In order to depict the formula of a salt graphically, you should:

1. Write the empirical formula of this compound correctly.

2. Considering that any salt can be represented as a product of neutralization of the corresponding acid and base, the formulas of the acid and base that form this salt should be depicted separately.

For example:

Ca(HSO 4) 2 - calcium hydrogen sulfate can be obtained by incomplete neutralization of sulfuric acid H 2 SO 4 with calcium hydroxide Ca(OH) 2.

3. Determine how many molecules of acid and base are required to obtain a molecule of this salt.

For example:

To obtain a Ca(HSO 4) 2 molecule, one molecule of base (one calcium atom) and two molecules of acid (two acid residues HSO 4  1) are required.

Ca(OH) 2 + 2H 2 SO 4 = Ca(HSO 4) 2 + 2H 2 O.

Next, you should construct graphic images of the formulas of the established number of molecules of the base and acid and, mentally removing the hydroxyl anions of the base and hydrogen cations of the acid that participate in the neutralization reaction and form water, obtain a graphic image of the formula of the salt:

O – H H - O O O O

Ca + → Ca + 2 H - O - H

O – H H - O O O O

H- O O H- O O

Physical properties of salts

Salts are crystalline solids. Based on their solubility in water, they can be divided into:

1) highly soluble,

2) slightly soluble,

3) practically insoluble.

Most salts of nitric and acetic acids, as well as potassium, sodium and ammonium salts, are soluble in water.

Salts have a wide range of melting and thermal decomposition temperatures.

Chemical properties of salts

The chemical properties of salts characterize their relationship to metals, alkalis, acids and salts.

1. Salts in solutions interact with more active metals.

A more active metal replaces a less active metal in the salt (see Appendix Table 9).

For example:

Рb(NO 3) 2 + Zn = Рb + Zn(NO 3) 2,

Hg(NO 3) 2 + Cu = Hg + Cu(NO 3) 2.

2. Salt solutions react with alkalis, this produces a new base and a new salt.

For example:

CuSO 4 + 2KOH = Cu(OH) 2  + 2K 2 SO 4,

FeCl 3 + 3NaOH = Fe(OH) 3 + 3NaCl.

3. Salts react with solutions of stronger or less volatile acids, this produces a new salt and a new acid.

For example:

a) as a result of the reaction, a weaker acid or a more volatile acid is formed:

Na 2 S + 2HC1 = 2NaCl + H 2 S

b) reactions of salts of strong acids with weaker acids are also possible if the reaction results in the formation of a slightly soluble salt:

СuSO 4 + Н 2 S = СuS + H 2 SO 4 .

4. Salts in solutions enter into exchange reactions with other salts, this produces two new salts.

For example:

NaС1 + AgNO 3 = AgCl + NaNO 3,

CaCI 2 + Na 2 CO 3 = CaCO 3  + 2NaCl,

CuSO 4 + Na 2 S = CuS+ Na 2 SO 4.

It should be remembered that exchange reactions proceed almost to completion if one of the reaction products is released from the reaction sphere in the form of a precipitate, gas, or if water or other weak electrolyte is formed during the reaction.

9TH GRADE

Continuation. See No. 34, 35, 36, 37, 38/2003

Practical work № 13.
Nitric acid. Nitrates
(end)

HNO 3 is a strong oxidizing agent. Concentrated nitric acid oxidizes nonmetals to higher oxidation states:

Passivation occurs due to the formation of insoluble films of metal oxides:

2Al + 6HNO 3 = Al 2 O 3 + 6NO 2 + 3H 2 O.

HNO 3 (conc.) can be stored and transported without air access in containers made of Fe, Al, Ni.
A qualitative reaction is the interaction of HNO 3 with Cu to form a brown NO 2 gas with a pungent odor (in addition, salt and water are formed).

As the concentration (dilution) decreases, HNO 3 with Zn can form various nitrogen-containing products:

and also in all cases salt and water.

Note . To recognize the nitrate anion, a diphenylamine indicator is used (a solution of (C 6 H 5) 2 NH in conc. H 2 SO 4).
Demonstration experience . Recognition is carried out “by traces” or by droplet contact: a dark blue color appears.

