Oxidation - Reduction Chemistry

Transcription

1 Oxidation - Reduction Chemistry Oxidation - Reduction Reactions Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances -- this can be a complete transfer to form ionic bonds or a partial transfer to form covalent bonds In all redox reactions: one substance loses electrons -- this substance is oxidized one substance gains electrons -- this substance is reduced There are lots of processes in the natural world (and in the laboratory) that involve redox reactions e.g., corrosion, batteries, photosynthesis/respiration, etc. Reaction between zinc and sulfuric acid Sulfuric acid (solution of H + and SO4 2- ions) Zn strip H2 bubbles Zinc loses electrons zinc is oxidized Zn(s) Zn 2+ (aq) + 2e - Hydrogen gains electrons hydrogen is reduced 2 H + (aq) + 2e - H 2 (g) Reaction between zinc and sulfuric acid Overall reaction Zn(s) + H 2 SO 4 (aq) ZnSO 4 (aq) + H 2 (g) Overall ionic reaction -- all dissolved ions are explicitly shown Zn(s) + 2 H + (aq) + SO 2 4 (aq) Zn 2+ (aq) + SO 2 4 (aq) + H 2 (g) Net ionic reaction -- includes only substances that undergo change -- ions that are present but do not react (spectator ions) are not shown Electrons are transferred from zinc to hydrogen Zn(s) + 2 H + (aq) Zn 2+ (aq) + H 2 (g)

2 Oxidation - Reduction Reactions Reaction between Cu and AgNO 3 Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances Cu(s) Ag(s) Oxidation occurs when a substance loses electrons Reduction occurs when a substance gains electrons Ag + (aq) NO 3 (aq) initial final Cu 2+ (aq) NO 3 (aq) In a redox reaction, oxidation and reduction occur simultaneously -- one cannot occur in the absence of the other initial final Cu loses electrons (oxidation) Cu(s) + 2 Ag + (aq) Cu 2+ (aq) + 2 Ag(s) Oxidation of Cu: Cu(s)! Cu 2+ (aq) + 2e - Reduction of Ag + : 2 Ag + (aq) + 2e -! 2 Ag(s) Electrons are transferred from Cu atoms to Ag + ions in solution Overall Reaction: Cu(s) + 2 AgNO 3 (aq)! 2 Ag(s) + Cu(NO 3 ) 2 (aq) Overall Ionic Equation: Cu(s) + 2 Ag + (aq) + 2 NO 3 (aq)! 2 Ag(s) + Cu 2+ (aq) + 2 NO 3 (aq) Cu(s) Ag + (aq) Ag(s) Cu 2+ (aq) Ag + gains electrons (reduction) Net Ionic Equation: Cu(s) + 2 Ag + (aq)! 2 Ag(s) + Cu 2+ (aq) Voltaic cells Zinc-copper voltaic cell A voltaic cell is a device that produces an electric current from a spontaneous redox reaction the oxidation reaction and the reduction reaction are physically separated and connected with a wire electrons transferred during the redox reaction must pass through the wire, producing an electric current the electric current is used to perform work - e.g., lighting a bulb, running an electric motor, etc. Chemical potential energy (energy stored in chemical bonds) is converted to electricity that is used to perform work Oxidation of Zn: Zn(s)! Zn 2+ (aq) + 2e - Reduction of Cu 2+ : Cu 2+ (aq) + 2e -! Cu(s) Net Ionic: Zn(s) + Cu 2+ (aq)! Cu(s) + Zn 2+ (aq) Electrons are transferred from Zn atoms to Cu 2+ ions in solution

3 Dry cell batteries A dry cell battery is a small, efficient voltaic cell that contains a non-liquid electrolyte metal cap (+) carbon rod (positive electrode) zinc case (negative electrode) manganese (IV) oxide moist paste of NaOH or KOH (electrolyte) metal bottom ( ) Oxidation of Zn: Zn(s) + 2 OH - (aq)! ZnO(s) + H2O(l) + 2e - Alkaline-type dry cell battery Electrolyte is NaOH or KOH Electrolyte is source of OH - and H2O for redox reactions Reduction of Mn 4+ : 2 MnO2(s) + H2O(l) + 2e -! Mn2O3(s) + 2 OH - (aq) The oxidation number of an atom is an integer value that represents the number of electrons gained, lost, or unequally shared by that atom an oxidation number of zero indicates that the atom has the same number of electrons assigned to it as there are in the free, neutral atom a positive oxidation number indicates that the atom has fewer electrons assigned to it than in the neutral atom a negative oxidation number indicates that the atom has more electrons assigned to in than in the neutral atom An element in its free state (uncombined with other elements) has an oxidation number of zero Ba Cl Cl barium (Ba) chlorine (Cl 2 ) The oxidation number of an element that has gained or lost electrons to form an ion is that same as its positive or negative charge Ba 2+ Cl BaCl 2 Oxidation numbers of common monoatomic ions barium ion (oxidation number: +2) chloride ion (oxidation number: 1) In an ionic compound, the ions retain their oxidation number

4 In covalent compounds (shared electrons), oxidation numbers are assigned by an arbitrary system based on relative electronegativities of the atoms H H Cl Cl hydrogen (H 2 ) chlorine (Cl 2 ) In covalent compounds containing different elements, the bonding atoms are shared unequally between the atoms -- the higher the electronegativity of the atom, the greater its affinity for the shared electrons In these types of compounds, oxidation numbers are determined by assigning both bonding electrons to the most electronegative atom In diatomic molecules containing only one element (nonpolar covalent bonding), the bonding pair of electrons is shared equally between the atoms (" electronegativity = 0) each atom is assigned an oxidation number of zero Example: Water Electronegativity Oxygen: 3.5 Hydrogen: both bonding electrons assigned to oxygen 2 H O H +1 both bonding electrons assigned to oxygen Many elements have multiple oxidation numbers It depends on the types of compounds they form Many elements have multiple oxidation numbers It depends on the types of compounds they form N 2 N 2 O NO N 2 O 3 NO 2 N 2 O 5 NO 3 N oxidation number Note: The oxidation number of oxygen is 2 in all of these compounds Elemental copper (Cu) Cu oxidation state: 0 Copper (II) sulfate (CuSO 4 ) Cu oxidation state: +2

