOH H2O

HClO4 HF CH3OH

HNO3 C2H5OH

H2SO4 C12H22O11(sucrose)

KBr Most organic compounds

NaCl

NaOH, KOHOther soluble ionic compounds

E.g. Calculating ion concentrations in a solution of strong electrolyte. What are the

concentration (in molarity) of Mg2+ and Cl- ions in 5.0g/L of MgCl2 solution?

[Mg2+] = 5.0g x 1 molMgCl2 x 1 mol Mg2+ =5.3 x10-2 mol/L L 95.21gMgCl2 1 molMgCl2

[Cl-] =2 x 5.3x10-2 = 0.11 mol/L

Concentration of Ions

A bottle labeled as 0.100 M Al2(SO4)3.

[Al3+] = _____ M (mol / L)

[SO42–] = _____ M

Assume sea water is 0.438 M NaCl, 0.0512 M MgCl2, and 0.001 M CaCl2

[Na+] = _____ M

[Mg2+] = _____ M

[Ca2+] = _____ M

[Cl–] = _____ M Know how to calculate your quantities

Aqueous Reactions

Aqueous reactions can be grouped into three general categories, each with its own kind of driving force:

1. Precipitation reactions 2. Acid base neutralization reactions 3. Oxidation-reduction reactions.

Precipitation ReactionsWhen ions form a solid that is not very soluble, a solid is formed.

Such a phenomenon is called precipitation.

Precipitate – insoluble solid that separates from solution

PbI2

Rules for converting molecular equations to ionic equations

• The rules for converting molecular equations to ionic equations follow:

• 1) Make sure the molecular equation is balanced• 2) Ionic substances indicated in the molecular equation as

dissolved in solution, such as NaCl(aq), are normally written as ions.

• 3) Ionic substances that are insoluble (do not dissolve) either as reactants or products (such as precipitate) are represented by formulas of the compounds

• 4) Molecular substances that are strong electrolytes, such as strong acids, are written as ions. Thus, HCl(aq) is written as H3O+(aq) + Cl-(aq) or as H+(aq) + Cl-(aq).

• 5) Molecular substances that are weak electrolytes or nonelectrolytes are represented by their molecular formulas.

1. All alkali metals (Group 1A) compounds are soluble.

Solubility Rules for Common Ionic Compounds In water at 25oC

2. All ammonium (NH4+) compounds are soluble.

3. All nitrate (NO3-), chlorate (ClO3

-), and perchlorate (ClO4-) compounds

are soluble.

4. Most hydroxides (OH-) are insoluble. The exceptions are barium hydroxide [Ba(OH)2], which is very soluble, calcium hydroxide [Ca(OH)2], which is slightly soluble, and previous examples.

5. Most compounds containing chlorides (Cl-), bromides (Br-), or iodides (I-), are soluble. The exceptions are those containing Ag+, Hg2

2+, and Pb2+.

7. All carbonates (CO32-), phosphates (PO4

3-), and sulfides (S2-) are insoluble, excepting previous rules.

6. Most sulfates (SO42-) are soluble, excepting previous rules. Calcium

sulfate [CaSO4] and silver sulfate [Ag2SO4] are slightly soluble. Barium sulfate [BaSO4], mercury (II) sulfate [HgSO4], and lead sulfate [PbSO4] are insoluble.

Writing Net Ionic Equations

1. Write the balanced molecular equation.

2. Write the ionic equation showing the strong electrolytes

3. Determine precipitate from solubility rules

4. Cancel the spectator ions on both sides of the ionic equation

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

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

-

Ag+ + Cl- AgCl (s)

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

write a net ionic equation for the reaction. (a) Al2(SO4)3 + NaOH

i) write down the reactants and interchange of anions to get

product

Al2(SO4)3 + 6NaOH2Al(OH)3 + 3Na2SO4

All common Na compounds are water soluble Na+ remain in

solution. The combination of Al3+ and OH- produce insoluble

Al(OH)3. Then the ionic equation is

2Al3+ +3SO4

2- + 6Na+ + 6OH-2Al(OH)3(s)+ 6Na+ + 3SO4

2-

The net ionic equation is :

Al3+ + 3OH-Al(OH)3(s)

Net Ionic Equations

Note that some ions appear on both side of equation. These ions go through the reaction unchanged- does not take part in the reaction. We called them spectator ions. We can cancel them from the equation. The resulting equation is a net ionic equation.

Ag+(aq) + I-(aq) AgI(s)

Net ionic EquationA net ionic equation is an equation that includes only the actual participants in a reaction, with each participant denoted by the symbol or formula that best represent it.

Precipitation Reactions

Precipitate – 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)

Precipitation Reactions

Ag+ (aq) + NO3– (aq) + Cs+ (aq) + I– (aq) AgI (s) + NO3

– (aq) + Cs+ (aq)

Ag+ (aq) + I– (aq) AgI (s) (net reaction)or

Ag+ + I– AgI (s)

Heterogeneous Reactions

Spectator ions

Soluble ions

Alkali metals, NH4+

nitrates, ClO4-,

acetate

Mostly soluble ions

Halides, sulfates

Mostly insoluble

Silver halidesMetal sulfides, hydroxidescarbonates, phosphates

Acids (according to Arrhenius)

• Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid.

• React with certain metals to produce hydrogen gas.