Nitrates– salts of nitric acid, crystalline solids, highly soluble in water. Nitrates alkali metals, calcium and ammonium – saltpeter.
Most nitrates are very good mineral fertilizers.
Nitrates are strong oxidizing agents! Coal, sulfur and other flammable substances burn in molten nitrate, since all nitrates (like HNO 3) release oxygen when heated and, depending on the chemical activity of the metal, salts give different products:

Operating procedure	Quests	Observations and conclusions
Assemble the device (according to the diagram), put a little crystalline sodium (Chilean) nitrate in a cup, melt it. Heat a piece of charcoal in the flame of a spirit lamp and lower it into the molten saltpeter	Why does coal catch fire? Write equations for the reactions occurring based on the electronic balance, draw appropriate conclusions
Take samples of all three solutions in test tubes No. 1–3 (see No. 38/2003) and first pour approximately an equal amount (volume) of concentrated sulfuric acid into each sample, then add a little copper shavings and heat a little. Observe characteristic changes in one of the samples	Three numbered test tubes contain solutions of sodium chloride, sulfate and sodium nitrate. Recognize the saltpeter solution. Why is concentrated sulfuric acid first added to the nitrate solution? Write molecular and ionic equations for the reaction. Check the output using a trace reaction with a diphenylamine indicator.

Complex substances (turpentine, wood, sawdust) can also burn in nitric acid. A mixture of concentrated nitric and sulfuric acids (nitrating mixture) with many organic substances forms nitro compounds (nitration reaction).
A mixture of 1 volume of HNO 3 (conc.) and 3 volumes of HCl (conc.) is called “aqua regia”. Even gold Au and platinum Pt dissolve in such a mixture:

Operating procedure	Quests	Observations and conclusions
Add some copper shavings (Cu) to a test tube with concentrated nitric acid (1 ml). If the effect is delayed, warm it up a little. Work under traction! Pour the products from the sanitary bottle into the sewer system and rinse with a stream of water.	What explains the release of brown gas with a pungent odor? Considering that water and copper(II) nitrate are still formed, write the reaction equation. Draw up an electron balance diagram and write the reaction equation in ionic form
Mix fine-crystalline sulfur (S) powder with 1 ml of concentrated HNO 3, heat the mixture (under draft). Take a sample of the reaction products and test it with 2-3 drops of barium chloride solution. Immediately pour products from the sanitary bottle into the sewer system	What explains the observed changes - dissolution of sulfur, release of brown, pungent-smelling gas (and water)? Write the equation for this reaction. Draw up an electron balance diagram and an ionic equation for the reaction. What do the changes observed when a sample of the reaction products interacts with a solution of barium chloride prove? Justify the answer

Practical work 14.
Determination of orthophosphates

Goals. Learn to recognize orthophosphates, hydroorthophosphates and dihydrogen orthophosphates by their solubility in water, hydrolysis, and qualitative reaction to the orthophosphate anion.
Equipment and reagents. A rack with test tubes, glass tubes with rubber rings, a sanitary bottle, spatulas (3 pcs.); crystalline Ca 3 (PO 4) 2, CaHPO 4, Ca(H 2 PO 4) 2, distilled water, universal indicator, solutions of H 3 PO 4, NaCH 3 COO (= 10%), AgNO 3.

Operating procedure	Quests	Observations and conclusions
Pour 1 cm 3 of calcium orthophosphate, hydrogen orthophosphate and calcium dihydrogen orthophosphate into three test tubes, add a little (the same amount) of water, mix	Draw a conclusion about the solubility of primary, secondary and tertiary orthophosphates. Can the different solubilities of these phosphates be considered a method for their recognition?	…
Using aqueous solutions and suspensions in three test tubes from the previous experiment, test them with a universal indicator	Determine the pH of all solutions on a scale and explain why the pH in this case has different values	…
K aqueous solution of orthophosphoric acid in one test tube (1 ml) and superphosphate solution in another (1 ml) add 10% sodium acetate solution and a few drops of silver(I) nitrate	What is the reagent for an ion?? Write the equations of the corresponding reactions in molecular and ionic forms, indicate the signs of the reactions	…

Practical work 15.
Determination of mineral fertilizers.
Solving experimental problems on the topic
"Nitrogen subgroup"

