Experiment 3
Chemical Reactions: Qualitative Analysis and Identification of Unknown Aqueous Salts, Acids, and Bases
Purpose: This experiment will use simple observations to describe the interactions of common aqueous salts, acids, and bases and to utilize the properties of color, odor, pH, solubility and precipitation to identify unknown compounds.
Background: Chemical reactions involve mixing together elements and/or compounds under varying conditions of temperature, pressure, and pH, in the presence of catalysts and solvents (generally water) to produce desired products. The reagents, either elements or molecules, interact as salts, polar & non-polar covalent molecules, acids, bases, and gases. The reactions, generally, but not always, take place in aqueous solutions. Salts: The term “Salt” originated in 1884 when Svante Arrhenius proposed that certain compounds, when dissolved in water, produce positive and negative ions. Salts are composed of a metal ion and a non-metal ion (or, non-metallic polyatomic ion). When the salts dissociate in water, the metals form positively charged ions called cations, while the nonmetals form negatively charged ions called anions. The dissociated ions are closely associated (“aquated” from the Latin aqua, or water) with water molecules in aqueous media. Aqueous ions are symbolized by the “(aq)” appendix in chemical equations. Not all water-soluble compounds are salts. Common table sugar (sucrose), for example, is a molecular compound that dissolves in water, but does not dissociated into ions. To summarize these differences, consider the examples below, which include water-soluble example ions and molecular compounds.
H O
2 NaI Na + (aq) + I- (aq)
Cu(NO3 )2
H O
2 Cu 2+ (aq) + 2NO3- (aq)
H O
C12H22O11 (s)
2
(NH4 )2 SO4 (aq)
2
No dissociation,simply C12H22O11 (aq)
H O
2NH4+ (aq)
+ SO-4 (aq)
Once dissociated, the positive and negative ions act independently from each other. This allows their aqueous solution behavior to be studied independently. Any aqueous solution containing the Silver (Ag) ion (Ag+), e.g., AgNO3, will react with any other solution containing the Chloride ion (Cl-), e.g., FeCl3, or ZnCl2, to produce the solid compound Silver Chloride, AgCl(s), which does not dissolve in water. Silver Chloride is an example of a nonsoluble salt termed a precipitate when formed in an aqueous solution. Ag+ (aq) + Cl- (aq)
H O
2
AgCl(s), white solid precipitate
Acids and Bases: Acids and bases are defined according to three different concepts. Arrhenius Concept
Acid: a substance that, when dissolved in water, increases the concentration of Hydrogen ion, H+(aq) or Hydronium ion, H3O+. Base: a substance that when dissolved in water increases the concentration of the Hydroxide ion, OH-(aq). Brønsted-Lowry Concept Acid: species that donates protons Base: species that accepts protons Lewis concept Acid: species that accepts an electron pair Base donates an electron pair Acidity/Basicity: The strength of an acid (acidity) or base (basicity) is the degree of dissociation of the substance producing either hydronium ions or hydroxide ions. A strong acid will dissociate completely into positively charged hydronium ions and the negatively charged anion (conjugate base). Similarly, a strong base will dissociate completely into a positively charged metal cation (conjugate acid) and the negatively charged hydroxide ion.
On the other hand, a weak acid or base dissociates only to a very small degree, generally less than 1%. The relative strength of an acid or base can be quantified by the acid dissociation constant, Ka, (also known as the acidity constant, or acid-ionization constant) and the base counterpart, the base dissociation constant, Kb.
Ka =
Kb =
[H3O+ ][A - ] [HA]
pK a = - log10[K a ]
[HA][OH- ]
pK b = - log10 [K b ]
-
[A ]
Strong acids will have Ka values > 1 and pKa values < 1 Weak acids will have Ka values < 0.01 and pKa values > 1 Strong bases will have Kb values >1 and pKb values 1 pH: pH is another measure of the acidity or basicity of an aqueous solution. It is defined as the negative logarithm of the hydronium ion concentration.
pH = - log[H3O+ ] The pH range of aqueous solutions is 0 – 14 Solutions with a pH less than 7 are said to be acidic and solutions with a pH greater than 7 are basic or alkaline. Pure water has a pH very close to 7. Strong acids have pH values less than 3, while weak acids have pH values between 3 and 7. Strong bases have pH values greater than 11, while weak bases have pH values between 7 and 11. Table 3.1 lists the common acids and bases and their relative strengths.
