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CHAPTER 4: A QUEOUS REACTIONS AND SOLUTION STOICHIOMETRY
GENERAL PROPERTIES OF AQUEOUS SOLUTIONS
ö A solution is a homogeneous mixture of two or more substances. ö A solution in which water is the dissolving medium is called an aqueous solution.
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SOLUTE VS . S OLVENT
ö The substance present in the greatest quantity is called the solvent. ö Water is considered the universal solvent because of its ability to dissolve many substances. ö The other dissolved substances are called the solutes . ö A solvent dissolves a solute.
ELECTROLYTES
ö A substance whose aqueous solutions contain ions is called an electrolyte because it will allow electric current to flow through it. Example: NaCl ö A substance that does not form ions in solution is
called a nonelectrolyte. Example: C12 H22 O11
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IONIC COMPOUNDS IN WATER ö When ionic compounds dissolve in water the ions separate and are surrounded by water molecules. ö Once the ions have separated, they can move and conduct electricity. In the crystalline form they cannot move or conduct electricity. ö The solvation process helps stabilize the ions in solution and prevents the cations and anions from recombining.
IONIC COMPOUND IN WATER
ö Water molecules are polar and have a δ+ and δ- side. ö The more electronegative oxygen atoms are attracted to the positive ions. ö The less electronegative hydrogen atoms are attracted to the negative ions.
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Ionic Compounds : undergo dissociation: process by which many ionic substances dissolve in water, the solvent pulls the individual ions from the crystal and solvates them. _
{NaCl + H2O} Polar water molecule
+ +
H2O
Dissociation H O NaCl 2 + - (s) Na (aq) + Cl (aq)
Hydration of sodium chloride
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MOLECULAR COMPOUNDS IN WATER
ö When a molecular compound dissolves in water, the solution usually consists of intact molecules (neutral) dispersed throughout the system. ö Consequently, most molecular compounds are nonelectrolytes. A few exceptions include strong acids like HCl.
Ionic (electrolyte) Molecular (non electrolyte)
Molecular (covalent) Compounds: mostly insoluble gases , except polar organic liquids containing O & N (polar : acids, bases, alcohols, etc.)
Insoluble gases : NO 2, CH 4, CO 2, O 2, P 2O5, N 2, CO, etc.
Polar Covalent {carbon (C) chains containing H,O or N}: CH 3OH, C 6H12 O6, C6H5OH, etc.
H2O C6H12 O6(s) C6H12 O6 (aq)
H2O C3H5OH (l) C3H5OH (aq)
Dissolve without dissociating into ions
{Ethanol+Water }
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STRONG AND WEAK ELECTROLYTES
ö Strong electrolytes are those solutes that exist in solution completely or nearly completely as ions. ö Weak electrolytes are those solutes that exist in solution mostly in the form of molecules with only a small fraction in the form of ions.
ELECTROLYTES : S TRONG AND WEAK
ö A strong electrolyte dissociates completely when dissolved in water.
H2O + - HCl (g) H (aq) + Cl (aq)
ö A weak electrolyte only dissociates partially when dissolved in water.
H2O + - NH 4OH (aq) NH 4 (aq) + OH (aq)
Acetic Acid,
H2O + - HC 2H3O2 (aq) H (aq) + C 2H3O2 (aq)
molecules ions
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ELECTROLYTES : A VISUAL
CHEMICAL EQUILIBRIUM
ö Chemical equilibrium is the state of balance in which the relative numbers of each type of ion or molecule in the reaction are constant over time. ö Equilibrium is represented by a double arrow for the yield sign. This shows that the reaction is occurring in both directions.
ö Dynamic Equilibrium
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STRONG ELECTROLYTES
ö Remember that soluble ionic compounds are strong electrolytes. ö We consider ionic compounds as those composed of
metals and nonmetals – such as NaCl, FeSO 4 and Al(NO 3)3 – or compounds containing the ammonium + ions, NH 4 - such as NH 4Br and (NH 4)2CO 3.
SAMPLE EXERCISE 4.1
ö The diagram below represents an aqueous solution of
one of the following compounds: MgCl 2, KCl, or K2SO 4. Which solution does the drawing best represent?