Goals. Review the composition and properties of nitrogen and phosphorus compounds, their interconversions and methods of recognition.
Equipment and reagents. Alcohol lamp, matches, blue glass, filter paper, test tube holder, rack with test tubes (2 pcs.), spatulas (3 pcs.), mortar, pestle, sanitary bottle;
in test tubes No. 1–3:
Option I – double superphosphate, NH 4 NO 3, (NH 4) 2 SO 4,
Option II – NH 4 Сl, NaNO 3, KCl,
Option III – KNO 3, (NH 4) 2 SO 4, double superphosphate;
crystalline salts (NH 4) 2 SO 4, NH4Сl, ammophos, aqueous solutions of CH 3 COONa (= 10%), AgNO 3, BaCl 2,
CH 3 COOH ( = 10%), NaOH, litmus paper, CuO, Cu (chips), HNO 3 (dil.), HNO 3 (conc.), H 2 SO 4 (conc.), diphenyl indicator, (C 6 H 5) 2 NH in concentrated H 2 SO 4,
Ca(OH) 2 (dry), distilled water, AgNO 3 in HNO 3, in test tubes No. 4–6 dry crystalline substances: Na 2 SO 4, NH 4 Cl, NaNO 3, in test tubes No. 7 and 8: H 3 PO 4 and H 2 SO 4 (diluted solutions), in test tubes No. 9 and 10: Na 3 PO 4 and Ca 3 (PO 4) 2.

Experimental task . Four numbered bottles contain aqueous solutions of sodium orthophosphate, ammonium sulfate, sodium nitrate, and potassium chloride. Using the most rational recognition methods (see table), determine where each substance is located.

Characteristic signs some salts
(recognition methods)

Table

Substance name	Appearance	Solubility (in water)	The interaction of a solution of this salt with				Flame coloring
			H2SO4 (conc.) and Cu	solutions of BaCl 2 and CH 3 COOH	NaOH solution when heated	AgNO 3 solution
Ammonium nitrate NH 4 NO 3		good	NO 2, brown, with a pungent odor	–	NH 3, colorless, with a pungent odor	–	Yellow (from impurities)
Ammonium chloride NH 4 Cl	White crystalline powder	good	–	–	NH 3	AgCl, white precipitate	Yellow (from impurities)
Potassium nitrate KNO 3	Light gray small crystals	good	NO 2	–	–	–	Purple
Ammonium sulfate (NH 4) 2 SO 4	Colorless large crystals	good	–	BaSO 4, white, insoluble in CH 3 COOH	NH 3	Ag 2 SO 4, white, highly soluble in acids	–
Superphosphate Ca(H 2 PO 4) 2 2H 2 O	Light gray powder or granules	Dissolves slowly	–	Ba 3 (PO 4) 2, white, partially soluble in CH 3 COOH	–	Ag 3 PO 4, yellow (in the presence of CH 3 COONa)	Brick- red
Silvinite KCl NaCl	Pink crystals	good	–	–		AgCl	Yellow with hints of purple
Potassium chloride KCl	Colorless crystals	good	–	–	–	AgCl	Purple

Solution

All ions in aquatic environment colorless, it is impossible to recognize them by color.
2) Since none of the substances (flasks No. 1–4) are characterized by worse solubility, the solutions cannot be distinguished by this criterion; all are transparent solutions.
3) The same cations are present in two solutions, but different anions are present in all of them, so qualitative recognition should be carried out based on the anions. Reagent for – AgNO 3 in the presence of a 10% solution of CH 3 COONa (or BaCl 2 and CH 3 COOH); reagent – ​​BaCl 2 solution; reagent for Cl – – solution of AgNO 3 in HNO 3; reagent - concentrated H 2 SO 4 and Cu (chips). You can immediately identify, then, using one reagent (AgNO 3), recognize all three remaining solutions (or vice versa). Other options take longer and require significantly more reagents.
4) Test all four solution samples with AgNO 3 solution (1–2 drops): the solution from bottle No. 4 remained unchanged - it should be a NaNO 3 solution; in flask No. 2 there is a white crystalline precipitate, insoluble in acids, this is a KCl solution; the other two samples give cloudy solutions, when added to which a 10% solution of CH 3 COONa, sample No. 3 gives a precipitate yellow is a solution of Na 3 PO 4, and sample No. 1 is a solution of (NH 4) 2 SO 4 (the turbidity disappears when the acid HNO 3 is added).

Verification of primary tests.

