Ionic equations reactions
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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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Equations l.jpg
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 l.jpg
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)


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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


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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


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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


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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


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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


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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


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Another Example

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


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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)


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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


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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


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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


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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


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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.


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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


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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!


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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-


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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


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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


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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


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Decomposition # 1

  • Metal oxyanion  metal oxide + nonmetal oxide (No Redox – NO ions)

  • A solid sample of calcium sulfate is heated

  • CaSO4 

  • CaSO4  CaO + SO3


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Decomposition # 2

  • Base  metal oxide + water (No Redox – NO ions)

  • Calcium hydroxide is decomposed

  • Ca(OH)2

  • Ca(OH)2 CaO + H2O


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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


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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-


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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


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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)


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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


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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-


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Combustion

  • Hydrocarbon + O2 CO2+ H2O (No ions)

  • Combustion of methane

  • CH4 + O2 CO2+ H2O

  • CH4 + 2O2 CO2+ 2H2O


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