1. Basic Tests

2. Objective Simple Tests
To identify the contents of different test tubes using a few simple tests and their mutual reactivities.
Simple Tests
Color – Transition metals tend to form brightly colored compounds.
Odor – Some compounds have very distinctive odors.
When testing for odor always remember to WAFT!
pH – Using universal litmus paper determine the pH of
the compound.
Solubility – By combining the unknowns and knowing the rules of solubility, one can determine the contents of the test tubes.
Flame Test – Metal ions when introduced into a flame give a distinct emission spectrum. The color of the flame can help identify the unknown metal.

3. Definitions Applicable to Ionic Reactions
Ions – Charged Species. Metals tend to form cations
and Nonmetals tend to form anions.
Ionic substances tend to dissolve readily in water to
form solutions because they are charged
particles that should electrostatically attract
the corresponding end of the water dipole. However, not all ionic substances are soluble in water,
indicating that they do not have enough energy to break apart the ionic crystal.
Cations – Positively charged ions. Cations in today’s
experiment include: H+, Na+, K+, Ca2+, Ba2+,
Fe3+, Cu2+, Ni2+, and Sn2+.
Anions – Negatively charged ions. Anions in today’s
experiment include: Cl-, S2-, NO3-, SCN-, SO42-,
CrO42-, and C2042-.
Lagowski & Sorum – Introduction to Semimicro Qualitative Analysis p77

4. What are ions
Ions are atoms that have lost or gained their electrons and as a result are no longer neutral in charge.
The resulting ion has an octet (8 valence electrons).
Metals lose electrons & form positive ions.
Nonmetals gain electrons & form negative ions.

5. Ions – Memorize these!
Cations are positive and are formed by elements on the left side of the periodic chart.
Anions are negative and are formed by elements on the right side of the periodic chart.
These ions are formed from a single atom and are called Monoatomic ions.
© 2009, Prentice-Hall, Inc.

6. Ions with Multiple Charges
Usually transition metals
Two Examples: Memorize these two! Cuprous (Cu 1+) vs Cupric (Cu 2+)
Ferrous (Fe 2+) vs Ferric (Fe 3+)
Pattern:
Lois Lane is superman’s girlfriend.
The lower charge is the –ous ion.
Get it? Low-ous (Lois)

7. Polyatomic Ions
Polyatomic Ion - A group of covalently bonded atoms that have lost or gained electron(s) and as a result have an electronic charge.
Example:
SO = Sulfate Ion

8. Common Anions - Memorize
© 2009, Prentice-Hall, Inc.

9. Common Cations - Memorize
© 2009, Prentice-Hall, Inc.

10. Soluble - The term soluble means that a substance dissolves. 
An aqueous solution is soluble.  If one mixes two solutions together and no precipitate forms, then only the ions are
in solution.  Thus, there is no reaction.
Note: Sometimes two solutions are mixed together, a reaction
can occur that does not form a precipitate.  Usually when this type of reaction takes place, there is a marked
color change when the product is formed or a large
temperature change is observed.
Insoluble - The term insoluble means a
substance does not dissolve. 
Precipitate – A solid beneath a liquid.  If one
mixes two solutions and a solid forms,
this is called a precipitation reaction.

11. Solubility Rules
1. All nitrates, chlorates, and acetates of all metals are soluble. Silver acetate is sparingly soluble.
2. All sodium, potassium, and ammonium salts are soluble.
3. All chlorides, bromides, and iodides are soluble except silver, lead (II), and mercury (I).
4. All sulfates are soluble except barium, calcium, strontium,
lead (II), and mercury (I).
5. Carbonates, phosphates, borates, sulfites, chromates, and
arsenates of sodium, potassium, and ammonium are soluble; all others are insoluble.
6. Sulfides of barium, calcium, magnesium, sodium, potassium,
and ammonium are soluble; all others are insoluble.
7. Hydroxides of sodium, potassium, and ammonium are soluble. Hydroxides of barium and calcium are moderately soluble.
8. Everything else will be considered insoluble!