Add 1-2 drops of BaCl 2 and CH 3 COOH solutions to the sample solution from bottle No. 1, the solution becomes milky in color, because a white crystalline precipitate precipitates:

You can check the same sample by adding an alkali solution with heating. NH 3 gas is released, determined by the characteristic odor and blueness of wet red litmus paper. Reaction equation:

Add concentrated H 2 SO 4 and Cu (shavings) to the solution sample from bottle No. 4 and heat slightly. A brown gas with a pungent odor is released, and the solution becomes greenish-azure in color:

5) Conclusion .

In bottles:
No. 1 – solution (NH 4) 2 SO 4,
No. 2 – KCl solution,
No. 3 – Na 3 PO 4 solution,
No. 4 – NaNO 3 solution.

Recognition scheme

Determined solutions
№ 1	№ 2	№ 3	№ 4
(NH 4) 2 SO 4	KCl	Na3PO4	NaNO3
All solutions are transparent and colorless
+AgNO3
Cloudiness of the solution (Ag 2 SO 4, soluble in acids)	White cheesy sediment (AgСl	According to the option, write down which salt solutions are given in test tubes No. 1–3. Determine where each of these substances is located. In the conclusions, write down the equations of the reactions carried out in molecular and ionic forms. Note the signs of each qualitative reaction	…
1) In a test tube with a small amount CuO (at the tip of a spatula), add HNO 3 solution, shake. 2) Place some copper shavings into a test tube with concentrated HNO 3 (if the effect is not immediately observed, warm the mixture a little)	Using the given reagents, prepare a solution of copper(II) nitrate in two ways. Note the signs of reactions and write molecular and ionic reaction equations. Which reaction is redox?	…
In a mortar, mix and grind the mixture of Ca(OH) 2 (slightly moistened) with ammonium salt, sniff carefully. Repeat the experiment with other ammonium salts	To prove experimentally that sulfate, Ammonium nitrate and chloride should not be mixed with lime. Give appropriate explanations	…
Draw up a recognition plan (order) that is most efficient in terms of time and reagent consumption	In test tubes No. 4–6, determine crystalline sodium sulfate, ammonium chloride and sodium nitrate. Write the reaction equations. Note observed signs of reactions	...
It is best to test solution samples in test tubes No. 7 and 8 with BaCl 2 and CH 3 COOH reagents, watching the result very carefully while shaking the reaction mixture	By qualitative recognition, determine Which of the test tubes No. 7 and 8 contains the solutions? sulfuric and orthophosphoric acids. Write reaction equations	...
Make a plan for recognizing the substances Na 3 PO 4 and Ca 3 (PO 4) 2 in test tubes No. 9 and 10	Determine practically in test tubes No. 9 and 10 crystalline sodium and calcium orthophosphates	...

HNO 3 is a strong acid. Its salts - nitrates-- obtained by the action of HNO 3 on metals, oxides, hydroxides or carbonates. All nitrates are highly soluble in water.

Salts of nitric acid - nitrates - decompose irreversibly when heated, the decomposition products are determined by the cation:

• a) nitrates of metals located in the voltage series to the left of magnesium:
  • 2NaNO3 = 2NaNO2 + O2
• b) nitrates of metals located in range of voltages between magnesium And copper:
  • 4Al(NO 3) 3 = 2Al 2 O 3 + 12NO 2 + 3O 2
• c) nitrates of metals located in the voltage series to the right mercury:
  • 2AgNO3 = 2Ag + 2NO2 + O2
• G) ammonium nitrate:

NH 4 NO 3 = N 2 O + 2H 2 O

Nitrates in aqueous solutions exhibit practically no oxidizing properties, but when high temperature in the solid state, nitrates are strong oxidizing agents, for example:

Fe + 3KNO 3 + 2KOH = K 2 FeO 4 + 3KNO 2 + H 2 O - when fusing solids.

Zinc And aluminum in an alkaline solution, nitrates are reduced to NH 3:

Nitric acid salts -- nitrates-- widely used as fertilizers. Moreover, almost all nitrates are highly soluble in water, so there are extremely few of them in nature in the form of minerals; the exception is Chilean (sodium) saltpeter and Indian saltpeter ( potassium nitrate). Most nitrates are obtained artificially.

Does not react with nitric acid glass, fluoroplastic-4.

