Question

Assign oxidation numbers to all the atoms in each of the following. a. \(\mathrm{SrCr}_{2} \mathrm{O}_{7} \quad\) g. \(\mathrm{PbSO}_{3}\) b. \(\mathrm{CuCl}_{2} \quad \quad\) h. \(\mathrm{PbO}_{2}\) c. \(\mathrm{O}_{2} \quad\quad\quad\) i. \(\mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4}\) d. \(\mathrm{H}_{2} \mathrm{O}_{2} \quad\quad \mathrm{j} . \mathrm{CO}_{2}\) e. \(\mathrm{MgCO}_{3} \quad\) k. \(\left(\mathrm{NH}_{4}\right)_{2} \mathrm{Ce}\left(\mathrm{SO}_{4}\right)_{3}\) f. \(\mathrm{Ag} \quad\quad\quad \)l. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\)

Short answer

a. Sr: +2, Cr: -3, O: -2 b. Cu: +2, Cl: -1 c. O: 0 d. H: +1, O: -1 e. Mg: +2, C: +4, O: -2 f. Ag: 0 g. Pb: +2, S: +6, O: -2 h. Pb: +4, O: -2 i. Na: +1, C: +3, O: -2 j. C: +4, O: -2 k. N: -3, H: +1, Ce: +3, S: +6, O: -2 l. Cr: +3, O: -2

Step-by-step solution

Rule 1: Oxidation Number of Individual Atoms

The oxidation number of an individual atom is always zero.

Rule 2: Oxidation Number of Monoatomic Ions

The oxidation number of a monoatomic ion is equal to its charge.

Rule 3: Oxidation Number of Oxygen

The oxidation number of oxygen is always -2, except in peroxides (like \(\mathrm{H}_{2}\mathrm{O}_{2}\)) where it is -1, and in compounds with fluorine where it can be positive.

Rule 4: Oxidation Number of Hydrogen

The oxidation number of hydrogen is +1, except in metal hydrides where it is -1.

Rule 5: Oxidation Number of Group 1 and 2 Metals

The oxidation number of Group 1 metals is always +1, and the oxidation number of Group 2 elements is always +2.

Rule 6: Oxidation Number of Fluorine

The oxidation number of fluorine is always -1.

Rule 7: Oxidation Number of Group 17 (Halogens)

The oxidation number of halogens is usually -1, but this can vary when they form a compound with more electronegative atoms or with other halogens.

Rule 8: Sum of Oxidation Numbers in Neutral Compounds

The sum of oxidation numbers in a neutral compound is always 0.

Rule 9: Sum of Oxidation Numbers in Polyatomic Ions

The sum of oxidation numbers in a polyatomic ion is equal to the charge of the ion. Now let's apply these rules to find the oxidation numbers of atoms in the given compounds. a. \(\mathrm{SrCr}_{2} \mathrm{O}_{7}\): - \(\mathrm{Sr}\) is a Group 2 element, so its oxidation number is +2. - \(\mathrm{Cr}_{2} \mathrm{O}_{7}\) is a polyatomic ion with a charge of -2. Since there are two Cr atoms, the total oxidation number of Cr should be -6 (using Rule 6: to cancel out the -14 from 7 O atoms with -2 oxidation number). Therefore, each Cr atom has an oxidation number of -3. - Oxygen has an oxidation number of -2. b. \(\mathrm{CuCl}_{2}\): - Cu has an oxidation number of +2 (to make the sum of oxidation numbers 0). - Cl has an oxidation number of -1. c. \(\mathrm{O}_{2}\): As O2 is an individual(atom) oxygen molecule, the oxidation number for both O atoms is 0. d. \(\mathrm{H}_{2} \mathrm{O}_{2}\): - H has an oxidation number of +1. - O has an oxidation number of -1 (since it is a peroxide). e. \(\mathrm{MgCO}_{3}\): - Mg is a Group 2 element, so its oxidation number is +2. - C in CO3 has an oxidation number of +4 (to cancel out the -6 from 3 O atoms). - Oxygen has an oxidation number of -2. f. \(\mathrm{Ag}\): As an individual Ag atom, the oxidation number is 0. g. \(\mathrm{PbSO}_{3}\): - Pb has an oxidation number of +2 (to cancel out the -6 from SO3). - S in SO3 has an oxidation number of +6 (to cancel out the -6 from 3 O atoms). - Oxygen has an oxidation number of -2. h. \(\mathrm{PbO}_{2}\): - Pb has an oxidation number of +4 (to cancel out the -4 from 2 O atoms). - Oxygen has an oxidation number of -2. i. \(\mathrm{Na}_{2} \mathrm{C}_{2}\mathrm{O}_{4}\): - Na is a Group 1 element, so its oxidation number is +1. - C in C2O4 has an oxidation number of +3 (to cancel out the -2 from 2 O atoms, the total oxidation number of C2 should be +6). - Oxygen has an oxidation number of -2. j. \(\mathrm{CO}_{2}\): - C has an oxidation number of +4 (to cancel out the -4 from 2 O atoms). - Oxygen has an oxidation number of -2. k. \(\left(\mathrm{NH}_{4}\right)_{2}\mathrm{Ce}\left(\mathrm{SO}_{4}\right)_{3}\): - N has an oxidation number of -3. - H has an oxidation number of +1. - NH4 has an oxidation number of +1 (from N: -3 and H: 1*4 = +4, sum = +1). - Ce has an oxidation number of +3 (to make the sum of oxidation numbers 0). - S in SO4 has an oxidation number of +6 (to cancel out the -8 from 4 O atoms). - Oxygen has an oxidation number of -2. l. \(\mathrm{Cr}_{2} \mathrm{O}_{3}\): - Cr has an oxidation number of +3 (to cancel out the -6 from 3 O atoms). - Oxygen has an oxidation number of -2.