The diagram shows twice as many cations as anions, consistent with the formulation
K2SO 4.
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PRACTICE EXERCISE
ö If you were to draw diagrams like the previous one representing aqueous solutions of each of the following ionic compounds, how many anions would you show if the diagram contained six cations?
6 a. NiSO 4
12 b. Ca(NO 3)2
c. Na 3PO 4 2
d. Al 2(SO 4)3 9
PRECIPITATION REACTIONS
ö Reactions that result in the formation of an insoluble product are called precipitate reactions. ö A precipitate is an insoluble solid formed by a reaction in solution.
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PRECIPITATION REACTIONS
Aqueous solutions, reacting to produce a
precipitate (an insoluble compound). Example: KI (aq) + Pb(NO 3)2 (aq) Predict the solubility of compounds in reaction:
Pb(NO 3)2(aq) + 2KI (aq) ‰ 2 KNO 3(aq) PbI 2(s) Precipitate (ppt)
Pb(NO 3)2
KI
PbI 2
SOLUBILITY
ö The solubility of a substance at a given temperature is the amount of substance that can be dissolved in a given quantity of solvent at a given temperature. ö Most important solubility rule: Alkali metal ions and ammonium ions are soluble in water.
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SOLUBILITY – ABILITY TO BE DISSOLVED
ö Soluble (aq) ó a substance which easily dissolves in water ö Slightly Soluble (s) ó a substance that only dissolves a tiny bit in water ö Insoluble (s) ó substance that does NOT dissolve in water
ó substance remains separate from the H2O molecules (no interaction) ó substances are insoluble when the ions attract so strongly that they CANNOT be
pulled apart by H2O molecules
SOLUBILITY : A VISUAL
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SOLUBILITY RULES
ö To know if an ionic substance is soluble or not, you must consult the solubility rules ó lists which ions are typically soluble and insoluble ó also lists exceptions to the rules
Note : Acids are not explicitly listed in the rules because all acids are aqueous solutions (thus soluble)
SOLUBILITY RULES: for Ionic Compounds (Salts) 1. All salts of alkali metals (IA) are soluble . + 2. All NH 4 salts are soluble . - - - 3. All salts containing the anions: NO 3 , ClO 3 , ClO 4 , - (C 2H3O2 ) are soluble . - - - 2+ + 4. All Cl , Br , and I are soluble except for Hg 2 , Ag , and Pb 2+ salts. 2- 2+ 2+ 2+ 5. All SO 4 are soluble except for Pb , Ba , and Sr .
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Solubility Rules for Common Ionic Compounds in Water
Soluble Ionic Compounds Exceptions: these are insoluble w/ respective anions - Nitrates, NO 3 none - Acetates, C 2H3O2 none - + 2+ 2+ Chlorides, Cl Ag , Hg 2 (2 Hg atoms), Pb - + 2+ 2+ Bromides, Br Ag , Hg 2 , Pb - + 2+ 2+ Iodides, I Ag , Hg 2 , Pb 2- + 2+ 2+ 2+ 2+ Sulfates, SO 4 Ag , Hg 2 , Pb , Sr , Ba Alkali metal cations none + Ammonium, NH 4 none
Insoluble Ionic Compounds Exceptions: these are soluble w/ respective anions 2- + 2+ 2+ 2+ Sulfides, S alkali metal cations, NH 4 , Ca , Sr , Ba 2- + Carbonates, CO 3 alkali metal cations, NH 4 3- + Phosphates, PO 4 alkali metal cations, NH 4 - + 2+ 2+ 2+ Hydroxides, OH alkali metal cations, NH 4 , Ca , Sr , Ba
SAMPLE EXERCISE 4.2
ö Classify the following ionic compounds as soluble or insoluble in water:
a. Sodium carbonate: Na 2CO 3,
Soluble
b. Lead (II) sulfate : PbSO 4
Insoluble
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PRACTICE EXERCISE
ö Classify the following compounds as soluble or insoluble in water:
a. Cobalt (II) hydroxide: Insoluble
b. Barium nitrate: Soluble
c. Ammonium phosphate: Soluble
EXCHANGE REACTIONS
ö Reactions in which positive ions and negative ions appear to exchange partners conform to the following generic equation: AX + BY ‰ AY + BX Such reactions are called exchange reactions, double displacement reactions, or metathesis reactions.