Table 3.1: Common Acids & Bases: Species
HBr HNO3 H2SO4 -
H3PO4 CH3OOH HF NaOH KOH Ca(OH)2 NH3 CO3 S
2-
2-
Type
Strength
Acid
Strong
2 HCl(g) H3O+ + Cl- (aq)
Acid
Strong
2 HBr(g) H3O+ + Br - (aq)
Acid
Strong
H2 O HNO3 (aq) H3O+ (aq) + NO3- (aq)
Acid
Strong
2 H2SO4 (aq) H3O+ + HSO4-
Acid
Weak
2 HSO4- (aq) H3O+ + SO42-
Acid
Weak
2 H3PO4 H2PO4-
Acid
Weak
2 CH3OOH CH3OO-
Acid
Weak
H2 O HF
Base
Strong
NaOH
Base
Strong
2 KOH K+ + OH-
Base
Strong
Ca(OH)2
Ammonia
Base
Weak
NH3
Carbonate Ion
Base
Weak
CO32-
Sulfide ion
Base
Weak
S2-
Hydrochloric Acid Hydrobromic Acid
HCl
HSO4
Name
Nitric Acid Sulfuric Acid Sulfurous Acid Phosphoric Acid Acetic Acid Hydrofluoric Acid Sodium Hydroxide Potassium Hydroxide Calcium Hydroxide
Ions Produced in Reaction with Water H O
H O
H O
H O
H O
+ H3O+
H O
F-
+ H3O+
H3O+
2 Na + + OH-
H O
H O
2 Ca ++
H O
2
H O
NH4+
+ +
2OHOH-
H2 O HCO3-(aq) + OH- (aq) 2
H O
HS- (aq)
+
OH- (aq)
Salts and acidity/basicity: When salts react with water the resulting solution can acidic, basic or neutral depending on the nature of the anion/cation pairs. A salt consisting of the cation of a strong base, such + 2+ as Na , Ca and the anion of a strong acid, such as Cl and NO3 (see table 3.1 below) will produce a neutral solution (pH = 7.0) because neither of the ions react with water to form + either hydronium ions (H3O ions or hydroxide ions (OH ) ions. Such ions are referred to as spectator ions.
Some salts will produce an acidic or basic solution when they react with water. Some cations act as weak acids by donating a proton water to form hydronium ions. Some anions act as weak bases by taking a proton from water producing hydroxide ions. Table 3.2 provides a review of the common spectator ions and some acidic anions and basic cations. Table 3.2: Common spectator ions, acidic anions, and basic cations Spectator Ions (do not react with water) -
Anions
- -
Cl Br I 2NO3 ClO4 SO4 +
Cations
Basic Anions (react with water to produce OH- ions) 2F CN CO3 HCO3 32- 2NO2 PO4 SO4 S 2C2H3O2- HPO4
+
Acidic Cations (react with water to + produce H3O ions)
2+
NH4+ Al3+ Pb2+ Sn H2PO4 HSO4 Transition Metals 3+ 2+ 2+ (Fe , Zn Pb , etc.)
