AP CHEMISTRY CHAPTER 4 SOLUTION STOICHIOMETRY Electrons aren’t shared evenly (oxygen is more electronegative) Electrons spend more time close to O than to H. This uneven distribution.
Download ReportTranscript AP CHEMISTRY CHAPTER 4 SOLUTION STOICHIOMETRY Electrons aren’t shared evenly (oxygen is more electronegative) Electrons spend more time close to O than to H. This uneven distribution.
AP CHEMISTRY CHAPTER 4 SOLUTION STOICHIOMETRY
Electrons aren’t shared evenly (oxygen is more electronegative) Electrons spend more time close to O than to H.
This uneven distribution of charge makes water polar. Because of this, water is a good solvent. The positive end (H) attracts negative ions or the negative end of another polar molecule. The negative end of water (O) attracts the positive ions or the positive end of another polar molecule.
When water surrounds an ionic crystal, the H end attracts the anion and the O end attracts the cation. This process is called hydration.
Hydration
causes salts (ionic compounds) to dissolve. H 2 O also dissolves polar covalent substances such as C 2 H 5 OH. H 2 O doesn’t dissolve nonpolar covalent substances because there is not enough attraction between the water and the nonpolar molecule.
Hydration
Show the association of the ions with some water molecules when 1 formula unit of KCl dissolves in excess water. Cl K+
A
solution
is a homogeneous mixture.
In a solution, a
solute
dissolves in the
solvent
. If the solute ionizes in the solution, electricity can be conducted and the solute is said to be an
electrolyte
. If the solute ionizes 100% or nearly 100%, it is called a
strong electrolyte
.
Lesser ionization occurs with
weak electrolytes
.
Svante Arrhenius determined that the extent to which a solution can conduct an electrical current depends directly on the number of ions present.
Solubility is usually shown as g/given volume solvent or moles/given volume solution
Strong electrolytes 1.soluble salts 2.strong acids –completely ionize HCl(
aq
), HNO 3 (
aq
), H 2 SO 4 (
aq
)
Ex. Show how HCl dissociates when dissolved in water.
HCl
H
+
+ Cl
-
Acid (Arrhenius) – a substance that produces H + ions in water solution
3.strong bases- completely ionize -contain OH -bitter taste and slippery feel -NaOH, KOH
Weak electrolytes -only ionize slightly (weak acids and bases) HC 2 H 3 O 2 H + + C 2 H 3 O 2 99% 1%
Ammonia (NH 3 ) -weak base NH 3 + H 2 O NH 4 + + OH
Molarity (
M
) = moles of solute liters of solution Molarity is the most common unit of concentration used in Chemistry. We may also see
mM
(millimolar) = moles of solute mL of solution
Ex. Calculate the molarity of a solution made by dissolving 23.4g of sodium sulfate in enough water to form 125 mL of solution.
23.4 g Na 2 SO 4 1 mol Na 2 SO 4 = 0.165 mol Na 2 SO 4 142.06g Na 2 SO 4 0.165 mol = 1.32
M
0.125 L
Ex. How many grams of Na 2 SO 4 are required to make 350 mL of 0.50
M
Na 2 SO 4 ?
0.350L 0.50 mol Na 2 SO 4 142.06g Na 2 SO 1 L 1 mol Na 2 SO 4 4 = 24.9g
Ex. What volume of 1.000
M
KNO 3 must be diluted with water to prepare 500.0 mL of 0.250
M
KNO 3 ?
Dilution problem (
M
1
V
1 =
M
2
V
2 ) (1.000
M
)(
V
1 ) = (0.250
M
)(500.0mL) V 1 = 125 mL Remember, this formula can only be used for dilution! Never use it for a chemical reaction (stoichiometry)!
Read procedure for using volumetric flasks and types of pipets. We will be using both in several labs this year.
Let’s Review Equation Writing from Chemistry I Some reactions fit neatly into a certain “category” of reaction type, some do not.
DECOMPOSITION REACTIONS
Reaction where a compound breaks down into two or more elements or compounds. Heat, electrolysis, or a catalyst is usually necessary.
A compound may break down to produce two elements.
Ex. Molten sodium chloride is electrolyzed.
2NaCl
(l)
2Na + Cl 2
A compound may break down to produce an element and a compound.
Ex. A solution of hydrogen peroxide is decomposed catalytically.
