and are key processes in electrochemistry. They involve the transfer of electrons between species, changing their oxidation states. Understanding these concepts is crucial for grasping how batteries work and how chemical reactions occur.
Oxidizing and reducing agents play vital roles in these electron transfers. Oxidizing agents accept electrons, while reducing agents donate them. Knowing how to identify and balance these reactions is essential for predicting and controlling chemical processes in various applications.
Oxidation and Reduction Fundamentals
Oxidation and reduction definitions
Top images from around the web for Oxidation and reduction definitions
electrochemistry - Anode and Cathode Understanding - Chemistry Stack Exchange View original
Is this image relevant?
Electrochemistry: cells and electrodes View original
Is this image relevant?
Galvanic Cells | Chemistry: Atoms First View original
Is this image relevant?
electrochemistry - Anode and Cathode Understanding - Chemistry Stack Exchange View original
Is this image relevant?
Electrochemistry: cells and electrodes View original
Is this image relevant?
1 of 3
Top images from around the web for Oxidation and reduction definitions
electrochemistry - Anode and Cathode Understanding - Chemistry Stack Exchange View original
Is this image relevant?
Electrochemistry: cells and electrodes View original
Is this image relevant?
Galvanic Cells | Chemistry: Atoms First View original
Is this image relevant?
electrochemistry - Anode and Cathode Understanding - Chemistry Stack Exchange View original
Is this image relevant?
Electrochemistry: cells and electrodes View original
Is this image relevant?
1 of 3
Oxidation involves the loss of electrons from a species resulting in an increase in its
Occurs at the in an electrochemical cell (battery)
Examples: Fe2+→Fe3++e−, 2Cl−→Cl2+2e−
Reduction involves the gain of electrons by a species resulting in a decrease in its oxidation number
Occurs at the in an electrochemical cell
Examples: Cu2++2e−→Cu, 2H++2e−→H2
Mnemonic device OIL RIG helps remember Oxidation Is Loss of electrons and Reduction Is Gain of electrons
Oxidizing and reducing agents
Oxidizing agent (oxidant) is a species that accepts electrons causing the oxidation of another species and is reduced in the process
Examples: H2O2 in the reaction 2Fe2++H2O2+2H+→2Fe3++2H2O, MnO4− in the reaction 5Fe2++MnO4−+8H+→5Fe3++Mn2++4H2O
Reducing agent (reductant) is a species that donates electrons causing the reduction of another species and is oxidized in the process
Examples: Na in the reaction 2Na+Cl2→2NaCl, Sn2+ in the reaction Sn2++2Fe3+→Sn4++2Fe2+
Oxidation numbers in compounds
Rules for assigning oxidation numbers to elements in compounds:
Free elements have an oxidation number of 0 (Na, H2, O2)
Monatomic ions have an oxidation number equal to their charge (Na+ is +1, Cl− is -1)
Hydrogen has an oxidation number of +1 except in metal hydrides where it is -1 (NaH)
Oxygen has an oxidation number of -2 except in peroxides where it is -1 (H2O2) and in compounds with fluorine (OF2)
Fluorine always has an oxidation number of -1 (HF, SF6)
In neutral compounds, the sum of oxidation numbers must equal 0 (KMnO4: +1 + +7 + 4(-2) = 0)
In polyatomic ions, the sum of oxidation numbers must equal the charge of the ion (SO42−: +6 + 4(-2) = -2)
Balancing half-reactions
Half-reactions separate oxidation and reduction processes
Oxidation half-reaction shows the species losing electrons (Zn→Zn2++2e−)
Reduction half-reaction shows the species gaining electrons (Cu2++2e−→Cu)
Steps to balance half-reactions:
Balance atoms other than H and O (Cr2O72−→2Cr3+)
Balance oxygen atoms by adding H2O (Cr2O72−+14H+→2Cr3++7H2O)
Balance hydrogen atoms by adding H+ ions (already balanced in the previous step)
Balance charge by adding electrons (Cr2O72−+14H++6e−→2Cr3++7H2O)
Combining half-reactions:
Multiply half-reactions to equalize electrons transferred