5 Rules for assigning oxidation numbers 1. All elements in their free state (uncombined with other elements) have an oxidation number of zero (e.g., Na, Cu, Mg, H 2, O 2, Cl 2, etc.) 2. H is +1, except in metal hydrides, where it is -1 (e.g., NaH, CaH 2 ) 3. O is -2, except in peroxides, where it is -1, and in OF 2, where it is The metallic element in an ionic compound has a positive oxidation number 5. In covalent compounds, the most electronegative element is assigned a negative oxidation number 6. The sum of the oxidation numbers of the elements in a neutral compound is zero Rules for determining the oxidation number of an element within a compound Step 1: Write the oxidation number of each known atom below the atom in the formula Step 2: Multiply each oxidation number by the number of atoms of that element in the compound Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero 7. The sum of the oxidation numbers of the elements in a polyatomic ion is equal to the charge of the ion Example: Determine the oxidation number of carbon in carbon dioxide CO 2-2 2(-2) = C + (-4) = C = +4 (oxidation number for carbon) Step 1: Write the oxidation number of each known atom below the atom in the formula Step 2: Multiply each oxidation number by the number of atoms of that element in the compound Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero Example: Determine the oxidation number of sulfur in sulfuric acid H 2 SO (+1) = +2 4(-2) = -8 ( 2) S + (-8) = 0 ( 2) + 8 S = +6 (oxidation number for sulfur) Step 1: Write the oxidation number of each known atom below the atom in the formula Step 2: Multiply each oxidation number by the number of atoms of that element in the compound Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero

6 Example: Determine the oxidation number of chromium in Cr 2 O 7 2- Cr 2 O (-2) = -14 2Cr + (-14) = -2 (the charge on the ion) Cr = +6 (oxidation number for chromium) Step 1: Write the oxidation number of each known atom below the atom in the formula Step 2: Multiply each oxidation number by the number of atoms of that element in the compound Step 3: Assign oxidation numbers for the other atoms in the compound in order to make the sum of the oxidation numbers equal to zero Example: Determine the oxidation number of potassium and nitrogen in KNO 3 KNO 3 Recognize that KNO3 is an ionic compound between K + and NO 3 - The oxidation number of potassium in K + is +1 (the charge on the ion) For nitrogen: NO 3-2 K + NO 3 3(-2) = -6 N + (-6) = -1 (the charge on the ion) N = +5 (oxidation number for nitrogen) Oxidation - Reduction Reactions Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances -- this can be a complete transfer to form ionic bonds or a partial transfer to form covalent bonds Oxidation - Reduction Reactions Oxidation - reduction (redox) reactions are chemical processes that involve a transfer of electrons between substances -- this can be a complete transfer to form ionic bonds or a partial transfer to form covalent bonds L E O ose lectrons xidation G E R ain lectrons eduction In redox reactions, the oxidation numbers of the elements involved in the reaction change oxidation of an element (loss of electrons) results in an increase in its oxidation number reduction of an element (gain of electrons) results in a decrease in its oxidation number

7 Reaction between zinc and sulfuric acid Zn(s) + H 2 SO 4 (aq) ZnSO 4 (aq) + H 2 (g) zinc loses electrons the oxidation number of Zn increases zinc is oxidized hydrogen gains electrons the oxidation number of H decreases hydrogen is reduced Reaction can be rewritten to emphasize electron transfer Zn + 2 H + + SO 4 2- Zn 2+ + SO H zinc loses electrons the oxidation number of Zn increases zinc is oxidized hydrogen gains electrons the oxidation number of H decreases hydrogen is reduced Electrons are transferred from zinc to hydrogen Electrons are transferred from zinc to hydrogen Oxidizing and reducing agents Oxidizing agent: The reactant that causes another substance to be oxidized i.e., the reactant that causes an increase in the oxidation state of another substance The oxidizing agent is reduced in a redox reaction Reducing agent: The reactant that causes another substance to be reduced i.e., the reactant that causes a decrease in the oxidation state of another substance The reducing agent is oxidized in a redox reaction Reaction between zinc and sulfuric acid Zn(s) + H 2 SO 4 (aq) ZnSO 4 (aq) + H 2 (g) zinc loses electrons the oxidation number of Zn increases zinc is oxidized zinc is the reducing agent (it causes hydrogen to be reduced) hydrogen gains electrons the oxidation number of H decreases hydrogen is reduced sulfuric acid is the oxidizing agent (it causes zinc to be oxidized)

8 Example: Is the following a redox reaction? Example: Is the following a redox reaction? Neutralization reaction between hydrochloric acid and potassium hydroxide Thermite reaction HCl (aq) + KOH (aq) H 2 O (l) + KCl (aq) 2 Al (s) + Fe2O3 (s) Al2O3 (l) + 2 Fe (l) Element before reaction after reaction H O 2 2 K Cl 1 1 All oxidation numbers unchanged No redox reactions occurred Element before reaction after reaction Al 0 +3 Fe +3 0 O 2 2 Which element is oxidized? Which element is reduced? Al Fe Homework assignment Chapter 6 Problems: 6.76, 6.77, 6.78, 6.79, 6.84, 6.87, 6.88, 6.89, 6.90, 6.91, 6.92