2H 2 O 2 2H 2 O + O 2
A compound may break down to produce two compounds.
Ex. Solid magnesium carbonate is heated.
MgCO 3 MgO + CO 2
Metallic carbonates break down to yield metallic oxides and carbon dioxide.
Metallic chlorates break down to yield metallic chlorides and oxygen.
Hydrogen peroxide decomposes into water and oxygen.
Sulfurous acid decomposes into water and sulfur dioxide.
Carbonic acid decomposes into water and carbon dioxide.
Hydrated salts decompose into the salt and water.
Na 2 CO 3 H 2 O Na 2 CO 3 + H 2 O
ADDITION REACTIONS
Also known as Synthesis, Combination, or Composition
Two or more elements or compounds combine to form a single product.
A Group IA or IIA metal may combine with a nonmetal to make a salt.
A piece of lithium metal is dropped into a container of nitrogen gas.
6Li + N
2
2Li
3
N
Two nonmetals may combine to form a molecular compound.
C + O 2 →CO 2
When an element combines with a compound, you can usually sum up all of the elements on the product side.
Ex. PCl 3 + Cl 2 PCl 5 This is a trick that works because the common positive oxidation states of P are + 3 and +5.
Two compounds combine to form a single product.
Ex. Sulfur dioxide gas is passed over solid calcium oxide.
SO 2 + CaO CaSO 3
A metallic oxide plus carbon dioxide yields a metallic carbonate.
(Carbon keeps the same oxidation state)
A metallic oxide plus sulfur dioxide yields a metallic sulfite.
(Sulfur keeps the same oxidation state)
A metallic oxide plus water yields a metallic hydroxide
.
A nonmetallic oxide plus water yields an acid.
Double Replacement (metathesis)
Two compounds react to form two new compounds.
No changes in oxidation numbers occur
. All double replacement reactions must have a "driving force" that removes a pair of ions from solution.
Formation of a precipitate:
A precipitate is an insoluble substance formed by the reaction of two aqueous substances. Two ions bond together so strongly that water can not pull them apart. You must know your solubility rules to write these net ionic equations!
Simple Rules for Solubility
1. Most nitrate (NO 3
) salts are soluble.
2. Most alkali (group 1A) salts and NH 4 +
3. Most Cl , Br , and I salts are soluble (
are soluble.
NOT Ag + , Pb 2+ , Hg 2 2+ ) 4. Most sulfate salts are soluble ( NOT BaSO 4 , PbSO 4 , HgSO 4 , CaSO 4 ) 5. Most OH salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH) 2 , Ca(OH) 2 are marginally soluble) 6. Most S 2 , CO 3 2 , CrO 4 2 , PO 4 3 salts are only slightly soluble.
SOLUBILITY SONG
To the tune of “ My Favorite Things ” from “The Sound of Music” Nitrates and Group One and Ammonium, These are all soluble, a rule of thumb.
Then you have chlorides, they’re soluble fun, All except Silver, Lead, Mercury I.
Then you have sulfates, except for these three: Barium, Calcium and Lead, you see.
Worry not only few left to go still.
We will do fine on this test. Yes, we will!
Then you have the-- Insolubles Hydroxide, Sulfide and Carbonate and Phosphate, And all of these can be dried!
Ex. Solutions of silver nitrate and lithium bromide are mixed.
AgNO 3 (
aq
) + LiBr(
aq
) AgBr(
s
) + LiNO 3 (
aq
)
Formation of a gas:
Gases may form directly in a double replacement reaction or can form from the decomposition of a product such as
H 2 CO 3 or H 2 SO 3
. H 2 CO 3 H 2 SO 3 NH 4 OH H 2 O and CO 2 H 2 O and SO 2 NH 3 and H 2 O
Ex. Excess hydrochloric acid solution is added to a solution of potassium sulfite.
2HCl(
aq
) + K 2 SO 3 (
aq
) H 2 SO 3 H 2 O(
l
) + SO 2 (
g
) + 2KCl(
aq
) Remember that sulfurous acid decomposes into water and sulfur dioxide!
Ex. A solution of sodium hydroxide is added to a solution of ammonium chloride.
Remember that ammonium hydroxide does not exist.