12. Ionic Equations & Precipitation Reactions
Overall Equation – Shows reactants and products as undissociated, electrically neutral compounds.
AgNO3(aq) + NaCl(aq) NaNO3(aq) + AgCl(s)
Complete Ionic Equation – Shows the state of reactants and products as hydrated or other phases.
Ag+1(aq) + NO3-1(aq) + Na+1(aq) + Cl-1(aq) Na+1(aq) + NO3-1(aq) + AgCl (s)
Net Ionic Equation – Shows only the reactants and products that are directly involved in the reaction.
Ag+1(aq) + Cl-1(aq) AgCl(s)

13. Spectator Ions – Ions which are not directly involved in the net ionic equation are called spectator ions. In the previous equation, this would be the sodium and nitrate ions.
Ag+1(aq) + NO3-1(aq) + Na+1(aq) + Cl-1(aq) Na+1(aq) + NO3-1(aq) + AgCl (s)
Ag+1(aq) + NO3-1(aq) + Na+1(aq) + Cl-1(aq) Na+1(aq) + NO3-1(aq) + AgCl (s)
(Spectator Ions)
Ag+1(aq) + Cl-1(aq) AgCl(s)
(Net Ionic Equation)

14. Ions
Ions are atoms that have lost or gained their electrons and as a result are no longer neutral in charge.
The resulting ion has an octet (8 valence electrons).
Metals lose electrons & form positive ions.
Nonmetals gain electrons & form negative ions.

15. Ions – Memorize these!
Cations are positive and are formed by elements on the left side of the periodic chart.
Anions are negative and are formed by elements on the right side of the periodic chart.
These ions are formed from a single atom and are called Monoatomic ions.
© 2009, Prentice-Hall, Inc.

16. Ions with Multiple Charges
Usually transition metals
Two Examples: Memorize these two! Cuprous (Cu 1+) vs Cupric (Cu 2+)
Ferrous (Fe 2+) vs Ferric (Fe 3+)
Pattern:
Lois Lane is superman’s girlfriend.
The lower charge is the –ous ion.
Get it? Low-ous (Lois)

17. Polyatomic Ions
Polyatomic Ion - A group of covalently bonded atoms that have lost or gained electron(s) and as a result have an electronic charge.
Example:
SO = Sulfate Ion

18. Common Anions - Memorize
© 2009, Prentice-Hall, Inc.

19. Common Cations - Memorize
© 2009, Prentice-Hall, Inc.

20. Cation Identification
When testing cation, the following methods may be employed:
Chemical Color Identification
Solubility
Flame Test for Identification
Addition of sodium hydroxide solution
Addition of ammonia solution
Addition of potassium iodide solution
Ammonium Identification
Addition of Hydrochloric Acid
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21. Anion Identification
When testing anion, the following methods may be employed:
Addition of silver nitrate to a slightly acidic solution of compound
Addition of sulfuric acid*
Polyatomic anion tests
Addition of Lead(II) Nitrate
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22. Chemical Color Identification
Compounds of some transition metals have unique color:
Copper: blue, blue/green
Chromium: green, purplish/green
Iron: pale green, yellow, amethyst
Nickel: green
Cobalt: red
• Students must realize that this is only a starting point to guide more testing.
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23. Solubility The following are insoluble: Halides of lead and silver
Sulfates of barium, calcium, strontium, and lead
Carbonates of all metals with the exception of alkali metals
• Students must realize that this is only a starting point to guide more testing.
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24. Sodium Hydroxide Addition
Carefully add dilute NaOH solution to solution of unknown
White precipitate: calcium, zinc, strontium, aluminum, and lead
Blue precipitate, but turns black upon heating: copper
Green precipitate: Nickel
Gray/green, turning brown on standing: Iron(II)
Red precipitate: Iron(III)
Blue precipitate, turning gray upon standing: Cobalt
• 1M stock solution of NaOH should be prepared by teacher and placed in dropper bottles.
• Litmus paper must be available
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25. Ammonia Solution Addition
Carefully add dilute NH3 solution to solution of unknown
White precipitate: calcium, aluminum, lead, and zinc
Blue precipitate, turning to deep blue solution as more NH3 is added : copper
Coral Blue Solution: nickel
Gray/green, turning brown on standing: Iron(II)
Red precipitate: Iron(III)
Blue precipitate, turning gray upon standing: Cobalt
• Stock solution of ammonia must be prepared and stored in labeled dropper bottles.
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26. Addition of Potassium Iodide Solution
Carefully add dilute KI solution to solution of unknown
Yellow precipitate: lead
• Potassium iodide solution must be prepared and stored in labeled dropper bottles.
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27. Addition of Potassium Chromate Solution
Carefully add dilute potassium chromate solution to solution of unknown
Fine crystalline precipitate: zinc, cobalt, iron(III), nickel, or lead
• Potassium iodide solution must be prepared and stored in labeled dropper bottles.
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28. Addition of HCl Solution
Carefully add dilute HCl solution to solution of unknown
White Precipitate: lead
• Potassium iodide solution must be prepared and stored in labeled dropper bottles.
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29. Addition of Silver Nitrate
Carefully add a dilute AgNO3 solution to a slightly acidic solution of unknown
Create slightly acidic solution by adding several drops of HNO3
White precipitate: chloride
Cream precipitate: bromide
Yellow precipitate: iodide
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30. Addition of Sulfuric Acid
*MUST BE CONDUCTED IN FUME HOOD
Carefully add several drops of concentrated sulfuric acid to solution of unknown
Brown acidic fumes: bromide or nitrite
Purple acidic fume: iodide
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31. Addition of Lead(II) Nitrate
Carefully add several drops of lead(II) nitrate solution to a solution of unknown
Yellow precipitate: iodide
White precipitate: bromide, chloride or sulfate
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32. Flame Tests
The flame test is a procedure used in chemistry to detect the presence of certain metal ions, based on each element's characteristic emission spectrum. The color of flames in general also depends on temperature.
The flame test is fast and easy to perform, and does not require any equipment not usually found in a chemistry laboratory. However, the range of detected elements is small, and the test relies on the subjective experience of the experimenter rather than any objective measurements.

