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Ionic Equations & Reactions

Ionic Equations & Reactions. Equations. Molecular equations – show the complete chemical formulas. Does not indicate ionic character Complete ionic equation – shows all ions. Actually how the particles exist in the solution. Steps for Writing Ionic Equations.

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Ionic Equations & Reactions

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  1. Ionic Equations & Reactions

  2. Equations • Molecular equations – show the complete chemical formulas. Does not indicate ionic character • Complete ionic equation – shows all ions. Actually how the particles exist in the solution

  3. Steps for Writing Ionic Equations • Write the balances molecular equation (balanced chemical equation) • Break every thing down into its ions EXCEPT the solid, gas, water, or weak electrolyte (complete ionic equation) • Cross out everything that is the same on both sides (spectator ions) • Write what is left (net ionic equation)

  4. Rules • When writing ionic equations, you must keep together the solid, gas, water, or weak electrolyte • Spectator ions – ions that appear on both sides of the equation. They have very little to do with the chemical reaction

  5. Example • Write the balanced chemical equation, the complete ionic equation, and the net ionic equation for the reaction between lead (II) nitrate and potassium iodide

  6. Example • Write the balanced chemical equation • Pb(NO3)2 + 2 KI  PbI2 + 2 KNO3 • You MUST identify the solid, gas, or water • Pb(NO3)2 + 2 KI  PbI2(s) + 2 KNO3 • Balanced chemical equation

  7. Example • Now break every thing except the solid, gas, or water into its ions • Remember ions are things with charges • Everything will be broken down into one positive charge and one negative charge

  8. Example • Pb(NO3)2 + 2 KI  PbI2(s) + 2 KNO3 • Pb+2 + 2NO3-1 + 2K+1 + 2 I -1 PbI2 (s) + 2K+1 + 2NO3-1 • Complete ionic Equation

  9. Example • Now cross out everything that is the same on both sides (spectator ions) • Pb+2 + 2NO3-1 + 2K+1 + 2 I -1 PbI2 (s) + 2K+1 + 2NO3-1 • Pb+2 + 2NO3-1 + 2K+1 + 2 I -1 PbI2 (s) + 2K+1+ 2NO3-1 • Now write what is left • Pb+2 + 2 I -1 PbI2 (s) • Net ionic equation

  10. Another Example • Write the balanced chemical equation, complete ionic equation, and net ionic equation for the reaction between calcium chloride and sodium acetate

  11. Another Example • Balanced chemical equation • CaCl2 + Na2CO3 CaCO3(s) + 2NaCl • Complete ionic equation • Ca+2 + 2Cl -2 + 2Na +1 + CO3 -2  CaCO3 (s) + 2Na +1 + 2Cl -1 • Net Ionic Equation • Ca+2 + 2Cl -2 + 2Na +1 + CO3 -2  CaCO3 (s) + 2Na +1 + 2Cl -1 • Ca+2 + CO3 -2  CaCO3 (s)

  12. What if water is formed? • Write the balanced chemical equation, complete ionic equation, and net ionic equation for the reaction between Calcium hydroxide and nitric acid

  13. Example with water • Balanced chemical equation • Ca(OH)2 + 2 HNO3 Ca(NO3)2 + 2 HOH • Complete ionic equation • Ca+2 + 2(OH) -1 + 2H+1 + 2NO3 -1  Ca+2 + 2NO3 -1 + 2 HOH • Net Ionic Equation • Ca+2 + 2(OH) -1 + 2H+1 + 2NO3 -1 Ca+2 + 2NO3 -1+ 2 HOH • 2(OH) -1 + 2H+1  2 HOH

  14. 5 Major Types of Reactions • We will be discussing 5 major types of reactions • Synthesis • Decomposition • Single Replacement • Double Replacement • Combustion • You need to know these reactions! • Note cards are an extremely effective way to remember them

  15. Synthesis # 1 • Metal oxide + nonmetal oxide  metal oxyanion (NO ions – No Redox) • No Redox simply means that the oxidation numbers of the elements stays the same

  16. Synthesis # 1 Example • Sulfur dioxide gas is passed over solid calcium oxide • SO2 + CaO  • We know that we have to get a metal oxyanion. • So we either get CaSO4 or CaSO3 • We need to check the oxidation states on sulfur to see which one is the same.