Historical information

The method of obtaining dilute nitric acid by dry distillation of saltpeter with alum and copper sulfate was apparently first described in the treatises of Jabir (Geber in Latinized translations) in the 8th century. This method, with various modifications, the most significant of which was the replacement of copper sulfate with iron sulfate, was used in European and Arab alchemy until the 17th century.

IN XVII century Glauber proposed a method for producing volatile acids by reacting their salts with concentrated sulfuric acid, including nitric acid from potassium nitrate, which made it possible to introduce concentrated nitric acid into chemical practice and study its properties. Method Glauber was used before XX century, and its only significant modification was the replacement of potassium nitrate with cheaper sodium (Chilean) nitrate.

In the time of M.V. Lomonosov, nitric acid was called strong vodka. Industrial production, application and effect on the body

Nitric acid production

Nitric acid is one of the largest volume products chemical industry.

Nitric acid production

The modern method of its production is based on the catalytic oxidation of synthetic ammonia on platinum-rhodium catalysts(process Ostwald) to the mixture oxides nitrogen(nitrous gases), with their further absorption water

• 4NH 3 + 5O2(Pt) > 4 NO + 6H2O
• 2NO + O2 > 2NO 2
• 4NO 2 + O2 + 2H2O> 4HNO 3 .

Concentration The amount of nitric acid obtained by this method varies depending on the technological design of the process from 45 to 58%. Alchemists were the first to obtain nitric acid by heating a mixture of saltpeter and iron sulfate:

4KNO 3 + 2(FeSO 4 · 7H 2 O)(t°) > Fe2O3 + 2K2SO4+2HNO3^+ NO 2^ + 13H2O

Pure nitric acid was first obtained by Johann Rudolf Glauber by treating nitrate with concentrated sulfuric acid:

KNO 3 + H2SO4(conc.) (t°) > KHSO 4+HNO3^

By further distillation the so-called “fuming nitric acid”, containing virtually no water.

With oxidation states +1, +2, +3, +4, +5.

The oxides N20 and N0 are non-salt-forming (what does this mean?), and the remaining oxides are acidic: N2O3 corresponds to nitrous acid HN02, and N205 corresponds to nitric acid HNO3. Nitrogen oxide (IV) NO2, when dissolved in water, simultaneously forms two acids - HNO2 and HNO3.

If it dissolves in water in the presence of excess oxygen, only nitric acid is obtained

4N02 + 02 + 2H20 = 4HNO3

Nitrogen oxide (IV) NO2 is a brown, very poisonous gas. It is easily obtained by the oxidation of colorless, non-salt-forming nitric oxide (N) by atmospheric oxygen:

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Oxides. Nitrogen forms five oxides with oxidation states +1, +2, +3, +4, +5.

The oxides N 2 O and NO are non-salt-forming (what does this mean?), and the remaining oxides are acidic: corresponds to nitrous acid, a - nitric acid. Nitrogen oxide (IV), when dissolved in water, simultaneously forms two acids - HNO 2 and HNO 3:

2NO 2 + H 2 O = HNO 2 + HNO 3.

If it is dissolved in water in the presence of excess oxygen, only nitric acid is obtained:

4NO 2 + O 2 + 2H 2 O = 4HNO 3.

Nitrogen oxide (IV) NO 2 is a brown, very poisonous gas. It is easily obtained by the oxidation of colorless, non-salt-forming nitrogen oxide (II) with air oxygen:

2NO + O 2 = 2NO 2.

Nitric acid HNO 3. It is a colorless liquid that “smoke” in air. When stored in the light, concentrated nitric acid turns yellow, as it partially decomposes to form brown gas NO 2:

4HNO 3 = 2H 2 O + 4NO 2 + O 2.

Nitric acid exhibits all the typical properties of strong acids: it interacts with metal oxides and hydroxides, with salts (make up the appropriate reaction equations).

Laboratory experiment No. 32
Properties of dilute nitric acid

Carry out experiments to prove that nitric acid exhibits the typical properties of acids. Place a little black powder or one granule of copper (II) oxide in a test tube, pour 1-2 ml of nitric acid solution into it. Secure the test tube in the holder and heat it on the flame of an alcohol lamp. What are you observing? Write the molecular and ionic equations. Pour 1-2 ml of alkali solution into a test tube, add 2-3 drops of phenolphthalein solution. What are you observing? Add nitric acid solution to the contents of the test tube until the color disappears. What is this reaction called? Write down its molecular and ionic equations. Pour 1 ml of copper sulfate solution into a test tube, add 1-2 ml of alkali solution. What are you observing? Add a solution of nitric acid to the contents of the test tube until the precipitate disappears. Write down the molecular and ionic equations of the reactions performed.