NaOH(
aq
) + NH 4 Cl(
aq
) NaCl(
aq
) + NH 4 OH NH 3 (
g
) + H 2 O(
l
)
Formation of a molecular substance:
When a molecular substance such as water or acetic acid is formed, ions are removed from solution and the reaction "works".
Ex. Dilute solutions of lithium hydroxide and hydrobromic acid are mixed.
LiOH(
aq
) + HBr(
aq
) LiBr(
aq
) +H 2 O(
l
)
Ex. Gaseous hydrofluoric acid reacts with solid silicon dioxide.
4HF(
g
) + SiO 2 (
s
) SiF 4 (
g
) + 2H 2 O(
l
) This reaction occurs when glass is etched.
Single Replacement
Reaction where one element displaces another in a compound. One element is oxidized and another is reduced.
A + BC B + AC
Active nonmetals replace less active nonmetals from their compounds in aqueous solution.
Each halogen will displace less electronegative (heavier) halogens from their binary salts.
Activity Series of Nonmetals Most Active Least Active F I 2 2 Cl 2 Br 2
Ex. Chlorine gas is bubbled into a solution of potassium iodide.
Cl 2 (
g
) + 2KI(
aq
) 2KCl(
aq
) + I 2 (
s
)
Active metals replace less active metals or hydrogen from their compounds in aqueous solution.
Use an activity series or a reduction potential table to determine activity. The more easily oxidized metal replaces the less easily oxidized metal.
ACTIVITY SERIES OF METALS Element Lithium Potassium Barium Decreasing Calcium Activity Sodium Magnesium Aluminum Zinc Iron Group I,II,+III Cadmium Nickel Tin Lead Transition Metals Hydrogen ( a nonmetal) Copper Mercury Hydrogen Silver Gold Jewelry Metals Platinum *Metals from Li to Na will replace H from water and acids; metals from Mg to Pb will replace H from acids only.
Ex. Magnesium turnings are added to a solution of iron(III) chloride.
3Mg(
s
) + 2FeCl 3 (
aq
) 2Fe(
s
)+3MgCl 2 (
aq
)
Ex. Sodium is added to water.
2Na(
s
) + 2H 2 O(
l
) 2NaOH(
aq
) + H 2 (
g
) Alkali metal demo
COMBUSTION REACTIONS
-Elements or compounds combine with oxygen to produce the oxides of each element.
The oxide of H is H 2 O, oxide of S is usually SO 2 , oxide of C is CO 2 , etc.
Hydrocarbons or alcohols combine with oxygen to form carbon dioxide and water
.
Nonmetallic hydrides combine with oxygen to form oxides and water.
Nonmetallic sulfides combine with oxygen to form oxides and sulfur dioxide.
Ex. Carbon disulfide vapor is burned in excess oxygen.
CS
2
+ 3O
2
CO
2
+ 2SO
2
Ex. Ethanol (C 2 H 5 OH) is burned completely in air.
C
2
H
5
OH + 3O
2
2CO
2
+ 3H
2
O
Diethyl ether (C 2 H 5 OC 2 H 5 ) is burned in air.
C 2 H 5 OC 2 H 5 + 6O 2 4CO 2 + 5H 2 O
Selective precipitation- process by which ions are caused to ppt one by one in sequence to separate mixtures of ions.
Qualitative analysis- process of separating and identifying ions
Ex. Separate Ag + , Ba 2+ , Fe 3+ 1. Add Cl to remove Ag + as AgCl.
2. Add SO 4 2 3. Add OH to remove Ba or S 2 2+ as BaSO to remove Fe 3+ as 4 .
Fe(OH) 3 or Fe 2 S 3 .
Ex. Separate Pb 2+ , Ba 2+ , Ni 2+ 1. Add Cl to remove Pb 2+ as PbCl 2 .
2. Add SO 4 2 3. Add OH to remove Ba or S 2 2+ as BaSO to remove Ni 2+ as 4 .
Ni(OH) 2 or NiS.
Quantitative analysis determines how much of a component is present.
Gravimetric analysis quantitative procedure where a ppt containing the substance is formed, filtered, dried & weighed.
Ex. The zinc in a 1.2000g sample of foot powder was precipitated as ZnNH 4 PO 4 . Strong heating of the ppt yielded 0.4089 g of Zn 2 P 2 O 7 . Calculate the mass percent of zinc in the sample of the foot powder.
