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Predicting Products of Reactions

Predicting Products of Reactions. Double and Single Replacement. Double replacement. Two main types a . formation of a precipitate b . acid/base reaction (produces a salt plus water) (c. or both!) use solubility rules and electrolytic properties

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Predicting Products of Reactions

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  1. Predicting Products of Reactions Double and Single Replacement

  2. Double replacement • Two main types a. formation of a precipitate b. acid/base reaction (produces a salt plus water) (c. or both!) • use solubility rules and electrolytic properties • only break apart into ions if soluble, in solution, and a strong electrolyte

  3. Must have a driving force that removes ions from the solution • a precipitate forms and/or • a gas forms and/or • a weak/nonelectrolyte (water usually) forms

  4. Formation of a precipitate • A precipitate is an insoluble compound (a solid) • Use solubility rules to determine if one or both of the products is a precipitate • Example 3AgNO3 (aq) + K3PO4 (aq)  Ag3PO4 (s) + 3KNO3 (aq) insoluble soluble precipitate Net ionic eqn: 3Ag+(aq) + PO4-3(aq)  Ag3PO4(s)

  5. Formation of a gas • Must know common gases • If you get H2CO3 when you do the double replacement it really gives H2O + CO2 (g) • If you getH2SO3 when you do the double replacement it really gives H2O + SO2 (g) • H2S is a gas

  6. For example Na2CO3(aq)+ H2SO4(aq)Na2SO4(aq)+ H2CO3(aq) REALLY IS: Na2CO3(aq)+H2SO4(aq)Na2SO4(aq)+H2O(l)+CO2(g) Net ionic eqn: CO3-2(aq) + 2H+(aq)  H2O(l) + CO2(g)

  7. Formation of a nonelectrolyte • Water is a common nonelectrolyte and is a result of an acid-base reaction HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l) Net ionic eqn: H+ (aq)+ OH- (aq) H2O(l)

  8. Redox Reactions • Some synthesis, some decomposition, single replacement, and other reactions are oxidation-reductions • OXIDATION-REDUCTION REACTIONS involve electron transfer

  9. Terms to Know: • OIL RIG – oxidation is loss, reduction is gain (of electrons) • Oxidation – the loss of electrons, increase in charge (becomes more positive) • Reduction – the gain of electrons, reduction of charge (becomes more negative) • Oxidation number – the assigned charge on an atom • Oxidizing agent (OA) – the species that is reduced and thus causes oxidation • Reducing agent (RA) – the species that is oxidized and thus causes reduction • *note that some in some reactions, the same species can be oxidized and reduced; these are called disproportionation reactions

  10. Rules for Assigning Oxidation States • 1. The oxidation state of an atom in an element is ZERO including allotropes [i.e. N2, P4, S8]. • 2. The oxidation state of a monatomic ion is the same as its charge. • 3. In its compounds, fluorine is always assigned an oxidation state of -1. • 4. Oxygen is usually assigned an oxidation state of -2 in its covalent compounds, such as CO, CO2, SO2, and SO3. • Exceptions to this rule include peroxides (compounds containing the O22- group), where each oxygen is assigned an oxidation state of -1, as in hydrogen peroxide (H2O2), • and OF2 in which oxygen is assigned a +2 oxidation state. • 5. In its covalent compounds with nonmetals and acids, hydrogen is assigned an oxidation state of +1. • Metal hydrides are an exception; H is at the end of the chemical formula since it has an oxidation state of -1. (such as NaH) • 6. The sum of the oxidation states must be zero for an electrically neutral compound. For a polyatomic ion, the sum of the oxidation states must equal the charge of the ion.

  11. There can be non-integer oxidation states like in Fe3O4. There’s a -8 for the 4 oxygens divided across 3 iron ions, therefore Fe’s charge is Fe8/3+

  12. Single Replacement Redox • How can we predict if a single replacement reaction will occur? • Reading the reduction potential chart •  elements that have the most positive reduction potentials are easily reduced (in general, non-metals) •  elements that have the least positive reduction potentials are easily oxidized (in general, metals) • Can also be used as an activity series. Metals having less positive reduction potentials are more active and will replace metals with more positive potentials. • The MORE POSITIVE reduction potential gets to indeed be reduced IF you are trying to set up a cell that can act as a battery.

  13. For example, Will the following reaction occur? Mg(s) + 2KCl(aq)  2K(s) + MgCl2(aq) Step1: Determine which element is reduced by assigning oxidation numbers before and after Mg(s) + 2KCl(aq)  2K(s) + MgCl2(aq) 0 +1 -1 0 +2 -1 So K was reduced b/c it went from +1 to 0

  14. Step 2: Look up the reduction potentials of the elements whose oxidation #’s changed K= -2.92 Mg=-2.37 • Step 3: Determine if the element with the most +reduction potential was reduced • NO! K was reduced in the reaction as written, but Mg has the most positive reduction potential! • K cannot be reduced by Mg, but Mg could be reduced by K • Step 4: Does this reaction occur? • NO! (but note that means that the reverse reaction would occur!)

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