a)How is it possible to determine if CaCO3 is Cl- free after synthesis?
b)How can the Cl- ions be remove from CaCO3 after synthesis?
I should answer the questions from the following experiment but if you know the answer and you are sure, yo do not need to read experiment.
Please answer correctly because i hav no chance to make wrong :((((
Physical and Chemical Properties of Pure Substances
The aim of today’s experiment is to learn handling chemicals in laboratory, how to observe chemical reactions and investigate the property of copper salts and Mg, Cu, Zn metals. In part I the solubility of different copper compounds will be investigated. In part II magnesium, copper and zinc metals will be compared and their reactivity will be determined by simple chemical reactions.
Part I. Solubility of salts
Barium sulfate (BaSO4) has a milk like solution, where the fine particles of barium sulfate is suspended in water. It is frequently used medically as a radiocontrast agent and administered orally. At the other hand barium is a heavy metal, and its water-soluble compounds are highly toxic such as barium carbonate (BaCO3) and used as a rat poison. Comparing the solubility of BaSO4 and BaCO3 in water and stomach acid (dilute hydrochloric 0,1-0,01 M) resolves this controversy (see Table 1.). Barium carbonate and barium sulfate are insoluble at 20 °C in water. In dilute hydrochloric acid barium sulfate is insoluble, but barium carbonate reacts with HCl and form BaCl2, a water soluble barium compound. The low solubility of barium sulfate protects the patient from absorbing harmful amounts of the metal.
The solubility of salts can be predicted by applying the 7
solubility rules and is essential in biological systems like body
fluids where more complex salts solutions are in present. The
paragraph below discusses the most important chemical concepts and
solubility rules of salts.
Inorganic compounds can be classified into four major groups: acids, bases, salts, and oxides.
We define an acid, in terms of Arrhenius Theory, as a substance that in water can give a proton (H+) to a water molecule to form the hydronium (H3O+) ion. A base is defined as any substance that generates hydroxide ions (OH-) when dissolved in water. Other theories of acids and bases proposed by Brönsted-Lowry and Lewis are covered in later experiments. Oxides are binary compounds of an element with oxygen. This experiment focuses on salts, ionic compounds formed between metallic and nonmetallic elements.
The use of the term "salt" in chemistry is not limited to sodium chloride. Table salt (sodium chloride) is one of many chemical substances identified as a salt. Chemically we define a salt as a compound that contains a positive cation other than hydrogen and a negative anion other than the hydroxide or oxide ions.
The most common property of all salts is their ability to dissociate into separated ions in aqueous solution. For a salt to dissociate, however, it must be soluble in water. Salts vary their ability to dissolve in water, and thus produce solutions with varying concentrations. The terms "slightly soluble", "soluble" and "insoluble" are used qualitatively to describe such solutions.
The solubility of ionic compounds can be examined by various
An indirect method, used in this experiment requires the mixing of one solution containing the cation of the desired compound with a second solution containing the anion of the desired compound. If a precipitate forms, a double displacement reaction has taken place and has formed at least one insoluble compound.
In the following experiment different types of reactions will be investigated and the solubility of copper compounds will be observed. The experiment goes through a series of chemical reactions similar to the one used for the recycling of copper metal.
8 HNO3(aq) + 3 Cu(s) + O2(g) → 3 Cu(NO3)2(aq) + 4 H2O(l) + 2 NO2(g) (1)
Cu(NO3)2(aq) + 2 NaOH(aq) → Cu(OH)2(s) +2 NaNO3(aq) (2)
Cu(OH)2(s) → CuO(s) + H2O(l) (3)
CuO(s) + H2SO4(aq) → CuSO4(aq) + H2O(l) (4)
Solubility Rules for Salts
The rules governing the solubility of common salts are given below:
1. All nitrates, chlorates, and acetates of all metals are soluble in water. Only silver acetate is sparingly soluble.
2. All sodium, potassium and ammonium salts are soluble in water.
3. The chlorides, bromides and iodides of all metals except lead, silver and mercury(I) are soluble in water. HgI2 is insoluble in water. PbCl2, PbBr2 and PbI2 are soluble in hot water. The water insoluble chlorides, bromides and iodides are also insoluble in dilute acids.
4. The sulfates of all metals except lead, mercury(I), barium and calcium are soluble in water. Silver sulfate is slightly soluble. The water-insoluble sulfates are also insoluble in dilute acids.
5. The carbonates, phosphates, borates, sulfites, chromates and arsenates of all metals except sodium, potassium and ammonium are insoluble in water but soluble in dilute acids. MgCrO4 is soluble in water; MgSO3 is slightly soluble in water.
6. The sulfides of all metals except barium, calcium, magnesium, sodium, potassium and ammonium are insoluble in water.
7. The hydroxides of sodium, potassium, and ammonium are soluble in water. The hydroxides of calcium and barium are moderatly soluble. The oxides and hydroxides of all other metals are insoluble.
Part II. Reactivity of metals
Metal sensitivity (metal hypersensitivity) is a form of an allergic reaction and can be caused by exposure to metals in jewelry, dental implants and orthopedic implants. Reactivity series of metals is based on empirical methods, by observation and experimentation of metals reactivity. It can be used in many way as it offer a simple tool to predict the behavior of the pure metals. An example is if a metal might release metal cations responsible to allergenic reactions, as well as tendency to corrosion attack or damage of the surrounding environment.