0.4089gZn
2 P 2 O 7 1 mol Zn 2 P 2 O 7 2 mol Zn 65.37g = 304.7 g 1 mol Zn 2 P 2 O 7 1 mol Zn 0.1754g Zn 1.200g sample × 100 = 14.62% Zn
Ex. A mixture contains only NaCl and Fe(NO 3 ) 3 . A 0.456g sample of the mixture is dissolved in water, and an excess of NaOH is added, producing a precipitate of Fe(OH) 3 . The ppt is filtered, dried, & weighed. Its mass is 0.128g.
Calculate: a. the mass of the iron b. the mass of Fe(NO 3 ) 3 c. the mass percent of Fe(NO 3 ) 3 in the sample 0.128g Fe(OH) 3 1 mol Fe(OH) 3 1 mol Fe 55.85g Fe= 106.9g Fe(OH) 3 1 mol Fe(OH) 3 1 mol Fe 0.0669g Fe 0.0669g Fe 1 mol Fe 1 mol Fe(NO 3 ) 3 241.9g Fe(NO 3 ) 3 = 0 .290g Fe(NO 3 ) 3 55.85g Fe 1 mol Fe 1 mol Fe(NO 3 ) 3 0.290g
0.456g
× 100 = 63.6% Fe(NO 3 ) 3
Acid-Base Reactions
Bronsted-Lowry acid-base definitions: acid- proton donor base- proton acceptor
When a strong acid reacts with a strong base the net ionic rxn is: H + (
aq
) + OH (
aq
) H 2 O(
l
) When a strong acid reacts with a weak base or a weak acid reacts with a strong base, the reaction is complete (the weak substance ionizes completely.) HC 2 H 3 O 2 (
aq
) + OH (
aq
) H 2 O(
l
) + C 2 H 3 O 2 (
aq
)
neutralization reaction - acid-base rxn When just enough base is added to react exactly with the acid in a solution, the acid is said to be neutralized.
Volumetric Analysis
titration- process in which a solution of known concentration (standard solution) is added to analyze another solution (analyte).
Titrations are most often used for acids and bases, but can be used for other types of reactions, also.
titrant- solution of known concentration (usually in buret) equivalence point or stoichiometric point point where just enough titrant has been added to react with the substance being analyzed
Indicator - chemical which changes color at or near the equivalence point End point- point at which the indicator changes color
Ex. 54.6 mL of 0.100 M HClO 4 solution is required to neutralize 25.0 mL of an NaOH solution of unknown molarity. What is the concentration of the NaOH solution?
HClO 4 + NaOH H 2 O + NaClO 4 0.0546 L HClO 4 0.00546 mol NaOH 0.100 mol HClO 4 1 L HClO 4 1 mol NaOH = 1 mol HClO 4 0.00546 mol NaOH = 0.218
M
0.025L
NaOH
Writing net-ionic equations 1. molecular equation -overall reaction stoichiometry 2. complete ionic equation ions -all strong electrolytes are represented as 3. net ionic equation -spectator ions are not included
1. NaCl(
aq
) + AgNO 3 (
aq
) NaNO 3 (
aq
) + AgCl(
s
) 2. Na + (
aq
) + Cl (
aq
) + Ag + (
aq
) + NO 3 (
aq
) Na + (
aq
) + NO 3 (
aq
) + AgCl(
s
) 3. Cl (
aq
) + Ag + (
aq
) AgCl(
s
)
Oxidation-Reduction Reactions Redox Rxns - reactions in which one or more electrons are transferred.
Electronegativity - attraction for shared electrons most electronegative elements F>O>N=Cl “Phone Call” These are most likely to have negative oxidation numbers.
Rules for Assigning Oxidation States 1. Oxidation state of an atom in an element = 0 2. Oxidation state of monatomic element = charge 3. Oxygen = 2 in covalent compounds (except in peroxides where it = 1) 4. H = +1 in covalent compounds 5. Fluorine = 1 in compounds 6. Sum of oxidation states = 0 in compounds Sum of oxidation states = charge of the ion
Review oxidation state rules on page 167.
+1-2 N 2 +3-1 +1+3-2 O PBr 3 HPO 3 2 +3 -2 -3 +1 P 4 O 6 NH 2 -
Noninteger states are rare, but possible.