Nitric acid behaves in a special way with metals - none of the metals displaces hydrogen from nitric acid, regardless of its concentration (for sulfuric acid this behavior is characteristic only in its concentrated state). This is explained by the fact that HNO 3 is a strong oxidizing agent; in it, nitrogen has a maximum oxidation state of +5. It is this that will be restored when interacting with metals.

The reduction product depends on the position of the metal in the stress series, on the acid concentration and on the reaction conditions. For example, when reacting with copper, concentrated nitric acid is reduced to nitric oxide (IV):

Laboratory experiment No. 33
Reaction of concentrated nitric acid with copper

Carefully pour 1 ml of concentrated nitric acid into the test tube. Using the tip of a glass tube, scoop up a little copper powder and pour it into a test tube with acid. (If there is no copper powder in your office, you can use a small piece of very thin copper wire, which must first be rolled into a ball.) What do you observe? Why does the reaction occur without heating? Why does this experiment not require the use of a fume hood? If the area of ​​contact between copper and nitric acid is less than the proposed experimental option, then what conditions must be observed? After the experiment, immediately place the test tubes with their contents in a fume hood. Write down the reaction equation and consider redox processes.

Iron and aluminum, when exposed to concentrated HNO 2, are covered with a durable oxide film, which protects the metal from further oxidation, i.e. the acid passivates the metals. Therefore, nitric acid, like sulfuric acid, can be transported in steel and aluminum tanks.

Nitric acid oxidizes many organic matter, decolorizes dyes. This usually releases a lot of heat and the substance ignites. So, if a drop of turpentine is added to nitric acid, a bright flash occurs, and a smoldering splinter in the nitric acid lights up (Fig. 135).

Rice. 135.
Burning a splinter in nitric acid

Nitric acid is widely used in the chemical industry for the production of nitrogen fertilizers, plastics, artificial fibers, organic dyes and varnishes, medicinal and explosives (Fig. 136).

Rice. 136.
Nitric acid is used to produce:
1 - fertilizers; 2 - plastics; 3 - medicines; 4 - varnishes; 5 - artificial fibers; 6 - explosives

Nitric acid salts - nitrates are obtained by the action of acid on metals, their oxides and hydroxides. Sodium, potassium, calcium and ammonium nitrates are called nitrates: NaNO 3 - sodium nitrate, KNO 3 - potassium nitrate, Ca(NO 3) 2 - calcium nitrate, NH 4 NO 3 - ammonium nitrate. Nitrate is used as nitrogen fertilizer.

Potassium nitrate is also used in the manufacture of black gunpowder, and ammonium nitrate, as you already know, is used to prepare the explosive ammonal. Silver nitrate, or lapis, AgNO 3 is used in medicine as a cauterizing agent.

Almost all nitrates are highly soluble in water. When heated, they decompose releasing oxygen, for example:

New words and concepts

1. Non-salt-forming and acidic nitrogen oxides.
2. Nitric oxide (IV).
3. Properties of nitric acid as an electrolyte and as an oxidizing agent.
4. Interaction of concentrated and dilute nitric acid with copper.
5. Application of nitric acid.
6. Nitrates, nitrate.

Tasks for independent work

1. Why does nitric acid not form acid salts?
2. Write molecular and ionic equations for the reactions of nitric acid with copper (II) hydroxide, iron (III) oxide and sodium carbonate.
3. Most nitric acid salts are soluble in water, however, propose an equation for the reaction of HNO 3 with the salt, resulting in the formation of a precipitate. Write the ionic equation for this reaction.
4. Consider the equations for the reactions of dilute and concentrated nitric acid with copper from the point of view of oxidation-reduction processes.
5. Propose two chains of transformations leading to the production of nitric acid, starting from nitrogen and ammonia. Describe redox reactions using the electron balance method.
6. How many kilograms of 68% nitric acid can be obtained from 276 kg (N.S.) of nitric oxide (IV)?
7. When calcining 340 g of sodium nitrate, 33.6 liters of oxygen were obtained. Calculate mass fraction impurities in saltpeter.