Redox Reactions and Other Chemical Reactions PDF

Cebu Doctors’ University College of Arts and Sciences Physical Sciences Department REDOX REACTION AND OTHER CHEMICAL REACTIONS Compiled by: Joselito R. Tumulak Jr., RChT, MS (cand.) Analytical Chemistry Professor I. OXIDATION-REDUCTION (REDOX) REACTION Oxidation-reduction reaction (commonly known as...

Cebu Doctors’ University College of Arts and Sciences Physical Sciences Department REDOX REACTION AND OTHER CHEMICAL REACTIONS Compiled by: Joselito R. Tumulak Jr., RChT, MS (cand.) Analytical Chemistry Professor I. OXIDATION-REDUCTION (REDOX) REACTION Oxidation-reduction reaction (commonly known as redox reaction) involves loss and gain of electrons. Oxidation is loss of electrons, while reduction is gain of electrons. The simple mnemonic "LEO GER" or Loss of Electrons - Oxidation; Gain of Electrons - Reduction has been used by several generations of students to remember these definitions. The equation below shows an obvious example of a simple redox reaction: Cu2+ + Mg à Cu + Mg2+ In the above equation, magnesium loses electrons to copper(II) ions which neutralizes its charge. In return, magnesium now has a positive charge. Now, if you’re going to track the movements of electrons, you would notice that copper(II) ion underwent reduction because it is the one that gained the electrons. On the other hand, Mg underwent oxidation for losing electrons. Other terminologies you would need to learn are the terms oxidizing agent and reducing agent. Oxidizing agent oxidizes something else (causes the other to lose electrons). Thus, it takes electrons from that other substance. Reducing agent reduces something else (causes the other to gain electrons). Thus, it donates electrons to that other substance. In the example above, Cu(II) ions is an oxidizing agent because it causes Mg to lose it electrons. Conversely, Mg is a reducing agent because it donated electrons to the Cu(II) ions. II. OXIDATION NUMBERS To recognize redox reactions, we must be able to identify when a species is oxidized or reduced; to do this we assign oxidation numbers to each atom in a reaction. Then, we compare the oxidation states of each atom on the reactants side and to the atoms on the products side. When changes occur, we know a redox reaction has taken place. A series of rules have been developed to determine oxidation numbers: 1. For free elements (uncombined state), each atom has an oxidation number of zero. H2, Br2, Na, Be, K, O2, and P4 all have an oxidation number of 0. 2. Monatomic ions have oxidation numbers equal to their charge. Li+=+1, Ba2+=+2, Fe3+=+3, I−=−1, O2−=−2, etc. Alkali metal oxidation numbers =+1. Alkaline earth oxidation numbers =+2. Aluminum =+3 in all of its compounds. Oxygen's oxidation number =−2 except when in hydrogen peroxide (H2O2), or a peroxide ion (O2−2) where it is −1. 3. Hydrogen's oxidation number is +1, except for when bonded to metals as the hydride ion forming binary compounds. In LiH, NaH, and CaH2, the oxidation number is −1. Page 1 of 3 4. Fluorine has an oxidation number of −1 in all of its compounds. 5. Halogens (Cl, Br, I) have negative oxidation numbers when they form halide compounds. When combined with oxygen, they have positive numbers. In the chlorate ion (ClO3-), the oxidation number of Cl is +5, and the oxidation number of O is −2. 6. In a neutral atom or molecule, the sum of the oxidation numbers must be 0. In a polyatomic ion, the sum of the oxidation numbers of all the atoms in the ion must be equal to the charge on the ion. Example: 1. What is the oxidation number of manganese in the polyatomic ion permanganate: MnO41-? Solution: o The oxidation number of O is -2 (Rule 2). o Since this is a polyatomic ion, the sum of the oxidation numbers of all atoms in the ion mus equal to the charge of the ion (Rule 6). o So we have: Mn + 4(-2) = - 1 Mn – 8 = -1 Mn = +7 2. What is the oxidation number of chromium in the molecule K2Cr2O7? Solution: o The oxidation number of K is +1 (Rule 2). o The oxidation number of O is -2 (Rule 2). o Since this is a compound (there is no charge indicated on the molecule), the net charge on the molecule is zero (rule 6). o So we have: 2(+1) + 2 Cr + 7 (-2) = 0 2 + 2 Cr – 14 = 0 2 Cr = 12 Cr = +6 Changes of Oxidation Number in Redox Reactions Consider the reaction below between elemental iron and copper sulfate: Fe + CuSO4 à FeSO4 + Cu2+ In the course of this reaction, the oxidation number of Fe increases from zero to +2. The oxidation number of copper decreases from +2 to 0. A loss of negatively charged electrons corresponds to an increase in oxidation number, while a gain of electrons corresponds to a decrease in oxidation number. Therefore, the element or ion that is oxidized undergoes an increase in oxidation number. The element or ion that is reduced undergoes a decrease in oxidation number. Example: Use changes in oxidation number to determine which atoms are oxidized and which atoms are reduced in the following reaction. Identify the oxidizing and reducing agent. Fe2O3 + 3 CO à 2 Fe + 3 CO2 Page 2 of 3 Solution: 1. Use the oxidation number rules to assign oxidation numbers to each atom in the balanced equation. Coefficients do not affect oxidation numbers. The oxidized atom increases in oxidation number and the reduced atom decreases in oxidation number. So we have: 2. Track the changes in oxidation numbers to ascertain which undergoes oxidation and reduction. § The element carbon is oxidized because its oxidation number increases from +2 to +4. The iron (III) ion within the Fe2O3 is reduced because its oxidation number decreases from +3 to 0. The carbon monoxide (CO) is the reducing agent since it contains the element that is oxidized. The Fe3+ ion is the oxidizing agent since it is reduced in the reaction. III. IDENTIFYING REACTION TYPES As was already mentioned in the previous section, A redox reaction must involve a change in oxidation number for two of the elements involved in the reaction. The oxidized element increases in oxidation number, while the reduced element decreases in oxidation number. That is why in reactions where oxidation numbers of the elements are not affected, they are not considered redox reactions. Types of Reactions That Are Not Redox 1. Double Displacement Reaction Observe the reaction below: Na2SO4 + Ba(NO3)2 à 2 NaNO3 + BaSO4 If you assign the oxidation numbers of the elements involved in the reaction, you will notice that oxidation numbers remain unchanged. That is why, Double-replacement reactions, such as the one above, are not redox reactions because ions are simply recombined without any transfer of electrons. 2. Acid-base reactions The reaction below is an acid-base reaction: HF + NH3 à NH4 + FAcid-base reactions, such as above, involve transfer of hydrogen ions. If you assign the oxidation numbers of the elements involved in the reaction, you will notice that oxidation numbers remain unchanged. Again, the transfer of an H+ ion leaves the oxidation numbers unaffected. Page 3 of 3