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Questions and Answers
The oxidation number of an elemental substance is always +1.
The oxidation number of an elemental substance is always +1.
False
In redox reactions, oxidation refers to the gain of electrons.
In redox reactions, oxidation refers to the gain of electrons.
False
The oxidation number of oxygen is -1 in peroxides.
The oxidation number of oxygen is -1 in peroxides.
True
The sum of the oxidation numbers of all atoms in a neutral compound equals zero.
The sum of the oxidation numbers of all atoms in a neutral compound equals zero.
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In the reaction 2Mg + O2 → 2MgO, magnesium is the reducing agent.
In the reaction 2Mg + O2 → 2MgO, magnesium is the reducing agent.
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The oxidation state of fluorine in all its compounds is +1.
The oxidation state of fluorine in all its compounds is +1.
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Potassium permanganate (KMnO4) acts as a reducing agent in reactions.
Potassium permanganate (KMnO4) acts as a reducing agent in reactions.
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In acidic solutions, electrons must be added to balance the charge in half-reactions.
In acidic solutions, electrons must be added to balance the charge in half-reactions.
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The oxidation number of hydrogen is +1 when combined with metals.
The oxidation number of hydrogen is +1 when combined with metals.
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Superoxides exhibit oxygen with a -1/2 oxidation state.
Superoxides exhibit oxygen with a -1/2 oxidation state.
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The balanced redox reaction involving zinc metal and nitrate ions produces zinc tetrahydroxide and ammonia.
The balanced redox reaction involving zinc metal and nitrate ions produces zinc tetrahydroxide and ammonia.
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When balancing the redox reaction, the total number of electrons lost must exceed the total number of electrons gained.
When balancing the redox reaction, the total number of electrons lost must exceed the total number of electrons gained.
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In acidic solutions, H+ ions can be used to balance the charge in a redox reaction.
In acidic solutions, H+ ions can be used to balance the charge in a redox reaction.
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The equation 2MnO4– + 6H+ + 5H2C2O4 → 2Mn2+ + 10CO2 + 8H2O is an unbalanced equation.
The equation 2MnO4– + 6H+ + 5H2C2O4 → 2Mn2+ + 10CO2 + 8H2O is an unbalanced equation.
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In the activity series, zinc can reduce copper ions but is unable to reduce hydrogen ions.
In the activity series, zinc can reduce copper ions but is unable to reduce hydrogen ions.
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When balancing the equation for Zn + NO3– → Zn(OH)42– + NH3, the product side contains more hydrogen atoms than the reactant side.
When balancing the equation for Zn + NO3– → Zn(OH)42– + NH3, the product side contains more hydrogen atoms than the reactant side.
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The process of combining H+ and OH– to form water is necessary when balancing half-reactions in basic solution.
The process of combining H+ and OH– to form water is necessary when balancing half-reactions in basic solution.
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The half-reaction for the reduction of MnO4– to Mn2+ involves the gain of electrons.
The half-reaction for the reduction of MnO4– to Mn2+ involves the gain of electrons.
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The overall charge is only important to check in acidic solutions when balancing redox reactions.
The overall charge is only important to check in acidic solutions when balancing redox reactions.
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The balanced reaction 4Zn + 7OH– + NO3– + 6H2O → 4Zn(OH)42– + NH3 contains an equal number of oxygen atoms on both sides.
The balanced reaction 4Zn + 7OH– + NO3– + 6H2O → 4Zn(OH)42– + NH3 contains an equal number of oxygen atoms on both sides.
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Study Notes
Oxidation Numbers
- The effective charge of an atom.
- The oxidation number of an elemental substance is zero.
- The oxidation number of a monoatomic ion equals the charge number of that ion.
- The sum of the oxidation numbers of all atoms in the species is equal to its total charge.
- Atoms in their elemental form are 0.
- Group I elements have +1 oxidation number.
- Group II elements have +2 oxidation number.
- Group III elements (excluding boron) have +3 oxidation number for M3+ and +1 for M+.
- Group IV elements (excluding carbon and silicon) have +4 oxidation number for M4+ and +2 oxidation number for M2+.
- Hydrogen has +1 oxidation number in non-metal combinations and -1 in metal combinations.
- Fluorine has -1 oxidation number in all its compounds.
- Oxygen has -2 oxidation number unless combined with fluorine, -1 in peroxides (O22-), -½ in superoxides (O2-), and -⅓ in ozonides (O3-).
Redox Reactions
- Oxidation is electron loss.
- Reduction is electron gain.
- A redox reaction is a reaction that involves both oxidation and reduction.
- An oxidizing agent removes electrons and becomes reduced in a reaction.
- An element in the oxidizing agent undergoes a decrease in oxidation number.
- A reducing agent supplies electrons and becomes oxidized in a reaction.
- An element in the reducing agent undergoes an increase in oxidation number.
Redox Reactions in Acidic Solution
- Redox reactions can be balanced in acidic solutions
- Steps for balancing redox reactions in acidic solutions:
- Identify the oxidized and reduced species based on oxidation number changes.
- Write two skeletal half-reactions.
- Balance all elements in the half-reactions by inspection, except oxygen, hydrogen, and charge.
- Balance oxygen atoms in each half-reaction by adding H2O.
- Balance hydrogen atoms in each half-reaction by adding H+.
- Balance electric charges in each half-reaction by adding electrons.
- Make the number of electrons in both half-reactions equal by multiplying one or both reactions by a factor.
- Combine the two half-reactions by adding them together and simplify.
- Check the equation is balanced by counting the number of each atom and overall charge on both sides of the equation.
- Example: 2MnO4– + 6H+ + 5H2C2O4 → 2Mn2+ + 10CO2 + 8H2O
Redox Reactions in Basic Solution
- Redox reactions can be balanced in basic solutions.
- Steps for balancing redox reactions in basic solutions:
- Follow the steps for balancing redox reactions in acidic solutions but add the following steps:
- After balancing hydrogen by adding H+, add the same number of OH- ions to both sides of each half-reaction as there are H+.
- Combine H+ and OH- in each half-reaction to form H2O.
- Remove excess H2O from each side of the half-reaction.
- Continue with the remaining steps for balancing redox reactions in acidic solutions.
- Follow the steps for balancing redox reactions in acidic solutions but add the following steps:
- Example: 2MnO4– + H2O + Br– → 2MnO2 + 2OH– + BrO3–
Element Activity Series
- A list of elements arranged in order of their reactivity, where a metal can reduce the cations formed by any of the metals below it in the list.
- For example, Zinc can reduce Copper ions, but Copper cannot reduce Zinc ions.
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Description
This quiz covers essential concepts related to oxidation numbers and redox reactions. It explains the effective charge of atoms, specific oxidation states for various elements, and the definitions of oxidation (electron loss) and reduction (electron gain). Test your understanding of these fundamental chemistry topics!