+
Cs Na K + 2+ 2+ Rb Mg Ca 2+ 2+ Sr Ba
Example: Salt the yields a neutral solution Sodium chloride (NaCl) reacts with water to produce the sodium and chloride spectator ions that do not react with water. NaCl(s) + H2O(l)
Na+(aq) + Cl-(aq) + H2O(l)
Example: Salt that yields an acidic solution -
Iron nitrate (Fe(NO3)3), reacts with water to produce the Nitrate spectator ion, NO3 , 3+
and the Iron(III) ion, Fe , which then reacts with water to produce Hydrogen ions. 3+
-
Fe(NO3)3(s) + 6H2O)l) Fe(H2O)6 (aq) + 3NO3 (aq) 3+
Fe(H2O)6 (aq) + H2O(l)
⇋
2+
+
Fe(H2O)5OH (aq) + H3O
(acidic)
Example: Salt that yields a basic solution Sodium acetate (CH3COONa(s) reacts with water to form the sodium spectator ion, + Na , and the acetate ion (CH3COO ), an anion of acetic acid, a weak acid. CH3COONa(s) + H2O(l) -
CH3OO (aq)
+
H2O
⇋
+
-
Na (aq) + CH3COO (aq) -
CH3COOH + OH (aq)
(basic)
See Appendix E for a more detailed overview of the equilibrium relationships relative to + [H3O ], pH, Kw, Ka, pKa, etc.
Evolution of Gases: A few of the possible unknown anions in the experiment will produce gases when reacted with aqueous acids. You may or may not see any bubbles. The three possibilities include 22Bicarbonate (HCO3 ), Carbonate (CO3 ) and Sulfide (S ). The first two produce odorless 2
Carbon Dioxide gas (CO ), while Sulfide produces a small amount of (H2S(g), which smells like rotten eggs. CO32- (aq) + 2H+ (aq) H2O(aq) + CO2 (g) HCO3- (aq) + H+ (aq) H2O(aq) + CO2 (g) S2- (aq) + 2H+ (aq) H2S(g)
Odor: Solutions of NH3 and Na2S in water have characteristic odors. You should carefully direct the vapors from the CLOSED TEST REAGENT BOATTLE toward your nose. Describe the odor as best you can. Color: Aqueous solutions of compounds of the transition metals have identifying colors. Copper compounds are usually blue, compounds of Nickel are usually green, and compounds of Iron are orange, yellow, or purple. The precipitates that form when two solutions are mixed may also have characteristic colors, which can help identify the unknown. The colors of certain common insoluble salts are given in Table 3.3.
Table 3.3: Colors of Some Common Water-Insoluble Materials:
Carbonates Ag(CO3
=
yellow
CuCO3
=
pale blue
BaCO3
=
white
CaCO3
=
white
AgCl
=
white
HgCl2
=
white
PbCl2
=
white
CuCl2
=
white
Ag(OH
=
grey-brown
Cu(OH)2
=
pale blue
Fe(OH)3
=
rust red
Ni(OH)2
=
pale green
=
white
Zn(OH)2
=
white
=
yellow
HgI2
=
yellow-orange
=
yellow
CuI2
=
brown
Ag2S
=
black
CuS
=
black
Fe2S3
=
black
HgS
=
black
=
black
PbS
=
black
=
white
PbSO4
=
white
Chlorides, Cl
-
Hydroxides, OH
Pb(OH)2
-
-
Iodides, I AgI PbI2
Sulfides, S
2-
NiS Sulfates, SO4 BaSO4
2-
Solubility: Insoluble precipitates will be formed when certain combinations of aqueous solutions are mixed. A precipitate can take the form of cloudiness, fine-grained particles, or heavy solids. A summary of the guidelines for the solubility of ionic salts in aqueous solution is provided in the Table 3.4.