33. Flame Tests
The test involves introducing a sample of the element or compound to a hot, non-luminous (blue) bunsen flame, and observing the color that results.
Flame Test Calcium
Samples are usually held on a nichrome wire cleaned with hydrochloric acid to remove traces of previous analytes.
Glass Rod with Nichrome Wire

34. Flame Tests Upon using cue tip dip flame test, color emission will be:
Sodium: bright yellow/orange
Potassium: lilac
Strontium/Lithium: bright red
Barium: pale green
Calcium: brick/orange red
Copper: green with blue/white center
Lead: Whitish in color
• Prior laboratory involving the teaching of the Flame Test Method must be completed.
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35. Flame Tests
Potassium - Purple
Barium - Green
Sodium - Yellow
Sodium is a common component or contaminant in many compounds and its spectrum tends to dominate over others. Thus the color yellow overpowers the true color.
The test flame is often viewed through cobalt blue glass to filter out the yellow of sodium and allow for easier viewing of other metal ions.

36. Flame Tests *Ba Barium - Apple Green *Na Sodium - Intense Yellow
Mg Magnesium - Brilliant white
Mn(II) Manganese(II) –
Yellowish green
Mo Molybdenum - Yellowish green
*Na Sodium - Intense Yellow
P Phosphorus - Pale bluish green
Pb Lead - Pale green
Rb Rubidium - Pale violet
Sb Antimony - Pale green
Se Selenium - Azure blue
Sr Strontium - Crimson Red
Te Tellurium - Pale green
Tl Thallium - Pure green
Zn Zinc - Bluish Green
As Arsenic - Blue
B Boron - Bright Green
*Ba Barium - Apple Green
Ca Calcium - Brick Red
Cs Cesium - Pale Violet
Cu(I) Copper(I) - Blue
Cu(II) Copper(II) (non-halide) -
Green
*Cu(II) Copper(II) (halide) -
Blue-Green
*Fe Iron - Gold
In Indium - Blue
*K Potassium - Lilac
Li Lithium – Carmine Red

37. Common Compounds H2SO4 K2CrO4 Fe(NO3)3 Na2S NiSO4 KNO3 Ba(NO3)2 NH3
NaCl
K2C2O4
Cu(NO3)2
SnCl2
KSCN

38. H2SO4
Bio: Strong Acid, most powerful industrial chemical in the world, may produce insoluble sulfates if “metal”ed (meddled) with.*
*Recall #4 of our solubility rules:
All sulfates are soluble except barium, calcium, strontium, lead (II), and mercury (I).
Sulfuric Acid

39. MSDS for H2SO4
Corrosive; highly exothermic reaction with water. Burns from sulfuric acid are potentially more serious than those of comparable strong acids (e.g. hydrochloric acid), as there is additional tissue damage due to dehydration and particularly due to the heat liberated by the reaction with water; i.e. secondary thermal damage.
The danger is greater with more concentrated preparations of sulfuric acid; however, even the "dilute" ~ 0.1 M H2SO4 will char paper by dehydration if left in contact for a sufficient while.

40. Ammonia
NH3
Bio: Alias Ammonium Hydroxide (NH4OH) has done important work in homes, last known employment as fertilizer, can turn ugly on
any nosey detectives.
A gas with a characteristic pungent odor. Caustic and can cause serious health damage. Exposure to very high concentrations of gaseous ammonia can result in lung damage and death.

