Chapter 5

Dr. S. M. Condren

Chapter 5

Reactions

in

Aqueous Solutions

Dr. S. M. Condren

Many reactions involve ionic compounds, Many reactions involve ionic compounds, especially reactions in water — especially reactions in water —

aqueous solutions.aqueous solutions.KMnOKMnO44 in water in water KK++(aq) + MnO(aq) + MnO44

--(aq)(aq)

Reactions in Aqueous Solution

Dr. S. M. Condren

An Ionic Compound, CuCl2, in Water

CCR, page 177CCR, page 177

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How do we know ions are How do we know ions are present in aqueous present in aqueous solutions?solutions?

The solutions The solutions conduct electricity!

They are called They are called ELECTROLYTES

HCl, CuClHCl, CuCl22, and NaCl are , and NaCl are strong electrolytes. . They dissociate completely They dissociate completely (or nearly so) into ions.(or nearly so) into ions.

Aqueous Solutions

Dr. S. M. Condren

HCl, CuClHCl, CuCl22, and NaCl are , and NaCl are strong electrolytes.. They dissociate They dissociate completely (or nearly so) into ions.completely (or nearly so) into ions.

Aqueous Solutions

Dr. S. M. Condren

Aqueous SolutionsAcetic acid ionizes only to a small extent, so Acetic acid ionizes only to a small extent, so

it is a it is a weak electrolyte..CHCH33COCO22H(aq)H(aq) ---> ---> CHCH33COCO22

--(aq) + H(aq) + H++(aq)(aq)

Dr. S. M. Condren

Aqueous SolutionsAcetic acid ionizes only Acetic acid ionizes only

to a small extent, so it to a small extent, so it

is a is a weak electrolyte.

CHCH33COCO22H(aq)H(aq) ---> --->

CHCH33COCO22--(aq) + H(aq) + H++(aq)(aq)

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Aqueous SolutionsSome compounds Some compounds

dissolve in water but dissolve in water but do not conduct do not conduct electricity. They are electricity. They are called called nonelectrolytes.

Examples include:Examples include:sugarsugarethanolethanolethylene glycolethylene glycol

Examples include:Examples include:sugarsugarethanolethanolethylene glycolethylene glycol

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Electrolytes

• Conduct electricity in solution due to the presence of ions

• Strong electrolyte – completely ionized in solution

• Weak electrolyte – partially ionized in solution

• Non-electrolyte – nonionic solution

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Solubility Rules

1. All nitrates (NO3-1) are soluble.

2. All compounds of Group IA metals and the ammonium ion, NH4

+, are soluble.

3. All chlorides are soluble except: AgCl, Hg2Cl2 and PbCl2.

4. All sulfates are soluble except: PbSO4, BaSO4, and SrSO4.

Dr. S. M. Condren

Solubility Rules

5. All hydroxides (OH-1)and sulfides (S-2)are insoluble except those of the Group IA metals and the ammonium ion.

6. All carbonates (CO3-2) and phosphates

(PO4-3) are insoluble except those of the

Group IA metals and the ammonium ion.

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Net Ionic Equation

Balanced Chemical Equation

Pb(NO3)2 + Na2SO4 ---> PbSO4 + 2NaNO3

Total Ionic Equation

Pb+2 + 2NO3-1 + 2Na+1 + SO4

-2

2Na+ + 2NO3 -1 + PbSO4

Net Ionic Equation

Pb+2 + SO4-2 PbSO4

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A solution of Ba(NO3)2 is added to a solution of Na2SO4 to make a precipitate. From a table of solubility rules, the product is

barium sulfate, sodium nitrate

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Types of Reactions

• synthesis reactions or combination reactions

• decomposition reactions

• precipitation reactions

• neutralization reactions– acid– base

• oxidation-reduction reaction

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Synthesis or Combination Reactions

Formation of a compound from simpler compounds or elements.

2Na(s) + Cl2(g) 2 NaCl(s)

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Decomposition Reactions

Separation into constituents by chemical reaction.

catalysis2 H2O2 (aq) 2H2O(l) + O2(g)

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Precipitation Reactions

The process of separating a substance from a solution as a solid.