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SOLUBILITY RULES -PRACTICE
ö Are the following soluble or insoluble?
ó Be(C 2H3O2)2 soluble
ó MgS insoluble
ó BaSO 4 insoluble
ó K3PO 4 soluble
STEPS FOR WRITING DOUBLE DISPLACEMENT REACTIONS
1. Write the formulas for the reactants (balance charges). 2. Write the formulas for the products by switching the cations of the two reactants (balance charges). 3. Balance the equation. 4. Include states of matter when needed.
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SAMPLE EXERCISE 4.3
a) Predict the identity of the precipitate that forms
when solutions of BaCl 2 and K2SO 4 are mixed.
BaSO 4 is insoluble and will precipitate from solution. KCl is soluble.
b) Write the balanced chemical equation for the reaction.
PRACTICE EXERCISE
a) What compound precipitates when solutions of
Fe 2(SO 4)3 and LiOH are mixed?
Fe(OH) 3 b) Write a balanced equation for the reaction.
Fe 2(SO 4)3(aq) + 6 LiOH(aq) 2Fe(OH) 3(s) +3Li 2SO 4(aq)
c) Will a precipitate form when solutions of Ba(NO 3)2 and KOH are mixed?
no (both possible products, Ba(OH) 2 and KNO 3, are water soluble)
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COMPLETE IONIC EQUATIONS
o A molecular equation is an equation showing the complete chemical formulas of the reactants and products without showing dissociated ions. o Ex:
Pb(NO 3)2(aq) + 2KI (aq) ‰ PbI 2(s) + 2KNO 3(aq)
o A complete ionic equations is an equation that shows all soluble strong electrolytes as ions. o Ex: 2+ - + - Pb (aq) + 2NO 3 (aq) + 2K (aq) + 2I (aq) ‰ PbI 2(s) + + - 2K (aq) + 2NO 3 (aq)
NET IONIC EQUATIONS
ö Ions that appear in identical forms on both sides of the yield sign are called spectator ions because they do not actually react. ö A net ionic equation is an equation in which the spectator ions have been removed. ö Complete Ionic Equation 2+ - + - + Pb (aq) + 2NO 3 (aq) + 2K (aq) + 2I (aq) ‰ PbI 2(s) + 2K (aq) + - 2NO 3 (aq)
ö Net Ionic Equation 2+ - Pb (aq) + 2I (aq) ‰ PbI 2(s)
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NET IONIC EQUATIONS
ö In a net ionic equation, the numbers of atoms and the charges are conserved in the reaction. ö If every ion in a complete ionic equation is a spectator, then no reaction occurs.
STEPS FOR WRITING NET IONIC EQUATIONS
1. Write the balanced equation. 2. Write the complete ionic equation which shows strong electrolytes (ionic compounds and strong acids) written as ions. 3. Identify and cancel spectator ions.
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WAYS OF EXPRESSING PRECIPITATION REACTIONS There are three different: (1) Molecular Equations
→ AgNO 3 (aq ) + KCl (aq ) AgCl (s) + KNO 3 (aq ) (2) Ionic Equations
+ - + - → Ag (aq ) + NO 3 (aq ) + K (aq ) + Cl (aq ) AgCl (s) + K+ (aq ) + NO - (aq ) (3) Net Ionic Equations 3
- Ag +(aq) + Cl (aq ) → AgCl (s)
SAMPLE EXERCISE 4.4 Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed. Molecular eq.
CaCl 2(aq) + Na 2CO 3(aq) CaCO 3(s) + 2NaCl (aq) Ionic Eq.
2+ - + 2- Ca (aq) + 2 Cl (aq) + 2Na (aq) + CO 3 (aq)
+ - CaCO 3(s) + 2Na (aq) + 2Cl (aq) Net Ionic eq.
2+ 2- Ca (aq) + CO 3 (aq) CaCO 3(s)
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PRACTICE EXERCISE
Write the net ionic equation for the precipitation reaction that occurs when aqueous solutions of silver (I) nitrate and potassium phosphate are mixed.