  17. Synthesis # 1 Example • In SO2, the oxidation number of O is -2 • So the oxidation number of S must be +4 • Our product choices are CaSO3 or CaSO4 • In CaSO3…S has an oxidation # of +4 • In CaSO4…S has an oxidation # of +6 • Therefore the product must be CaSO3 • SO2 + CaO  CaSO3

  18. Synthesis # 2 • Metal oxide + water  strong base (IONS) • Strong acids & bases ionize completely in water & are therefore electrolytes. • They will be written as ions • Strong bases…Group !a or 2A hydroxides • There are 7 strong acids… • HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4 • You MUST know these!

  19. Synthesis # 2 Example • Solid sodium oxide is added to water • Na2O + H2O  • Na2O + H2O  NaOH • Na2O + H2O  2NaOH • Na2O stays together because it is solid • H2O stays together because it is water • NaOH is separated because it is a strong base • Na2O + H2O  2Na+ + OH-

  20. Synthesis # 3 • Non metal oxide + water  oxyacid (weak molecules…strong ions…No Redox) • Sulfur dioxide gas is placed in water • SO2 + H2O  • We are going to get an oxyacid…so we either have H2SO3 or H2SO4 • The S needs to have the same oxidation number

  21. Synthesis # 3 Example • In SO2, O has an oxidation # of -2…so S has an oxidation # of +4 • In H2SO3…S has an oxidation # of +4 • In H2SO4…S has an oxidation # of +6 • Therefore we will get In H2SO3 • SO2 + H2O  H2SO3 • Since H2SO3 is a weak acid…we will keep it together

  22. Synthesis # 4 • Metal + nonmetal  salt (NO ions) • A salt is just an ionic compound ( a positive charge & a negative charge) • Magnesium metal is combusted in nitrogen gas • Mg + N2 • Mg + N2 Mg3N2 • 3Mg + N2 Mg3N2

  23. Decomposition # 1 • Metal oxyanion  metal oxide + nonmetal oxide (No Redox – NO ions) • A solid sample of calcium sulfate is heated • CaSO4  • CaSO4  CaO + SO3

  24. Decomposition # 2 • Base  metal oxide + water (No Redox – NO ions) • Calcium hydroxide is decomposed • Ca(OH)2 • Ca(OH)2 CaO + H2O

  25. Single Replacement # 1 • Metal + ionic solution  Metal ion + metal (will have ions) • Must look at activity series! • Aluminum metal is added to a solution of copper (II) chloride • Al + CuCl2 • Al + CuCl2 AlCl3 + Cu • 2Al + 3CuCl2 2AlCl3 + 3Cu • 2Al + 3Cu +2  2Al +3 + 3Cu

  26. Single Replacement # 2 • Active metal (Group 1A, Ba, Ca, Sr) + water  H2 + strong base (IONS) • Sodium is placed in water • Na + H2O  • Na + H2O  H2 + NaOH • 2Na + 2H2O  H2 + 2NaOH • 2Na + 2H2O  H2 + 2Na+ + 2OH-

  27. Single Replacement # 3 • Halogen + metal halide  new metal halide + halogen (REDOX…will have ions) • Chlorine gas is bubbled into a solution of sodium bromide • Cl2 + NaBr  • Cl2 + NaBr  NaCl + Br2 • Cl2 + 2NaBr  2NaCl + Br2 • Cl2 + 2Br- 2Cl- + Br2

  28. Double Replacement # 1 • Precipitate (must know solubility rules)…the precipitate will stay together • A saturated solution of barium hydroxide is mixed with a solution of iron (III) sulfate • Ba(OH)2 + Fe2(SO4)3 • Ba(OH)2 + Fe2(SO4)3 Fe(OH)3 + BaSO4(s) • 3Ba(OH)2 + Fe2(SO4)3 2Fe(OH)3 + 3BaSO4(s) • 3Ba+2 + 3SO4-2 3BaSO4(s)

  29. Double Replacement # 2 • Formation of a gas (acid + sulfide, carbonate, or bicarbonate) • Hydrobromic acid is added to a solution of potassium bicarbonate • HBr + KHCO3 • HBr + KHCO3 H2CO3 + KBr • H2CO3 ALWAYS breaks down into CO2 + H2O • HBr + KHCO3CO2 + H2O + KBr • H+ + HCO3-CO2 + H2O

  30. Double Replacement # 3 • Metal hydride + water  H2 + strong base (IONS) • Sodium hydride is placed into water • NaH + H2O  • NaH + H2O  H2 + NaOH • NaH + H2O  H2 + Na+ + OH-

  31. Combustion • Hydrocarbon + O2 CO2+ H2O (No ions) • Combustion of methane • CH4 + O2 CO2+ H2O • CH4 + 2O2 CO2+ 2H2O

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