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Redox Reactions: Exercises in Oxidation-Reduction Chemistry, Exercises of Chemistry

Chemistry

A comprehensive set of exercises focused on oxidation-reduction (redox) reactions, designed for high school students. It includes problems on determining oxidation states, identifying oxidized and reduced elements, and balancing half-reactions in both acidic and basic conditions. The exercises cover a range of chemical species and reactions, offering practical application of redox principles. The document also includes notes on oxidation state rules and terminologies, enhancing understanding of redox processes. It serves as a valuable resource for students to practice and master redox reaction concepts, improving their problem-solving skills in chemistry. Structured to facilitate step-by-step learning, making it an effective tool for both classroom and self-study.

Typology: Exercises

2024/2025

Uploaded on 08/20/2025

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shanjida-raihan 🇺🇸

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SAC Chemistry Department Page 1 of 7
Learning
Activity
C
Oxidation-Reduction
Reactions v1.2025
Name
Student ID#
Instructor
Due Date
Reading: Chemistry 2nd ed., OpenStax section 4.2 with focus on redox reaction.
1. Determine the oxidation state for each element in each formula below
Ex: Na2C2O4 Na: +1, O: -2, C: +3
KNO3 ________________________________ H2SO4 ________________________________
Cr2O72- ________________________________ H2O2 ________________________________
O3 ________________________________ KMnO4________________________________
CaH2 ________________________________ H2CO ________________________________
C6H5OH ________________________________ CuClO3________________________________
2. For each of the following chemical changes, first identify the element that changes oxidation state, second
indicate its oxidation state before and after, third indicate if it is oxidized or reduced, and then indicate if the
element has gained or lost e- and how many per atom. If there is no change in oxidation state of any element,
state so.
Ex: NO3- → NO N is going from +5 to +2; it is reduced; each N atom gains 3 e-
MnO41- → Mn(OH)3 __________________________________________________________________
NH4+ → NH3 __________________________________________________________________
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Learning Activity

C

Oxidation-Reduction

Reactions v1.

Name

Student ID#

Instructor

Due Date

Reading: Chemistry 2nd^ ed., OpenStax section 4.2 with focus on redox reaction.

  1. Determine the oxidation state for each element in each formula below Ex: Na 2 C 2 O 4 Na: +1, O: - 2, C: + KNO 3 ________________________________ H 2 SO 4 ________________________________ Cr 2 O 72 -^ ________________________________ H 2 O 2 ________________________________ O 3 ________________________________ KMnO 4 ________________________________ CaH 2 ________________________________ H 2 CO ________________________________ C 6 H 5 OH ________________________________ CuClO 3 ________________________________
  2. For each of the following chemical changes, first identify the element that changes oxidation state, second indicate its oxidation state before and after, third indicate if it is oxidized or reduced, and then indicate if the element has gained or lost e- and how many per atom. If there is no change in oxidation state of any element, state so. Ex: NO 3 -^ → NO N is going from +5 to +2; it is reduced; each N atom gains 3 e- MnO 41 -^ → Mn(OH) 3 __________________________________________________________________ NH 4 +^ → NH 3 __________________________________________________________________

S 22 -^ → S __________________________________________________________________

HCHO → HCHO 2 __________________________________________________________________

NaCl → NaClO __________________________________________________________________

  1. For each of the following half-reactions, circle oxidation (O) or reduction (R) then balance it using half- reaction method. For the half-reactions that are in basic condition, write the balanced half in acidic condition on the line and the balanced half in basic condition underneath. a. Cl-^ to ClO 3 -^ (acid) O R _______________________________________________________ b. H 2 S to H 2 SO 3 (acid) O R _______________________________________________________ c. S 2 O 3 -^2 to H 2 S (acid) O R _______________________________________________________ d. HNO 2 to NH 4

(acid) O R _______________________________________________________ e. Cr 2 O 72 -^ to Cr3+^ (basic) O R _______________________________________________________ f. ClO 21 -^ to Cl^1 -^ (acid) O R _______________________________________________________

b) H 2 O 2 à O 2 + H 2 O (in acidic solution) c) CN^1 -^ + MnO 41 -^ à CNO^1 -^ + MnO 2 (in basic solution)

d) Br 2 à BrO 31 -^ + Br^1 -^ (in basic solution) e) ClO 31 -^ + Cl^1 -^ à Cl 2 + ClO 2 (in acidic solution)

NOTES FOR REDOX REACTIONS v7.24.

  1. Rules for oxidation state: a) Oxidation state (like charge) is per 1 atom b) Elements have oxidation state = 0 in its free element form (H 2 , O 2 , Mg, Al…) c) Hydrogen has ox. state = +1, except when it’s in a binary compound with a metal where it’s - 1. d) Oxygen has oxidation state = - 2, except in peroxides (Na 2 O 2 , H 2 O 2 …) where it’s - 1. e) The sum all of all oxidation states in a species must be equal to its net charge f) For ions, oxidation state equals charge.
  2. The oxidation state range for an atom a) Each atom has a range of oxidation state. b) The highest oxidation state of an atom is equal to its positive charge when all of its valence e- are lost. If an atom is at its highest oxidation state, it can only decrease its oxidation state (be reduced) c) The lowest oxidation state of an atom is equal to its negative charge when it has gained all e- needed to become next noble gas. If an atom is at its lowest oxidation state, it can only increase its ox. state (be oxidized). For example, S has lowest oxidation state of - 2 and highest oxidation state of +6. d) Non-metals can have positive and negative oxidation states. e) Metals only have positive (or zero) oxidation states, but not negative.
  3. Some terminologies: a) Oxidation: increase in oxidation state; also means loss of e- b) Reduction: decrease in oxidation state; also means gain of e- c) Oxidizing agent: a species (chemical) that can oxidize something else; it itself is reduced (its oxidation state decreases aka it gains electron(s)) d) Reducing agent: a species (chemical) that can reduce something else; it itself is oxidized (its oxidation state increases aka it loses electron(s)) e) In a redox reaction, there is always a reagent (chemical) that has its oxidation state increased and another reagent that has its oxidation state decreased. A reagent can undergo both oxidation and reduction in a redox reaction.
  4. Steps to do to balance a half-reaction (oxidation or reduction); must do these steps in order below i. Balance atoms that are not H and O first using coefficients ii. Balance O by adding H 2 O to the side that needs O iii. Balance H by adding H+^ to the side that needs H (all atoms now must be balanced) iv. Balance charge by adding e- (not subtracting) to the correct side so charges on both sides are equal. For an oxidation half, e- needs to show up on the right; for reduction half, e- needs to show up on the left. If this is not the case, something is wrong. v. If in basic condition, add – OH to both sides of the equation; the number of – OH added must be equal to the coefficient in front of H+. Combine H+^ and – OH to give H 2 O. vi. Combine H 2 O on both sides
  5. Steps to do to balance a full redox reaction; must do these steps in order below i. Complete steps 4(i-iv) above for each half reaction. ii. Find the least common denominator for the # of e- of both half reactions. iii. Multiply each half reaction so # of e- in each half is the same iv. Add the 2 half-reactions (there can be no e- left); combine like terms (check to make sure all atoms are balanced and charges are balanced) v. Convert from acidic condition to basic condition by adding – OH and combine terms if needed.