Practical Work 2. Studying the Properties of Acids, Bases, and Salts in Light of the Theory of Electrolytic Dissociation

Lecture

I. Studying the properties of acids

1. Determine the pH of the given hydrochloric acid solution.

2. Investigate which of the proposed substances hydrochloric acid will react with: magnesium, copper, potassium hydroxide, sodium carbonate, sodium sulfate.

Recall that when carrying out a neutralization reaction, an indicator (phenolphthalein or another) can first be added to the alkali, followed by the acid.

Note the qualitative signs of the reactions taking place.

Draw conclusions about the properties of acids using hydrochloric acid as an example.

II. Studying the properties of alkalis

1. Determine the pH of the given sodium hydroxide solution.

2. Investigate how sodium hydroxide behaves toward acids and salts, for example sulfuric acid and iron(III) chloride.

Note the qualitative signs of the reactions taking place.

Draw conclusions about the properties of alkalis using sodium hydroxide as an example.

III. Studying the properties of salts

Precipitate the metal cation and the acid-residue anion contained in a solution of iron(II) sulfate.

Draw a conclusion about the properties of salts using iron(II) sulfate as an example.

In the report on the work performed, present all reaction equations in molecular, full ionic, and net ionic forms.

I. Studying the properties of acids

Practical Work 2. Studying the Properties of Acids, Bases, and Salts in Light of the Theory of Electrolytic Dissociation

1. Determination of the pH of hydrochloric acid

Hydrochloric acid solution — a strong acid
Solution pH: ≈ 1–2

Sign:

Indicator (litmus) → red
phenolphthalein → colorless

2. Reactions of hydrochloric acid

Testing substances:

1. Magnesium (Mg) reacts

Equation:

$M g + 2 H C l \to M g C l_{2} + H_{2} ↑$

Signs:

gas (H₂) is evolved
bubbles
the metal dissolves

2. Copper (Cu) does not react

Reason:

copper stands after hydrogen in the activity series

Signs:

no reaction

3. Potassium hydroxide (KOH) reacts (neutralization)

Equation:

$H C l + K O H \to K C l + H_{2} O$

Signs:

the crimson color of phenolphthalein disappears
heat is released

4. Sodium carbonate (Na₂CO₃) reacts

Equation:

$N a_{2} C O_{3} + 2 H C l \to 2 N a C l + H_{2} O + C O_{2} ↑$

Signs:

gas (CO₂) is evolved
fizzing

5. Sodium sulfate (Na₂SO₄) does not react

Reason:

a salt of a strong acid and a strong base

Signs:

no changes

Summary (schematic)

Substance	Reaction	Sign
Mg	+	gas H₂
Cu	-	no reaction
KOH	+	color disappears
Na₂CO₃	+	CO₂, fizzing
Na₂SO₄	-	no reaction

Conclusion

Using hydrochloric acid as an example:

Acids:

have a low pH
react with:
- metals (up to H₂ in the activity series)
- alkalis (neutralization)
- carbonates (with evolution of CO₂)
do not react with:
- low-activity metals (Cu)
- certain salts

II. Studying the properties of alkalis

1. Determination of the pH of a sodium hydroxide solution

Sodium hydroxide (NaOH) is a strong alkali.

Solution pH: ≈ 13–14
Indicators:
- litmus → blue
- phenolphthalein → crimson
- universal indicator → dark blue/purple

Conclusion: the medium is strongly alkaline

2. Interaction of NaOH with substances

a) With an acid (sulfuric acid H₂SO₄)

Molecular equation:

$2 N a O H + H_{2} S O_{4} \to N a_{2} S O_{4} + 2 H_{2} O$

Full ionic:

$2 N a^{+} + 2 O H^{-} + 2 H^{+} + S O_{4}^{2 -} \to 2 N a^{+} + S O_{4}^{2 -} + 2 H_{2} O$

Net ionic:

$H^{+} + O H^{-} \to H_{2} O$

Signs of the reaction:

heat is released
the indicator changes color (neutralization)

b) With a salt (iron(III) chloride FeCl₃)

Molecular equation:

$F e C l_{3} + 3 N a O H \to F e (O H)_{3} ↓ + 3 N a C l$

Full ionic:

$F e^{3 +} + 3 C l^{-} + 3 N a^{+} + 3 O H^{-} \to F e (O H)_{3} ↓ + 3 N a^{+} + 3 C l^{-}$

Net ionic:

$F e^{3 +} + 3 O H^{-} \to F e (O H)_{3} ↓$

Signs of the reaction:

a brown precipitate of iron(III) hydroxide forms

Conclusion

Alkalis (using NaOH as an example):

have a high pH (an alkaline medium)
react with acids (neutralization reaction)
react with salts, forming insoluble hydroxides (precipitates)
change the color of indicators

Sodium hydroxide is a typical strong alkali, displaying all the characteristic chemical properties of alkalis.

Practical Work 2. Studying the Properties of Acids, Bases, and Salts in Light of the Theory of Electrolytic Dissociation

Studying the properties of salts (using FeSO₄ as an example)

1. Precipitation of the metal cation (Fe²⁺)

Add an alkali solution (for example, NaOH).

Observation:

A light green precipitate of iron(II) hydroxide forms, which turns brown over time in air.

Molecular equation:

FeSO₄ + 2NaOH → Fe(OH)₂↓ + Na₂SO₄

Full ionic equation:

Fe²⁺ + SO₄²⁻ + 2Na⁺ + 2OH⁻ → Fe(OH)₂↓ + 2Na⁺ + SO₄²⁻

Net ionic equation:

Fe²⁺ + 2OH⁻ → Fe(OH)₂↓

2. Precipitation of the acid-residue anion (SO₄²⁻)

Add a barium salt solution (for example, BaCl₂).

Observation:

A white precipitate of barium sulfate forms.

Molecular equation:

FeSO₄ + BaCl₂ → BaSO₄↓ + FeCl₂

Full ionic equation:

Fe²⁺ + SO₄²⁻ + Ba²⁺ + 2Cl⁻ → BaSO₄↓ + Fe²⁺ + 2Cl⁻

Net ionic equation:

Ba²⁺ + SO₄²⁻ → BaSO₄↓

Conclusion

Iron(II) sulfate displays the typical properties of salts:

in solution it dissociates into ions (Fe²⁺ and SO₄²⁻);
it undergoes exchange reactions;
it can form precipitates:
- by cation (Fe²⁺ → Fe(OH)₂↓),
- by anion (SO₄²⁻ → BaSO₄↓).

Salts are electrolytes that break down into ions in aqueous solutions and undergo ion exchange reactions with the formation of precipitates, gases, or weak electrolytes.

Practical Work 2. Studying the Properties of Acids, Bases, and Salts in Light of the Theory of Electrolytic Dissociation