Oxidation-Reduction Reactions Redox Reactions Reactions in which one or more electrons are transferred 2Na (s) + Cl 2(g) 2NaCl (s) (Redox) Species loosing electrons is said to be oxidized and the species gaining electrons is said to be reduced Not Redox Reactions Types of Redox Rxn s Precipitation Double displacement Neutralization Ions Rearrange Metathesis Synthesis Decomposition Single Replacement Electrons are transferred in most cases Combustion Combustion of methane CH 4 + O 2 CO 2 + H 2 O Carbon looses electrons and oxygen gains electrons Notice that in the reaction, there are no ions present as either reactants or products, however, electrons are still transferred An understanding of oxidation states will show the transfer of electrons Oxidation State (Oxidation Numbers) A system of bookkeeping for electrons in molecules or ions The imaginary charges atoms would have in molecules and real charges in ionic
Oxidation states provide a way to track electrons in chemical reactions Especially useful for reactions containing covalent which have no charges Oxidation and Reduction Note: Charges are written: Oxidation states are written: Fe 2+, Cl 1- Fe +2, Cl -1 A species is oxidized when it loses electrons. Here, zinc loses two electrons to go from neutral zinc metal to the Zn 2+ ion. Oxidation and Reduction Oxidation and Reduction A species is reduced when it gains electrons. Here, each of the H + gains an electron and they combine to form H 2. What is reduced is the oxidizing agent. H + oxidizes Zn by taking electrons from it. What is oxidized is the reducing agent. Zn reduces H + by giving it electrons. Movie 1. The oxidation state of an atom in an element is zero Na (s), O 2(s), Hg (l) 2. The oxidation state of a monatomic ion is the same as its charge Na +, Cl -, N -3, Ca +2 Alkali metals, alkaline earth metals, Al, Ga, Ge.
3.Fluorine is always -1 in its HF, PF 3 H +1 F -1, P +3 F -1 Notice the oxidation states are written +1, whereas a charge would be written 1+ 4.Oxygen is always -2 in its H 2 O, CO 2 Exceptions: Peroxides (containing O 2 2- ) in which oxygen is -1 5.Hydrogen is +1 in its covalent - In ionic (metal bonded to non-metal) hydrogen can have a negative oxidation state. H 2 O, HCl, HN 3 6. Other oxidation states calculated from algebraic sum of known states Transitions, metalloids, and usually carbon, silicon, and boron. 7. In non-described the atom with the greatest electronegativity is assigned a negative oxidation state equal to its imaginary charge in an ionic compound CN, CN - SF 6 PbS CO 2 AsH 2 NO 3 - NH 4 + K 2 Cr 2 O 7 P 4 S 2 EXAMPLES
Noninteger Oxidation States Rare, but do occur due to the arbitrary way electrons are divide up using rules for oxidation states Used the same way as integer oxidation states Example: Fe 3 O 4 Identify the oxidation state for each of the following in the reaction CH4 + O 2 CO 2 + H 2 O Notice, carbon has changed from a 4 to a + 4 oxidation state and oxygen has changed from 0 to 2 and 2 Carbon was oxidized and oxygen was reduced Oxidation oxidation state becomes more positive loss of electrons Reduction oxidation state becomes more negative gain of electrons 1. Identify the species being reduced and oxidized in the reaction of aluminum metal with solid iodine to produce aluminum triiodide. Oxidizing agent Reducing Agent chemical species containing element that is oxidized (donates electrons) chemical species containing element that is reduced (accepts electrons)
Common Oxidizing and Reducing Agents Metals (Cu) are reducing agents HNO 3 is an oxidizing agent Cu + HNO 3 --> Cu 2+ + NO 2 2 K + 2 H 2 O --> 2 KOH + H 2 Metals (Na, K, Mg, Fe) are reducing agents This is why oxidizers are stored separate from reducers 2. Identify the reducing and oxidation agent in the demo reaction Recognizing a Redox Reaction Reaction Type Oxidation Reduction In terms of oxygen gain loss In terms of halogen gain loss In terms of electrons loss gain Synthesis A.K.A. (combination) A + B AB elements Any atom or group of atoms combine to form a more complex compound Examples: Formation of metal oxides or metal sulfides Mg (s) + O 2(g) MgO (s) Formation of metal and non-metal halides 2C (s) + 3F 2(g) 2 CF 3(s) Decomposition AB A + B compound elements and/or simpler A compound is broken down into simpler substances
Examples: Electrolysis 2H 2 O 2H 2 (g) + O 2 (g) Decomposition of polyatomic 2KClO 3 (s) e 2KCl (s) + 3O 2 (g) Combustion Substance + O 2 Oxide of Elements hydrocarbon + O 2 CO 2 + H 2 O Oxides are formed from each element in the compound being combusted. A hydrocarbon in combustion reactions refer to only carbon, hydrogen and oxygen containing Examples: Formation of metal oxides Ca (s) + O 2 (g) CaO (s) Formation of non metal oxides S 8 (s) + 8O 2 (g) 8SO 2(s) Hydrocarbon combustion Single Displacement A + BX AX + B a more reactive element replaces a less reactive element in a compound (involves ions) Predicted using the Activity Series Examples: Hydrogen replacement Na (s) + H 2 S (aq) Na 2 S (aq) + H 2(g) Halogen replacement Look through the reaction reference and recognize how the different reaction classes are redox in nature. Br 2(l) + 2KI (aq) 2KBr (aq) + I 2(g) Metal replacement Cu (s) + AgNO 3(aq) Cu(NO 3 ) 2(aq) + Ag (s) Compare activity series to electrical potential series
Exceptions to the general patterns of chemical reactivity Identify the types of chemical reactions that appear to be redox by nature but show no change in oxidation states listed in the reaction reference.