ACID -BASE REACTIONS
ö Acids and bases are also common electrolytes. ö Acids are substances that ionize in aqueous solutions to form hydrogen ions. ö Since hydrogen ions consist of only 1 proton, acids are often called proton donors .
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Bases ACIDS
öSubstances that öSubstances that produce H+ ions produce OH − ions when dissolved in when dissolved in water water (Arrhenius). (Arrhenius). ö Ex) ö Ex) + – + – HNO 3 ‰ H + NO 3 NaOH ‰ Na + OH ö Proton donor (L&B) ö HCl (g) ö Proton Acceptor (L&B)
ö NH 3 (g)
STRONG VS . W EAK
ö Both strong & weak acids and bases exist ó Strong acids & bases dissociate completely in solution
öExist as 100% ions
öStrong electrolytes & conductors ó Weak acids & bases only partially dissociate in solution
öFew ions exits
öWeak electrolytes & conductors
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STRONG VS . W EAK ELECTROLYTES Strong Acid: Weak Acid:
HCl HNO 2
BASES
The strong bases are : • Alkali metals (IA) Hydroxides • Barium Hydroxide • Strontium Hydroxide • (weaker: Ammonium , Calcium Hydroxides )
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Acids
ö Monoprotic acids contain one ionizable hydrogen. Example: HCl ö Diprotic acids contain two ionizable hydrogens.
Example: H 2SO 4 ö Triprotic acids contain three ionizable hydrogens.
Example: H 3PO 4 ö Most multiprotic acids only lose one hydrogen to a reasonable extent.
Citric acid
Acids
ö Acids taste sour. ö Some acids also contain hydrogens that will not ionize. ö Normally the hydrogen(s) that will ionize are listed at the from of the chemical formula.
ö Example: HC 2H3O2 ö Sometimes the hydrogen that will ionize is listed at the end of a COOH group.
ö Example: CH 3COOH
Diprotic acid Triprotic acid
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Bases
ö Bases taste bitter. ö Bases are substances that accept (react with ) hydrogen ions. ö Bases produce hydroxide ions (OH -) when they dissolve in water. ö Bases do not have to contain the hydroxide ion.
Examples: NaOH vs. NH 3
NaOH(aq) Na + (aq) + OH - (aq)
Strong vs. Weak
ö Strong acids and bases are strong electrolytes so they fully ionize in solution. ö Weak acids and bases are weak electrolytes so they only partially ionize in solution.
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The Strong Acids
ö You need to memorize the strong acids.
HCl (aq) hydrochloric acid
HBr (aq) hydrobromic acid
HI (aq) hydroiodic acid
HClO 3 (aq) chloric acid
HClO 4 (aq) perchloric acid
HNO 3 (aq) nitric acid
H2SO 4 (aq) sulfuric acid
STRONG BASES
ö You need to memorize the strong bases.
Group 1 metal hydroxides: LiOH, NaOH, KOH, RbOH, and CsOH
Heavy Group 2 metal hydroxides: Ca(OH) 2, Sr(OH) 2, and Ba(OH) 2
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SAMPLE EXERCISE 4.5
ö The following diagrams represent aqueous solutions of three acids (HX, HY, and HZ) with water molecules omitted for clarity. Rank them from strongest to weakest.
PRACTICE EXERCISE
ö Imagine a diagram showing 10 Na + ions and 10 OH - ions. If this solution were mixed with the one pictured on the previous question for HY, what would the diagram look like that represents the solution after any possible reaction? (H + ions will - react with OH ions to form H2O.)
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STRONG VS . W EAK ELECTROLYTE
Strong Weak Non-electrolyte Electrolyte Electrolyte
Ionic All None None
Molecular Strong Acids Weak acids and All other weak bases compounds
SAMPLE EXERCISE 4.6
ö Classify the following substances as a strong
electrolyte, weak electrolyte, or nonelectrolyte: CaCl 2, HNO 3,C2H5OH (ethanol), HCOOH (formic acid), KOH.
