Chemical Reactions

Classifying and Balancing Chemical Reactions

Reactants: Substances that undergo reaction. These are placed on the left side of a chemical equation. 

Products: Substances that are produced through a chemical reaction. These are placed on the right side of a chemical equation. 

There are three common types of chemical reactions:

• Precipitation
•  Acid-base 
• Oxidation-reduction

Precipitation Reactions and Solubility Rules

Solubility: a substance's ability to dissolve in water and produce a solution. Based on its solubility numbers, a substance can be classified as either soluble, slightly soluble, or insoluble. 

In a precipitation reaction, dissolved substances react together to produce a solid product, which is called the precipitate. 

Example I: 

When solutions of potassium iodide and lead nitrate are mixed, lead iodide and potassium nitrate are formed. 

2KI(aq) + Pb(NO₃)₂(aq) ⟶ PbI₂(s) + 2KNO₃(aq)

Using the solubility table below, we can determine that when iodide ions form a compound with a lead ion, it is insoluble. Therefore, it forms a precipitate (solid).

Example II:

Predict the result of mixing solutions of  potassium sulfate and barium nitrate.

The two products for this reaction are KNO₃ and BaSO₄. In the solubility table, it’s shown that BaSO₄ is insoluble; therefore, a precipitation reaction is expected. The ionic equation for this precipitation reaction becomes:

Ba²⁺_(aq) + SO₄²⁻_(aq) ⟶ BaSO_4(s)

With that, the final net ionic equation for the reaction becomes:

Ba(NO₃)₂ + K₂SO₄ → BaSO₄ + 2KNO₃

Table 1: Solubility Table

Soluble Ionic Compounds
Contains these ions	Exceptions
Group I cations: Li+ Na+ K+ Rb+ Cs+	None
Cl- Br- I-	Compounds with Ag+, Hg22+, and Pb2+
F-	Compounds with group 2 metal cations, Pb2+ and Fe3+
C2H3O2- HCO3- NO3- ClO3-	none
SO42-	Compounds with Ag+, Ba2+, Ca2+, Hg22+, Pb2+, and Sr2+
Insoluble Ionic Compounds
Contains these ions	Exceptions
CO32- CrO42- PO43- S2-	Compounds with group 1 cations and NH4+
OH-	Compounds with group 1 cations and Ba2+

 

Acid-Base Reactions

In this section, we will focus only on the common types of acid-base reactions that happen in aqueous solutions.

An acid is a substance that dissolves in water to produce hydronium ions.

When an acid is dissolved in water, a hydrogen ion, H⁺, is transferred from one chemical species to the other. A base, on the other hand, yields hydroxide ions, OH⁻, when dissolved in water.

Strong acids and bases are acids and bases that completely react with water and ionize when dissolved in water.

Example of a strong acid: when hydrochloric acid is dissolved in water, H₃O⁺ ions are produced, where H⁺ ions are transferred from HCl molecules to H₂O molecules, as shown below:

HCl_(aq) + H₂O_(aq) ⟶ Cl⁻_(aq) + H₃O⁺_(aq)

Example of a strong base: when sodium hydroxide is dissolved in water, it produces the following products:

NaOH_(s) ⟶ Na⁺_(aq) + OH⁻_(aq)

Weak acids and bases: weak acids and bases do not completely dissociate in water and only partially react with water.

Example of a weak acid:

CH₃CO₂H_(aq) + H₂O_(l) ⇌ CH₃CO₂⁻_(aq) + H₃O⁺_(aq)

Example of a weak base:

NH_3(aq) + H₂O_(l) ⇌ NH₄⁺_(aq) + OH⁻_(aq)

Oxidation-Reduction Reactions

An oxidation-reduction reaction, also known as a redox reaction, is where electrons transfer from one specie to the other. In other words, one specie gives electrons while the other accepts them. The loss of electrons is known as oxidation half-reaction, and the gain of electrons is known as reduction half-reaction.

Oxidation = loss of electrons (oxidation number of species increases)

Reduction = gain of electrons (oxidation number of species decreases) 

Oxidizing agent = species that are reduced

Reducing agent = species that are oxidized

In a redox reaction, the charge of the atoms involved is known as the oxidation number. For example, Na⁺ has the oxidation number of 1, and O^2- has the oxidation number of 2. An elemental substance (what you find in the periodic table) has an oxidation number of 0. The oxidation number of a monatomic ion is equal to its charge. In polyatomic ions, the sum of all the oxidation numbers of each ion is equal to its charge.

Oxidation numbers for common nonmetals are assigned as follows:

• Hydrogen: +1 when combined with nonmetals, −1 when bonded with metals.
• Oxygen: −2 in most compounds, sometimes −1 (peroxides, O2²⁻), rarely −1/2 (superoxides, O2⁻), positive values when bonded with F (values change).
• Halogens: −1 for F always, −1 for other halogens except when bonded with oxygen or other halogens (positive oxidation numbers in these cases, changing values).

Example I

Assign oxidation numbers to all the elements in the following compound:

H₂S (Hydrogen sulfide)

The oxidation number for H is +1. Knowing this, the oxidation number for sulfur can be calculated as follows:

Total charge on H₂S = 0 = (2 * +1) + (1 * x)

x = 0 − (2 * +1) = −2

Therefore, sulfur in this compound has a charge of -2.

Example II

Assign oxidation numbers to all the elements in the following polyatomic ion:

SO₃²⁻

SO₃²⁻ = −2 = (3 * −2) + (1 * x)

x = −2 − (3 * −2) = +4

Note: A chemical reaction is classified as a redox reaction if there is a change in oxidation numbers.

Examples

Identify whether the following chemical reactions are redox reactions:

1) ZnCO₃(s) ⟶ ZnO(s) + CO₂(g)

 

 

The above reaction is not a redox reaction since the oxidation numbers of the species in the reactants and the oxidation numbers of the same species in the products are the same.

2) 2Ga_(l) + 3Br_2(l) ⟶ 2GaBr_3(s)

 

 

The above reaction is a redox reaction since the oxidation numbers of the species in the reactants and the oxidation numbers of the same species in the products are not the same.