• React with carbonates (CO32-) and bicarbonates (HCO3

-) to produce carbon dioxide gas.

• Are electrolytes.

• Cause certain plant dyes to change color. For example, they turn litmus from blue to red.

• Have a bitter taste.

• Feel slippery (they make water insoluble organic molecules into water soluble molecules). Many soaps contain bases.

Bases (according to Arrhenius)

• Are electrolytes.

• Cause certain plant dyes to change color. For example, they turn litmus from red to blue.

Arrhenius acid is a substance that produces H+ (H3O+) in water

Arrhenius base is a substance that produces OH- in water

A Brønsted acid is a proton donorA Brønsted base is a proton acceptor

acidbase acid base

A Brønsted acid must contain at least one ionizable proton!

Monoprotic acids

HCl H+ + Cl-

HNO3 H+ + NO3-

CH3COOH H+ + CH3COO-

Strong electrolyte, strong acid

Strong electrolyte, strong acid

Weak electrolyte, weak acid

Diprotic acidsH2SO4 H+ + HSO4

-

HSO4- H+ + SO4

2-

Strong electrolyte, strong acid

Weak electrolyte, weak acid

Triprotic acidsH3PO4 H+ + H2PO4

-

H2PO4- H+ + HPO4

2-

HPO42- H+ + PO4

3-

Weak electrolyte, weak acid

Weak electrolyte, weak acid

Weak electrolyte, weak acid

2. Neutralization Reaction

acid + base salt + water

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

H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O

H+ + OH- H2O

net ionic equation for strong acid / strong base reaction

Neutralization Reaction

acid + base salt + water

HCl (aq) + NH3 (aq) NH4Cl(aq)

H+ + Cl- + NH3(aq) NH4+ + Cl-

H+ + NH3(aq) NH4+

net ionic equation for strong acid / weak base reaction

Oxidation is the lose of one or more electrons by a substance. Historically oxidation is combination of an element with oxygen

Reduction is the gain of one or more electrons by a substance. Reduction is referred to the removal of oxygen from an oxide

A redox reaction is a process in which electrons are transferred between substance or in which atoms change oxidation number.

Oxidizing agent: contains an element whose oxidation state decreases in a redox reaction (it make possible for some other substance to be oxidized and itself reduced.) gains electrons (electrons are found on the left side of its half-equation).

Reducing agent: contains an element whose oxidation state increases in a redox reaction, loses electrons (electrons are found on the left side of its half-equation).

Oxidation-Reduction Reaction

Oxidation number

The charge the atom would have in a molecule (or anionic compound) if electrons were completely transferred.

1. Free elements (uncombined state) have an oxidation number of zero.

Na, Be, K, Pb, H2, O2, P4 = 0

2. In monatomic ions, the oxidation number is equal to the charge on the ion.

Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2

3. The oxidation number of oxygen is usually –2. In H2O2

and O22- it is –1.

4. The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1.

6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.

5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.

HCO3-

O = -2 H = +1

3x(-2) + 1 + ? = -1

C = +4

Oxidation numbers of all the elements in HCO3

- ?

Fig. 4.10

NaIO3

Na = +1 O = -2

3x(-2) + 1 + ? = 0

I = +5

IF3

F = -1

3x(-1) + ? = 0

I = +3

K2Cr2O7

O = -2 K = +1

7x(-2) + 2x(+1) + 2x(?) = 0

Cr = +6

Oxidation numbers of all the elements in the following ?

Oxidation-Reduction Reactions(electron transfer reactions)

2Mg (s) + O2 (g) 2MgO (s)

2Mg 2Mg2+ + 4e-

O2 + 4e- 2O2-

Oxidation half-reaction (lose e-)

Reduction half-reaction (gain e-)

2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e-

2Mg + O2 2MgO

Fe (s) + 2HCl (aq) FeCl2 (aq) + H2 (g)

Fe is oxidizedFe Fe2+ + 2e-

H+ is reduced2H+ + 2e- H2

Fe is the reducing agent

H+ is the oxidizing agent

Iron (Fe) reacting with hydrochloric acid (HCl)

An oxidizing agent: contains an element whose oxidation state decreases in a redox reaction (it make possible for some other substance to be

oxidized and itself reduced.) gains electrons (electrons are found on the left side of its half-equation).

A reducing agent: contains an element whose oxidation state increases in a redox reation loses electrons (electrons are found on the left side of its

half-equation).

Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)

Zn is oxidizedZn Zn2+ + 2e-

Cu2+ is reducedCu2+ + 2e- Cu

Zn is the reducing agent

Cu2+ is the oxidizing agent

Copper wire reacts with silver nitrate to form silver metal.What is the oxidizing agent in the reaction?

Cu (s) + 2AgNO3 (aq) Cu(NO3)2 (aq) + 2Ag (s)

Cu Cu2+ + 2e-

Ag+ + 1e- Ag Ag+ is reduced Ag+ is the oxidizing agent

Cu is oxidized Cu is the reducing agent

Ca2+ + CO32- CaCO3

NH3 + H+ NH4+

Zn + 2HCl ZnCl2 + H2

Ca + F2 CaF2

Precipitation

Acid-Base

Redox (H2 Displacement)

Redox (Combination)

Classify the following reactions.

Chapter 4: Reactions in Aqueous Solutions