Table 3.4: Solubility of ionic salts in aqueous solution
Soluble
Exceptions -
All common salts of Nitrate ion (NO3 ) +
All common salts of Na, K, NH4 cations -
-
-
Most common salts of Cl , Br , I anions
Insoluble when accompanying + 2+ 2+ cation is Ag , Pb , or Hg2
2-
All common salts of Sulfate ion (SO4 )
2-
Salts of Sulfide ion (S ) when accompanying cation is group 1A metal (Na+, K+), a group 2A metal (Ba2+, Ca2+) or NH4+ Hydroxides of Group 1A metals and larger member of 2+ group 2A metals beginning with Ca 2-
Carbonates (CO3 ) of Group 1A metals and NH4 3-
+
2+
All other common hydroxides are insoluble in water All other common carbonates are insoluble in water
Phosphates of (PO4 ) of Group 1A metals and NH4 +
Insoluble when accompanying 2+ 2+ 2+ cation is Ba , Sr , Ca , 2+ Ag+, Pb All other sulfide compounds are insoluble
+
2+
Most of the transition metal cations (Ag , Cu , Ni , 2+ 2+ Hg , Zn form soluble complexes with aqueous ammonia. Sometime ammonia is used to “dissolve” otherwise insoluble precipitates of these metal cations AgCl(s) + 2NH3 (aq) Ag(NH3 )2+ (aq) + Cl- (aq)
All other common Phosphates are insoluble
The Experiment: Each student or group will be assigned 5 unknowns. The information on the reagent bottle, which includes an unknown identifier, must be recorded accurately in the Pre-lab document. Each unknown will be from one of the following categories:
A strong acid – HCl, HNO3, or H2SO4 A strong base – KOH, NH3, or Na2S A salt of the chloride ion – NaCl, CuCl2, or BaCl2 A salt of the nitrate ion – AgNO3, Cu(NO3)2, or Fe(NO3)3 A salt of the sodium ion – NaI, Na2SO4, or Na2CO3
Each unknown will be tested against eight (8) solutions in addition to observations of color, pH, and odor. The 8 solutions are: 1. KOH(aq)
5. NH3(aq)
2. AgNO3(aq)
6. Na2S(aq)
3. BaCl2(aq)
7. Na2SO4(aq)
4. NaCl(aq)
8. H2SO4(aq)
All solutions are in sets of dropper bottles scatted about the lab. Only a few drops are required for each test. All tests are to be performed where the test solutions are found. Do not remove the caps. Do not remove the test solution bottles from their location. The results of each test are to be recorded as you interpret them. Unknowns are to be identified by number and chemical name. A balanced chemical reaction will also be required for each unknown reaction.
Pre-Lab Report & Notebook: Download from the department data base to your hard drive or flash drive a copy of the lab report template and the data summary table for the Chemical Reactions experiment. http://chem.gmu.edu/templates Print the “Summary Results” table for the Chemical Interactions experiment. Prepare the Pre-lab report according to instructor’s instructions.
Materials and Equipment: Materials 5 Unknowns Aqueous Solutions KOH AgNO3 BaCl2 NaCl Na2SO4 NH3(aq) Na2S H2SO4
Equipment Small, clean test tubes
Procedure: Note: Avoid contamination of test results by yourself or other students, by rinsing the test tubes carefully before and after each test. Do not touch the tip of any dropper bottle to any other test or unknown solution. All test solutions are to be placed in the assigned waste bottle. Any clinging solids should also be rinsed into this waste bottle. 1. Obtain 5 unknowns from instructor’s desk 2. Add approximately 5 drops of an assigned unknown from the dropper bottle to a clean, small test tube. 3. Add approximately 10 drops of the test solution. 4. Shake the mixture gently. 5. Describe the color, odor and reaction results. 6. Repeat any test that gave ambiguous results.
Caution: If you have trouble differentiating between NH3 and KOH: a) Use 5 drops AgNO3 + 10 drops of known NH3 b) Use 5 drops AgNO3 + 10 drops of known KOH c) Use 5 drops AgNO3 + 10 drops unknown; compare to above results.
A false positive can result when Ba2+ and S2- are mixed producing a cloudy whiteness when it should not. It comes about from tests that are not definitive for either ion.
Data Processing: Use the printed Pre-lab report as a notebook to the experimental results in the results section of the applicable procedure. Enter the experimental results into the printed “Chemical Interactions” summary results table and attach to report. If required by the instructor, transfer the laboratory results to the electronic files and finalize the laboratory report.
Analysis and Conclusions: Develop arguments from the test results justifying the identity of each unknown.
Chemical Interactions Test Results Test Color
Odor
pH
KOH
AgNO3
BaCl2
NaCl
NH3
Na2S
Na2SO4
H2SO4
Unk #
Unk #
Unk #
Unk #
Unk #