8/3 -2 Fe 3 O 4 O = 4(-2) = -8 Fe = 8/3 = 2 2/3 or Fe 2+ , Fe 3+ , Fe 3+
Oxidation - loss of electrons - increase in oxidation number Reduction - gain of electrons - decrease in oxidation number
OIL RIG
Oxidation Is Loss (of e ), Reduction Is Gain (of e )
LEO the lion goes GER
Lose Electrons = Oxidation, Gain Electrons = Reduction
Oxidizing agent - electron acceptor - substance that is reduced Reducing agent - electron donor - substance that is oxidized The terms oxidizing agent and reducing agent are not tested on the AP test.
I +1-1 0 +1-1 0 2KI + F oxidized 2 2KF + I 2 reduced F OA F 2 RA KI
+4 -2 +2 -2 0 2PbO 2 2PbO + O 2 oxidized: O reduced: Pb OA: PbO 2 RA: PbO 2
Balancing redox reactions by the half-reaction method 1.Write skeleton half-reactions.
2.Balance all elements other than O and H.
3.Balance O by adding H 2 O.
4.Balance H by adding H + .
5.Balance charge by adding e to the more positive side.
6.Make the # of e lost = # of e gained by multiplying each half-rxn by a factor.
7.Add half-reactions together.
8.Cancel out anything that is the same on both sides.
9.If the reaction occurs in basic solution, add an equal number of hydroxide ions to both sides to cancel out the hydrogen ions. Make water on the side with the hydrogen ions. Cancel water if necessary.
10.Check to see that charge and mass are both balanced.
Practice: Sn 2+ + Cr 2 O 7 2 (acidic solution) Sn 4+ + Cr 3+ Sn Sn Sn 2+ 2+ 2+ Sn Sn Sn 4+ 4+ 4+ Cr Cr 2 O 7 2 Cr 2 O 7 2 2 O 7 2 Cr 3+ 2 2Cr Cr 3+ 3+ +7H 2 O Sn 2+ Sn 2+ 3( Sn 2+ 3Sn 2+ Sn 4+ Sn 4+ +2e Sn 4+ +2e + Cr 2 O 7 2 Cr 2 O 7 2 ) Cr 2 O 7 2 Cr 2 O 7 2 + 14H + + 14H + +6e 2Cr 3+ +7H 2 O 2Cr 3+ +7H 2 O + 14H + +6e 2Cr 3+ +7H 2 O + 14H + +6e 3Sn 4+ +6e + 2Cr 3+ +7H 2 O 3Sn 2+ 3Sn 2+ + Cr 2 O 7 2 + Cr 2 O 7 2 + 14H + + 14H + +6e 3Sn 4+ +6e + 2Cr 3Sn 4+ + 2Cr 3+ +7H 2 O 3+ +7H 2 O
MnO 4 2 + I MnO 2 + I 2 (basic solution) MnO 4 2 MnO 4 2 MnO MnO 2 2 I 2 I I 2 I 2 MnO 4 2 MnO 2 + 2H 2 O 2I I 2 MnO 4 2 MnO 4 2 +4H +4H + + MnO 2 + 2e MnO 2 + 2H 2 O 2I + 2H 2 O 2I I 2 I 2 +2e MnO 4 2 MnO 4 2 MnO 4 2 MnO 4 2 MnO 4 2 MnO 4 2 +4H + 2H + 2 + 2e O +2I +2I MnO 2 +4H +4H +4H + 4H + + + 2 + 2e O +2I + 4OH MnO 2 +2I + 2H 2 O +I 2 MnO 2 + 2H 2 O +I 2 + 4OH +2I +2I MnO 2 MnO 2 + 2H 2 O +I 2 +2e + 2H 2 O +I 2 + 4OH MnO 2 + I 2 + 2H 2 O +I 2 +2e + 4OH -
OXIDATION-REDUCTION TITRATIONS Most common oxidizing agents: KMnO 4 & K 2 Cr 2 O 7
Potassium permanganate used to disinfect ponds and fish in Egypt.
MnO MnO 4 4 in acidic solution: + 8H + + 5e Mn 2+ + 4H 2 O Purple colorless When you titrate with MnO 4 , the solution is colorless until you use up all of the reducing agent (substance being oxidized).
In calculations, work redox titrations like acid-base titrations. You must have a balanced reaction to know the mole ratio.