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CH301 - chapter 11b notes

CH301 - chapter 11b notes - 1 CH 301 Chapter 11b Reactions...

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1 CH 301 Chapter 11b Reactions in Aqueous Solutions II: Calculations Oxidation – Reduction (Transfer of electrons) Ch. 4 4-4: Oxidation Numbers To assign oxidation numbers to elements in compounds, there are a few simple rules: For simple Group A binary ionic compounds (salts), the oxidation state is the charge of the element: 1. Metals (Group I A, IIA, and IIIA) are assigned a positive oxidation state – determined by the number of electrons the element has lost. 2. Nonmetals (Group VIA, VIIA, VIIIA) are assigned a negative oxidation state determined by the number of electrons the element has gained. For simple binary ionic compounds that are comprised of transition metal (with more than one possible oxidation state), the oxidation state is the charge of the element: 1. For our purpose, all transition metals in salts (except zinc and silver) have oxidation states that are determined by the balancing the charge of the anion. e.g. FeO vs. Fe 2 O 3 -2 FeO Since the oxidation state of oxygen is 2 in the compound, the iron atom must have a charge of + 2 to give the formula unit an overall net charge of zero. x + (-2) = 0 x = +2 Fe 2 O 3 Since the oxidation state of oxygen is 2 in the compound and there are three oxygen atoms, the iron must have a charge of +3 to give the formula unit an overall net charge of zero. 2x + 3(-2) = 0 x = +3 2. Zinc always has a +2 charge, and silver always has a +1 charge. Group IVA elements tin and lead have two possible oxidations states that each may have in an ionic bond. For covalently bonded compounds, the oxidation states of the elements are determined by assigning either positive or negative oxidation states to each element by following the rules given on page 138 in your text. Some of rules for assigning oxidation states to all elements in a covalently bonded molecule are indicated below:

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2 1. Br 2 I 2 N 2 Cl 2 H 2 O 2 F 2 (the diatomic molecules), have zero oxidation states. This makes sense— there is no dipole (difference in Electronegativity) between the shared electrons in the diatomic molecules. 2. Metals, in their natural uncombined state, have zero oxidation states. E.g. Na(s), K(s), Mg(s), Fe(s), Zn(s) 3. Nonmetals, in their natural uncombined state, have zero oxidation states. E.g. P 4 (s) and S 8 (s) 4. Hydrogen is almost always +1 in a compound (except as a hydride), and oxygen is almost always 2 in a compound (except when it is bound to an element that forces it to be 1 or 1/2) E.g. H 2 O (hydrogen is +1) NaH (hydrogen is 1) H 2 O (oxygen is 2) H 2 O 2 (oxygen is 1)
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CH301 - chapter 11b notes - 1 CH 301 Chapter 11b Reactions...

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