AgNO3 + NaCl ---> AgCl + NaNO3

precipitate

Total Ionic EquationAg+1

(aq) + NO3(aq)-1 + Na(aq)

+1 + Cl(aq)-1 AgCl(s) + NO3(aq)

-1 + Na(aq)+1

Net Ionic Equation

Ag+1(aq) + Cl(aq)

-1 AgCl(s)

Dr. S. M. Condren

Neutralization Reactions

acid + base ---> “salt” + water


HCl + NaOH ---> NaCl + H2O


H+1 + Cl-1 + Na+1 + OH-1 Na+ + Cl-1 + H2O

Net Ionic Equation

H+1 + OH-1 H2O

Dr. S. M. Condren


acid + base ---> “salt” + water


H2SO4 + 2KOH ---> K2SO4 + 2H2O


2H+ + SO4-2 + 2K+ + 2OH-1 2H2O + SO4

-2 + 2K+

Net Ionic Equation

2H+1 + 2OH-1 2H2O

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• acid

• base

• salt

Household acids and Bases

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• acid– Any of a large class of sour-tasting

substances whose aqueous solutions are capable of turning blue litmus indicators red, of reacting with and dissolving certain metals to form salts, and of reacting with bases or alkalis to form salts.

– Substance that donates H+ ions to solution

Dr. S. M. Condren

An acid -------> HAn acid -------> H++ in water in waterAn acid -------> HAn acid -------> H++ in water in water

Some Some strongstrong acids areacids are

HClHCl hydrochlorichydrochloric

HH22SOSO44 sulfuricsulfuric

HClOHClO44 perchloricperchloric

HNOHNO33 nitricnitric HNOHNO33

ACIDS

Dr. S. M. Condren

WEAK ACIDS = weak WEAK ACIDS = weak electrolyteselectrolytes

CHCH33COCO22HH acetic acidacetic acid

HH22COCO33 carbonic acidcarbonic acid

HH33POPO44 phosphoric acidphosphoric acid

HFHF hydrofluoric acidhydrofluoric acid

Acetic acid

Weak Acids

Dr. S. M. Condren

Nonmetal oxides can be acidsNonmetal oxides can be acids

COCO22(aq) + H(aq) + H22O(liq) O(liq)

---> H---> H22COCO33(aq)(aq)

SOSO33(aq) + H(aq) + H22O(liq) O(liq)

---> H---> H22SOSO44(aq)(aq)

and can come from burning coal and can come from burning coal and oil.and oil.

ACIDS

Dr. S. M. Condren


• base– Any of a large class of compounds, including

the hydroxides and oxides of metals, having a bitter taste, a slippery solution, the ability to turn litmus blue, and the ability to react with acids to form salts.

– Substance that donates a OH-1 ion to solution

Dr. S. M. Condren

Base ---> OHBase ---> OH-- in water in waterBase ---> OHBase ---> OH-- in water in water

NaOHNaOH(aq)(aq) Na Na++(aq)(aq) + OH + OH--

(aq)(aq)

NaOH is a strong base

BASES

Dr. S. M. Condren

Ammonia, NH3

An Important Base

Dr. S. M. Condren


• salt– The term salt is also applied to substances

produced by the reaction of an acid with a base, known as a neutralization reaction.

– Salts are characterized by ionic bonds, relatively high melting points, electrical conductivity when melted or when in solution, and a crystalline structure when in the solid state.

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pH Scale

@ -2 -----> @ +16

pH = - log [H3O+]

[H3O+] = -antilog pH = 10-pH

For log problems, only decimal places are significant, and all decimal places count

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Example: What is the pH of a 1.0 M solution of HCl?pH = - log [H3O+]

[H3O+] = MHCl = 1.0 M

• because it is a strong electrolyte

pH = -log(1.0) = 0.002 significant figures

Dr. S. M. Condren

Example: What is the pH of a 1.0 M solution of HCl?pH = - log [H3O+]

[H3O+] = MHCl = 1.0 M


pH = -log(1.0) = 0.002 SF

Dr. S. M. Condren

Example: What is the pH of a 0.010 M solution of HCl?