In the study of the physical and chemical properties of the elements, a periodic recurrence of similar properties is clearly observed. This observation is formalized in the Periodic Law, and forms the basis for the periodic table of the elements. One readily recognizes a relationship between the position of the element in the table and its atomic and ionic sizes, ionization energy, and electron affinity.
Metals are found to have low ionization energies, which means that less energy is required to remove their outermost (valence) electrons than is required to remove equivalent electrons from the nonmetals. This ionization energy directly reflects the ease with which the isolated atom forms a cation.
The ready loss of the valence electron(s) causes certain metals to react with water to liberate hydrogen. This tendency for metals to produce hydrogen gas from water decreases as their metallic character decreases. Many of the less reactive metals react only in acidic solution, whereas others may even appear to be entirely unreactive.
When a metal generates hydrogen gas from water or acid solution,
atoms of the metal lose electrons to hydrogen ions (H+), converting
them to hydrogen atoms (H) that combine to form hydrogen molecules
(H2). The net process involves the transfer of electrons from the
metal atom to a hydrogen ion, producing the metal cation and
molecular hydrogen. The gain of the electrons by the H+ ions is a
process known as reduction. Chemical changes involving the transfer
of electrons are called oxidation-reduction reactions.
The metals that react spontaneously with acids are said to displace hydrogen ions from the acid and hydrogen gas evolves. These metals are said to be more active than hydrogen. Metals have been listed according to their activities in the reaction of metals with acids. This list is referred to as the reactivity series of elements. The series provides a method of predicting whether as designated redox reaction will or will not occur spontaneously.
In the second part of the experiment Mg, Cu and Zn will be compared based on their reaction with acids and salt solutions.
Comparison of copper salts solubility
Place a paper behind the reaction mixture to be able to distinguish if a precipitate (= insoluble compound) formed during the reaction. If the text on the paper is not readable insoluble compound formed. Observe the reactions carefully!
Write your observation and the chemical reactions to your report!
A piece of copper will be given in a test tube. Another test tube will have 1 ml concentrated HNO3. Make the reaction in the fume hood. Make sure that the fume hood fan is switched on and fully functional. Lower the fume hood door to prevent toxic fumes from entering into the room. Add the copper piece into the test tube containing the HNO3. Observe the reaction and wait till copper is completely dissolved. Add to a beaker 10 ml H2O. Slowly pour the copper solution into the beaker.
Slowly adding 7,5 ml of 3 M NaOH to the solution in the beaker, use a glass rod and gently stir the mixture.
Heat the mixture with a moderate heat or in water bath. Ensure to stir the mixture by gently rotating the beaker. The reaction is complete when the color of the mixture changes entirely.
Continue to stir for an additional minute, then allow the copper (II) oxide to settle to the bottom of the beaker (sedimentation). While the copper (II) oxide is settling, take approximately 50 ml of hot deionized water from the kettle.
Carefully decant the supernatant liquid from the reaction
mixture into a waste container as shown in Figure 1, make sure not
to lose any copper (II) oxide.
Add the hot deionized water to the beaker containing copper (II) oxide. Allow the copper (II) oxide to settle to the bottom of the beaker again and decant the supernatant liquid a second time. As before, avoid losing any copper (II) oxide in the decanting process.
While stirring with a glass rod, slowly add 4 ml of 6 M H2SO4 to the beaker containing copper (II) oxide. The reaction is completed when homogenous clear solution is observed. Set the solution aside for part II.
Part II Relative Reactivity of Mg, Cu and Zn with HCl Acid
1. Pour approximately 10 drops of 6 M HCl into 3 clean test tubes each.
2. Carefully add a small piece-equal in size - of each of the following metals to each subsequent tube: Mg, Cu, Zn.
3. Observe each test tube for signs of heat, color change, or evolution of gas.
4. If possible, indicate the time order in which any changes are observed.
5. Write complete equations for each reaction.
Relative Reactivity of Metal Ions in Solution
1. In the fume hood, into the beaker containing the copper solution remained from part I (4 reaction) add a piece of zinc metal. Use a Pasteur pipette to stir the mixture. Observe the reaction (the surface of the Zn piece)!
CuSO4(aq) + Zn(s) → ZnSO4(aq) + Cu(s)
2. Remember that parallel to the reaction a side reaction occurs. The reaction between H2SO4 and zinc metal (similar to HCl reaction).
Zn(s) + H2SO4(aq) → ZnSO4(aq) + H2(g)
3. Place 10 drops of MgSO4 solution into 1 test tube. Add one small piece of Zn into the test tube. Write your observation into the report sheet!
We can use silver nitrate test to determine the presence of cl- ions. AgNO3 + Hno3 is added to the salt solution taken. In the presence of chlorine ions, a white precipitate is formed which dissolved in ammonium hydroxide. AgNo3 forms yellow precipitate with caco3. You can also use chromyl chloride test or mno2. Once the presence of cl- is confirmed, you can dissolve the salt in water . Pure Caco3 remains as precipitate. Or Cl- can be adsorbed. It can be treated with a cation that reacts only with cl-. ( Hg or Ag)
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