ö Strong: CaCl 2, KOH, HNO 3 ö Weak: HCOOH
ö Non: C2H5OH
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PRACTICE EXERCISE
ö Consider solutions in which 0.1 mol of each of the following compounds is dissolved in 1L of water:
Ca(NO 3)2,C6H12 O6 (glucose), CH 3COONa (sodium acetate), and CH 3COOH (acetic acid). Rank the solutions in order of increasing electrical conductivity based on the fact that the greater the number of ions in solution the greater the conductivity.
NEUTRALIZATION REACTIONS
ö When a solution of an acid and a solution of a base are mixed, a neutralization reaction occurs. ö The products of the reaction are water and a salt. ö A salt is any ionic compound whose cation comes from a base and whose anion comes from an acid.
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NEUTRALIZATION REACTIONS (A RRHENIUS ). Generally, when solutions of an acid and a base are combined, the products are a salt and water.
→ HCl (aq ) + NaOH (aq ) NaCl (aq ) + H 2O ( l)
H+ (aq ) + Cl - (aq ) + Na + (aq ) + OH -(aq ) → + - Na (aq ) + Cl (aq ) + H 2O ( l)
+ → H (aq ) + OH- (aq ) H2O (l)
NEUTRALIZATION REACTIONS
Observe the reaction between a weak base , Milk of Magnesia, Mg(OH) 2 (s) and a strong acid HCl (aq).
How would you write the net ionic equation for such a reaction?
{Movie }
Mg(OH) 2(s) + 2HCl (aq) ‰ MgCl 2(aq) + 2H2O (l)
+ - 2+ - Mg(OH) 2(s) + 2H (aq) + 2Cl (aq) ‰ Mg (aq) + 2Cl (aq) + 2H2O (l)
+ 2+ Mg(OH) 2(s) + 2H (aq) ‰ Mg (aq) + 2H2O (l)
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SAMPLE EXERCISE 4.7
a. Write a balanced equation for the reaction between aqueous solutions of acetic acid
(CH 3COOH) and barium hydroxide. b. Write the net ionic equation for this reaction.
PRACTICE EXERCISE
a. Write a balanced equation for the reaction of
carbonic acid (H 2CO 3) and potassium hydroxide.
b. Write the net ionic equation for this reaction.
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GAS FORMATION
ö The sulfide ion, bicarbonate and the carbonate ion will react with acids to form gases.
ö Example: S 2-
2HCl (aq) + Na 2S(aq) ‰ H2S(g) + 2NaCl (aq)
2- ö Example: HCO 3
HCl (aq) + NaHCO 3(aq) ‰ NaCl (aq) + H 2CO 3(aq)
H2CO 3(aq) ‰ H2O(l) + CO 2(g) So …
HCl (aq) + NaHCO 3(aq) ‰ NaCl (aq) + H 2O(l) + CO 2(g)
OXIDATION -REDUCTION REACTIONS
ö Oxidation-reduction reactions (redox reactions) are reactions in which electrons are transferred between reactants. ö Oxidation is loss of electrons (Increase in ON). ö Reduction is gain of electrons (Decrese in ON). ö Oxidation and reduction always occur together.
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RULES FOR OXIDATION NUMBERS
1. For an atom in its elemental form, the oxidation number is always zero.
Ex: Na = 0 S8 = 0 N2 = 0
F2 = 0 P4 = 0 O2 = 0 2. For any monatomic ion the oxidation number equals the charge on the ion. Ex: Mg 2+ = +2, N3- = -3 Al 3+ = +3 3. The oxidation number of hydrogen is usually +1 when bonded to nonmetals and -1 when bonded to metals Ex: HCl H = +1 LiH H = -1