41. K2CrO4 Potassium Chromate
Bio: Best known for its bright disposition, potassium ion is almost inert, but the chromate may drop out if faced with silver, lead or barium. Remains bright even when it lays low.
Potassium Chromate

42. K2CrO4
Potassium Chromate is very toxic and may be fatal if swallowed. It may also act as a carcinogen, and can create reproductive defects if inhaled or swallowed.
It is a strong oxidizing agent. It may react rapidly, or violently.
It is also possible that it may react explosively with other reducing agents and flammable objects.

43. *i.e., the flame test will yield a yellow flame.
NaCl
Bio: Nothing but a common salt, almost impossible to recognize in a crowd, but shows quite a yellow streak when the real heat is on.*
Sodium Chloride
*i.e., the flame test will yield a yellow flame.

44. Ferric
Fe 3+
Ferrous
Fe 2+
Fe(NO3)3
Bio: Alias “Iron III” – Ferric is more reactive than younger brother Ferrous; may be recognized by color
Ferric Nitrate
(*More about that in the KSCN slide.)

45. K2C2O4
Potassium Oxalate
Bio: Actions not well known, but moderate toxicity noted, handle with care,

46. Na2S
Bio: Alias “Le Pew”, a real loner, tends to linger on the skin if touched.
(Do NOT Touch!)
Caution: Na2S + H2SO4 yields which smelly gas?
Sodium Sulfide

47. Cu(NO3)2
Bio: First name officially changed to “Copper II”; leading chemical citizen, many business ventures include electrical wire manufacturing and production of alloys, notably brass; in solution is easily recognizable by “blue colour”.
Cupric Nitrate
ammonia make darker blue.

48. Caution: Nickel salts are considered carcinogenic.
NiSO4
Bio: Once very valuable, now net worth greatly reduced, “Nick” is easily recognized by his (green) nature.
Nickel Sulfate
Caution: Nickel salts are considered carcinogenic.

49. SnCl2 Stannous Chloride
Bio: a.k.a. Tin Chloride, a hard worker, Tin known since ancient times, currently employed in food packaging industry, recyclable; fluoride salt prevents tooth decay;
Note: Stannous Chloride was
prepared in 1 M HCl, so it will
appear quite acidic.
Stannous Chloride

50. Unfortunate Confrontation
with “Le Pew”
Solutions of tin ( II) chloride can also serve simply as a source of Sn 2+ ions, which can form other tin (II) compounds via precipitation reactions, for example brown (or black) tin (II) sulfide:
SnCl2(aq) + Na2S(aq) → SnS(s) + 2 NaCl(aq)
- Stannous chloride is an inorganic compound made from the soft metal known as tin. The economically significant substance is utilized in a variety of applications, but primarily electroplating, the dyeing of textiles, tin galvanizing, and removing ink stains. Stannous chloride is also sometimes added to foods as a preservative and to oils as an antisludge agent. The compound typically exists as colorless crystals that decompose under high levels of heat. Stannous chloride is soluble in less than its own weight of water, but in the presence of large amounts of water, it forms an insoluble basic salt.
stannous sulfide (SnS) Dark crystals; insoluble in water, soluble (with decomposition) in concentrated hydrochloric acid; melts at 880°C; used as an analytical reagent and catalyst, and in bearing material.
Also known as tin monosulfide; tin protosulfide; tin sulfide.

51. KNO3
Potassium Nitrate
Can be distinguished from the other “common salt” by its pale violet response to any “trial by fire.”

52. KNO3 (aq)
KNO3 (s)
The solid is a critical oxidizing component of black powder gunpowder. In the past it was also used for burning fuse technologies including slow matches.
It readily precipitates and was widely "harvested" since the Late Middle Ages and Early Modern era through the 19th century from urine from which it was forced to crystallize in various odorous ways. Its common names include saltpeter, & Nitrate of potash.
The name Chile saltpetre is also applied to sodium nitrate, which while related to explosives as well, is a very different compound.

53. KSCN Potassium Thiocyanate
Bio: Poisonous little creature, approach with caution,.
Potassium Thiocyanate
by mixing Iron(III) compounds with potassium Thiocyanate. The Chemical we get is similar to the Iron-containing part of hemoglobin and is a blood red color.*
*Note: It may appear black in the well, use a toothpick to smear some on a piece of filter paper to verify color.
KSCN + FeCl3

54. Ba(NO3)2 .
Barium Nitrate
Toxic by ingestion or inhalation. Symptoms of poisoning include tightness of muscles (especially in the face and neck), vomiting, diarrhea, abdominal pain, muscular tremors, anxiety, weakness, labored breathing, cardiac irregularity, and convulsions. Death may result from cardiac or respiratory failure, and usually occurs a few hours to a few days following exposure to the compound. Barium nitrate may also cause kidney damage.