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Net Ionic and Redox reactions

Net Ionic and Redox reactions. Honors Chemistry. Net Ionic Reactions. All ions break up or “dissociate” in aqueous solutions Net ionic reactions only show the elements that are actually reacting. Reaction = change states ie . a queous  solid, precipitate, liquid, or gas

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Net Ionic and Redox reactions

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  1. Net Ionic and Redox reactions Honors Chemistry

  2. Net Ionic Reactions • All ions break up or “dissociate” in aqueous solutions • Net ionic reactions only show the elements that are actually reacting. • Reaction = change states • ie. aqueous  solid, precipitate, liquid, or gas • Spectator ions- are not changed (stay in an aqueous solution on both sides)

  3. AgNO3(aq) + NaCl(aq) → AgCl(ppt) + NaNO3(aq) • The compounds in aqueous solution can really be written as the separate ions, because they will break up in the solution. • ONLY the aqueous solutions break up • Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(ppt) + Na+(aq) + NO3-(aq) • This is a complete ionic equation – shows all ions in the reacting solutions.

  4. Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) → AgCl(ppt) + Na+(aq) + NO3-(aq) • Some of these ions do not participate in the reaction (they stay the same state on each side). • These are called spectator ions. • You can remove them from the equation, giving the net ionic equation • Ag+(aq) + Cl-(aq) → AgCl(ppt) • In net ionic equations you must balance the atoms AND THE CHARGE.

  5. HCl(aq) + ZnS(aq) → H2S(g) + ZnCl2(aq) • Pb(ClO4)2(aq) + NaI(aq) → PbI2(ppt) + NaClO4(aq) • Zn(s) + HCl(aq) → ZnCl2(aq) + H2(g)

  6. Ca(OH)2(aq) + H3PO4(aq) → Ca3(PO4)2(aq) + H2O(l) • Pb(C2H3O2)2(aq) + HCl(aq) → PbCl2(ppt) + HC2H3O2(aq) • (NH4)2S(aq) + Co(NO3)2(aq) → CoS(s) + NH4NO3(aq)

  7. Redox Reactions • Redox (reduction-oxidation) reactions occur when there is a transfer of electrons • Oxidation reaction:Loss of electrons • Often see a gain of oxygen/ loss of hydrogen • Reduction reaction:Gain of electrons • Often see a loss of oxygen /gain of hydrogen • LEO the lion says GER • Reduction and Oxidation always happen TOGETHER

  8. Oxidizing and Reducing Agents • Reducing Agent - donates electrons(is oxidized) • Oxidizing Agent - accepts electrons (is reduced)

  9. Oxidation Numbers • The oxidation number of a monatomic (ie. not polyatomic) ion is equal in # and sign to its ionic charge. • The oxidation number of hydrogen in a compound is always +1 (except in metal hydrides where it is -1. ex: NaH) • The oxidation number of oxygen in a compound is always -2. (Except in peroxides, where it is -1). • The oxidation number of an uncombined element is zero. • For any neutral compound, the sum of the oxidation numbers of the atoms in the compound must equal zero. • For a polyatomic ion, the sum of the oxidation numbers must equal the ionic charge of the ion.

  10. Oxidation # Practice • SO4-2 • CO3-2 • NO3- • NH4+

  11. Balancing redox • Step 1: Write the equation in ionic form. • Step 2: Write separate half-reactions for oxidation and reduction. • Step 3: Balance the atoms in the half-reactions. Use H2O and H+ to balance oxygen and hydrogen in an acid solution. Use H2O and OH- for reactions in a basic solution. • Step 5: Multiply each half-reaction by an appropriate number to make the electron charges equal. • Step 6: Add the half-reactions and subtract terms that appear on both sides of the equation. • Step 7: Check the final equation to be sure that atoms are conserved, charge is conserved, and all electrons have canceled. If spectator ions are known, add them to balance the equation.

  12. Ex. MnO4-(aq) + Cl-(aq)  Mn2+(aq) + Cl2(g) [acid solution]

  13. CN-(aq) + MnO4-(aq) → CNO-(aq) + MnO2(s) [Basic Solution]

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