4. The oxidation number of oxygen is usually -2 (except 2- in peroxide, O2 , where it is -1).
Example: CO 2 O = -2
H2O2 O = -1 5. The oxidation number of fluorine is -1 in all compounds. The other halogens have an oxidation number of -1 except when combined with oxygen in oxyanions. Example : NaF F = -1 HBr Br = -1 ClO - Cl = +1
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6. The sum of the oxidation numbers of all atoms in a neutral compound is zero. Example:
P2O5 = 0 O=-2,soP=+5 7. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion. Example: 2- CO 3 = 0 O=-2,soC=+4
NH 4 = +1 H=+1,soN=-3 8. For compounds contain elements of: G1A=+1,G2A=+2, Al=+3
SAMPLE EXERCISE 4.8
ö Determine the oxidation number of sulfur in each of the following:
a. H 2S 2(+1) + S = 0 S = -2
b. S 8 S = 0
c. SCl 2 S + 2(-1) = 0 S= +2
d. Na 2SO 3 Na 2(+1) + S + 3(-2) = 0 S = +4 2- e. SO 4 S+4(-2) = -2 S =+6
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PRACTICE EXERCISE
ö What is the oxidation state of each element in the following:
a. P 2O5
b. NaH
2- c. Cr 2O7
d. SnBr 4
e. BaO 2
OXIDATION OF METALS
ö Many metals undergo displacement reactions with acids to form salts and hydrogen gas.
ö Write the oxidation numbers for each element to determine which element is oxidized and which is reduced. ö Since Mg goes from 0 to +2, it is oxidized. Since H goes from +1 to 0, it is reduced.
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SAMPLE EXERCISE 4.9
ö Write the balance equation and net ionic equation for the reaction of aluminum with hydrobromic acid.
o +1 +3 0
PRACTICE EXERCISE
a. Write a balanced equation and net ionic equation for the reaction between magnesium and cobalt (II) sulfate.
b. What is oxidized and what is reduced in the reaction?
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THE ACTIVITY SERIES
ö A list of metals arranged in order of decreasing ease of oxidation is called an activity series. ö Active metals are at the top of the series and are most easily oxidized. ö Noble metals are at the bottom of the series are not very reactive. ö Any metal on the list can be oxidized by the ions of elements below it.
ACTIVITY SERIES
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USING THE ACTIVITY SERIES EXAMPLES
1) magnesium sulfate+ zinc metal
2) iron metal + copper (II) nitrate ‰ NR
Fe (s) + Cu(NO 3)2 (aq) ‰ Cu (s) + Fe(NO 3)2
We can also write CIE & NIE!!! (only aqueous dissociate)
2+ – CIE: Fe (s) + Cu (aq) + 2 NO 3 (aq) 2+ – ‰ Cu (s) + Fe (aq) + 2 NO 3 (aq)
NIE: Fe (s) + Cu 2+ (aq) ‰ Cu (s) + Fe 2+ (aq)
METALS REACTING W/A CIDS
ö Only metals listed ABOVE hydrogen in the series are able to react with acids
ö If the reaction does occur, H2 gas is produced. ö Example #1:
Ni (s) + 2 HCl (aq) ‰ NiCl 2(aq) + H 2 (g)
‹ Ni is above H 2, so the rxn occurs! ö Example #2: Cu (s) + 2HCl (aq) ‰ NR
‹ Cu is below H 2, so no rxn occurs!
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ACTIVITY SERIES OF HALOGENS :
ACTIVITY SERIES ö The idea is the same here… OF HALOGENS : ó Fluorine is the most reactive ó Iodine is the least reactive Fluorine ö Only a MORE reactive halogen can replace a Chlorine LESS reactive one Bromine ö In these single replacement rxns, the single Iodine element (halogen gas) must be listed ABOVE the halogen anion in the 2nd reactant for a rxn to occur
Cl 2 + 2NaBr (aq) ‰ 2NaCl (aq) + Br 2 (g)
I2 + 2KCl ‰ NR
SAMPLE EXERCISE 4.10
ö Will an aqueous solution of iron (II) chloride oxidize magnesium metal? If so, write the balanced equation and net ionic equation for the reaction.
Mg (s) + FeCl 2 (aq) MgCl 2(aq) + Fe (s)
Mg(s) + Fe 2+ Mg 2+ (aq) + Fe(s)
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PRACTICE EXERCISE
ö Which of the following metals will be oxidized by
Pb(NO 3)2: Zn, Cu, Fe?
CONCENTRATIONS OF SOLUTIONS
ö Scientists use the term concentration to designate the amount of solute dissolved in a given quantity of solvent or solution. ö Molarity (M) expresses the concentration of a solution as the number of moles of solute in a liter of solution.