Key concepts

Redox Reactions

Redox reactions, short for reduction-oxidation reactions, are processes where electrons are transferred between substances. One substance gets oxidized, meaning it loses electrons, while another gets reduced, meaning it gains electrons. Understanding redox reactions is crucial because they are at the heart of many important biological and chemical processes, such as photosynthesis, respiration, and combustion.
When looking at redox reactions, it's important to focus on the changes in oxidation numbers. This tells us which species are being oxidized and reduced:

• Oxidation: This is when the oxidation number increases. This means the substance loses electrons.
• Reduction: This is when the oxidation number decreases. It tells us the substance gains electrons.

To identify the oxidizing and reducing agents, remember:

• The substance that gets reduced (gains electrons) is the oxidizing agent.
• The substance that gets oxidized (loses electrons) is the reducing agent.

By practicing with equations and keeping track of oxidation numbers, one can easily identify these processes in chemical reactions.

Chemical Bonding

Chemical bonding is the force of attraction between atoms or ions that enables the formation of chemical compounds. This is fundamental to everything in chemistry as it determines the structure and properties of substances. There are several types of chemical bonding, each with unique characteristics:

• Ionic Bonding: Occurs between metals and non-metals, involving the transfer of electrons from one atom to another, creating ions.
• Covalent Bonding: Happens between non-metal atoms, where they share pairs of electrons to achieve stability.
• Metallic Bonding: Found between metal atoms, where electrons are delocalized, allowing metals to conduct electricity and heat.

Chemical bonds affect the oxidation numbers of atoms in a compound. For instance, in ionic compounds like \( \mathrm{CuCl}_2 \), the copper (Cu) has an oxidation number that reflects its ionic state (^+\(2\)), indicating its charge after losing two electrons to the chlorine atoms.

Oxidation States

Oxidation states, often referred to as oxidation numbers, are essential in determining the electron distribution in a molecule or ion. They provide insight into the degree of oxidation or reduction a particular atom has undergone in a compound. Here are a few rules to keep in mind when assigning oxidation states:

• An individual atom not combined with others has an oxidation number of 0.
• For monoatomic ions, the oxidation number is equal to the charge of the ion. For instance, in \( \mathrm{Ni}^{2+} \), the oxidation number is +2.
• Oxygen generally has an oxidation number of -2 in compounds, except in peroxides (like \( \mathrm{H}_2\mathrm{O}_2 \)) where it's -1.
• Hydrogen usually has an oxidation number of +1, except when bonded with metals in hydrides, where it's -1.
• In a neutral molecule, the sum of all oxidation numbers must be zero, whereas, in a polyatomic ion, it must equal the ion's charge.

For example, in \( \mathrm{SrCr}_2\mathrm{O}_7 \), one can determine the oxidation number of each element by applying these rules, ensuring the overall neutrality of the compound.