pH = - log [H3O+]

[H3O+] = MHCl = 0.010 M


pH = -log(0.010) = 2.002 SF

Dr. S. M. Condren

Example: What is the pH of a 2.0 M solution of HCl?

pH = - log [H3O+]

[H3O+] = MHCl = 2.0 M


pH = -log(2.0) = -0.302 SF

Dr. S. M. Condren

Example: What is the [H3O+] of a solution that has a pH = 2.30? [H3O+] = -antilog pH = 10-pH

[H3O+] = 10-pH = 10-2.30 = 5.0x10-3M2 SF 2 SF

decimal point pacement

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pH Scale

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Indicators

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Oxidation-Reduction Reaction

Oxidation - loss of electrons

Reduction - gain of electrons

Redox reaction

oxidizing agent - substance that causes oxidation

reducing agent - substance that cause reduction

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Identify the oxidizing agent in the reaction:

2Al(s) + 6 H+ ==> 2 Al3+

(aq) + 3 H2(g)

Al, H+, Al3+, H2

Identify the oxidizing agent in the reaction:

Dr. S. M. Condren

Redox Reactions

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

4 Al(s) + 3 O2(g) 2 Al2O3(s)

2HgO2HgO(s)(s) 2 Hg 2 Hg(l)(l) + O + O2(g)2(g)

2 Al2 Al(s)(s) + 3 Br + 3 Br2 (l)2 (l) Al Al22BrBr6 (s)6 (s)

Fe2O3(s) + 2 Al(s) 2 Fe(l) + Al2O3(s)

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Solution

Solutions, in chemistry, homogeneous mixtures of two or more substances.

Dr. S. M. Condren

Solute

The substance that is present in smallest quantity is said to be dissolved and is called the solute. The solute can be either a gas, a liquid, or a solid.

Dr. S. M. Condren

Solvent

The substance present in largest quantity usually is called the solvent. The solvent can be either a liquid or a solid.

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Preparing a Solution

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Preparing a Solution by Dilution

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Molarity

The number of moles of solute per liter of solution.

molarity => M moles of soluteM = -------------------- liter of solution

units => molar = moles/liter = M

Dr. S. M. Condren

Example: A sample of NaNO3 weighing 0.38g is placed in a 50.0-mL volumetric flask. The flask is then filled with water to the mark on the neck, dissolving the solid. What is the molarity of the resulting solution?

M =(0.38g NaNO3)(50.0-mL soln)

= 0.089 mol/L = 0.089 molar = 0.089 M

Dr. S. M. Condren

Stoichiometric Roadmap

Gramsof A

Molesof A

Molesof B

Gramsof B

mol A x (mol B/mol A)Multiply by the

stoichiometric factor

mol B x (g B/mol B)Multiply by the

molar mass

Volumesolution B

g B x (mol A/g A)Divide by themolar mass

Volumesolution A

Vol A x (mol A/L A)Multiply by molarity

Mol B x (L B/mol B)Diviide by molarity

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Example: What mass of Na2CO3, in grams, is required for complete reaction with 50.0mL of 0.125 M HNO3?

Na2CO3(aq) + 2 HNO3(aq) 2 NaNO`3(aq) + CO2(g) + H2O(l)

#g Na2CO3 = (50.0 mL)

= 0.331 g Na2CO3

Dr. S. M. Condren

EXAMPLE: Lye, which is sodium hydroxide, can be neutralized by sulfuric acid. How many milliliters of 0.200 M H2SO4 are needed to react completely with 25.0 mL of 0.400 M NaOH?

2 NaOH(aq) + H2SO4(aq) -----> Na2SO4(aq) + 2 H2O

(25.0 mL NaOH) #mL H2SO4 = ----------------------

(0.400 mol NaOH) ------------------------- (1 L NaOH)

(1 L)--------------(1000 mL)

(1 mol H2SO4) -------------------(2 mol NaOH)

(1000 mL H2SO4) ------------------------(0.200 mol H2SO4)

= 25.0 mL H2SO4

Chapter 5

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