M = mol/Vl
Units: M, mol/L or mol L -
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SAMPLE EXERCISE 4.11
ö Calculate the molarity of a solution made by dissolving 23.4g of sodium sulfate in enough water to form 125mL of solution.
PRACTICE EXERCISE
ö Calculate the molarity of a solution made by dissolving 5.00g of glucose in sufficient water to form exactly 100mL of solution. Answer: 0.278 M
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IONIC COMPOUNDS
ö When calculating concentration, remember that ionic compounds break into ions in solutions.
Example: MgCl 2 breaks into Mg 2+ and 2Cl -
ö NaCl(aq) Na + (aq) + Cl - (aq) 0.1M 0.1M 0.1M
+ 2- Na 2SO 4 (aq) 2Na (aq) + SO 4 (aq) 1 M 2 M 1 M
SAMPLE EXERCISE 4.12
ö What are the molar concentrations of each of the ions present in a 0.025M aqueous solution of calcium nitrate?
2+ - ö Ca (NO 3)2 (aq) Ca (aq) + 2NO 3 (aq) 0.025 M 0.025 M 2 x 0.025M = 0.05 M
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PRACTICE EXERCISE
ö What is the molar concentration of K + ions in a 0.015M solution of potassium carbonate? Answer : 0.030 M
MOLARITY
ö Since molarity equals mol/L, it can be used as a conversion factor. Example: 0.3 M NaOH ie: 0.3 mol NaOH 1 L soln 0.3mol NaOH 1000 mL soln
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SAMPLE EXERCISE 4.13
ö How many grams of Na 2SO 4 are required to make 0.350L of 0.500M Na 2SO 4?
PRACTICE EXERCISE
a) How many grams of Na 2SO 4 are there in 15mL of 0.50M Na 2SO 4? Answers: 1.1 g
b) How many milliliters of 0.50M Na 2SO 4 solution are needed to provide 0.038 mol of this salt? Answer: 76 mL
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DILUTION
ö Stock solutions are solutions that are purchased at a specific concentration, but sometimes a different concentration is needed for an experiment. ö Solutions of lower concentration can then be obtained by adding water, a process called dilution.
ö When solvent is added to dilute a solution, the number of moles of solute remains the same.
M1 x V 1 = M2 x V 2 (before dilution) (after dilution)
SAMPLE EXERCISE 4.14
ö How many milliliters of 3.0M H2SO 4 are needed to make 450mL of 0.10M H2SO 4?
M1 x V 1 = M2 x V 2
3 x V 1 = 450 x 0.10
V1= 15 mL
Note: Acid added to water for dilution
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PRACTICE EXERCISE a) What volume of 2.50M lead (II) nitrate solution contains 0.0500mol of Pb 2+? Answer: 0.0200 L= 20 ml
b) How many milliliters of 5.0M K2Cr 2O7 solution must be diluted to prepare 250mL of 0.10M solution? Answer: 5.0 mL
c) If 10.0mL of a 10.0M stock solution of NaOH is diluted to 250mL, what is the concentration of the resulting stock solution? Answer: 0.40 mL
EXAMPLE
ö If 50.0 mL of 0.10 M NaOH aq is added to 30 mL of 0.25 M NaOH aq, Calculate the molarity of the final solution.
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PRACTICE EXERCISE
ö If 75.0 mL of 0.15 M NaCl (aq) is added to 50 mL of 0.30 M CaCl 2(aq), Calculate the molarity of the final solution.
SOLUTION STOICHIOMETRY AND CHEMICAL ANALYSIS
ö You can use molar mass and molarity to solve stoichiometry problems.
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SAMPLE EXERCISE 4.15
ö How many grams of Ca(OH) 2 are needed to neutralize 25.0mL of 0.100M HNO 3?
2NHO 3 (aq) + Ca(OH) 2(aq) Ca(NO 3)2 (aq) + H2O(l) 0.1 M, 25.0 ml g ??
PRACTICE EXERCISE a) How many grams of NaOH are needed to neutralize
20.0mL of 0.150M H2SO 4 solution? Answer:0.240 g
b) How many liters of 0.500M HCl are needed to react
completely with 0.100 mol of Pb(NO 3)2, forming a precipitate of PbCl 2? Answer: 0.400L
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TITRATIONS
ö To determine the concentration of a particular solute in a solution, chemists often carry out a titration, which involves combining a sample of the solution with a reagent solution of known concentration, called a standard solution.
öBalanced chemical equation is needed to calculate the unknown con.
TITRATIONS
ö The point at which stoichiometrically equivalent quantities are brought together is known as the equivalence point . ö In acid-base titrations, dyes known as acid-base indicators are used to identify the equivalence point. ö The color change of the indicator signals the end point of the titration, which usually coincides very nearly with the equivalence point.
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• Suppose we know the molarity of a NaOH solution and we want to find the molarity of an HCl solution. • We know: – molarity of NaOH, volume of HCl. • What do we want? – Molarityof HCl. • Whatdowe do? – Take a known volume of the HCl solution, measure the mL of NaOH required to react completely with the HCl.
TITRATION : EXAMPLE : N EUTRALIZATION REACTION
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IMPRTANT NOTES :
Or
SAMPLE EXERCISE 4.16 a) How many grams of chloride ion are in a sample of the water if 20.2mL of 0.100M Ag + is needed to react with all the chloride in the sample? + - Ag (aq) + Cl (aq) ‰ AgCl (s) 20.0 mL, 0.1 M ? g
Solution
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ö If the sample has a mass of 10.0g, what percent Cl - does it contain?
PRACTICE EXERCISE a) A solution that contains Fe 2+ is titrated with 47.20mL - - of 0.02240M MnO 4 . How many moles of MnO 4 were added to the solution? - 2+ + 2+ 3+ MnO 4 (aq) + 5Fe (aq) + 8H (aq) ‰ Mn (aq) + 5Fe (aq) + 4H2O(l) Answers: × –3 – 1.057 10 mol MnO 4
b) How many moles of Fe 2+ were in the sample? Answer: 5.286 × 10 –3 mol Fe 2+
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PRACTICE EXERCISE
c) How many grams of iron were in the sample? Answer: 0.2952 g
d) If the sample had a mass of 0.8890g, what is the percentage of iron in the sample? Answer: 33.21%
SAMPLE EXERCISE 4.17
The concentration of NaOH is normally in the range of 3 to 6M. The NaOH is analyzed periodically. In one such
analysis, 45.7mL of 0.500M H2SO 4 is required to neutralize a 20.0mL sample of NaOH solution. What is the concentration of the NaOH solution?
45.7mL of 0.500M 20.0mL, M?? Note: The eq. must be balanced Then apply the neutralization law:
acid base
M2 = 2.28 M
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PRACTICE EXERCISE
ö What is the molarity of an NaOH solution if 48.0mL is
needed to neutralize 35.0mL of 0.144M H2SO 4? ö Answer: 0.210 M
SAMPLE INTEGRATIVE EXERCISE
ö A sample of 70.5mg of potassium phosphate is added to 15.0mL of 0.050M silver nitrate, resulting in the formation of a precipitate.
a. Write the molecular equation for the reaction.
K3PO 4(aq) + 3AgNO 3(aq) Ag 3PO 4(s) 3KNO 3(aq)
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b. What is the limiting reagent in the reaction?
K3PO 4(aq) + 3AgNO 3(aq) Ag 3PO 4(s) 3KNO 3(aq) 70.5 mg 15.0mL, 0.050M
moles K 3PO 4 :
Then /1 = 3.32x10 -4
moles AgNO 3 : =
Then /3 = 2.5x10 -4 (LR)
SAMPLE INTEGRATIVE EXERCISE
c) Calculate the theoretical yield, in grams, of the precipitate that forms.
3AgNO 3(aq) Ag 3PO 4(s)
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أ أ ءك أو : ا ر زو إ ل و ا م ا ول و و ن و ا ر ا و ةً ز ج ا ة ا ا ا و ا : ا ي إن ر ،.. أ ي ر .. ودأ ا و إ ء ا ة ى ر ا .. و ا ھ ا و رأت ا ً ل ر و و : ا .. أ ا . ب ا وج : ، ُ ا ھ ا ا ح ُو َز ج ِ ا ... !!
ا : ن أ ؤ ھ ا َأ ل ا س ً، أن ا .
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