Transcript Ionic Equations

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
1. Write the balances molecular equation
(balanced chemical equation)
2. Break every thing down into its ions
EXCEPT the solid, gas, water, or weak
electrolyte (complete ionic equation)
3. Cross out everything that is the same on
both sides (spectator ions)
4. Write what is left (net ionic equation)
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
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
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
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
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
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
Another Example
• Write the balanced chemical equation,
complete ionic equation, and net ionic
equation for the reaction between
calcium chloride and sodium acetate
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)
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
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
5 Major Types of Reactions
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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
Synthesis # 1
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
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.
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
Synthesis # 2
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!
Synthesis # 2 Example
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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-
Synthesis # 3
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
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
Synthesis # 4
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
Decomposition # 1
1. Metal oxyanion  metal oxide +
nonmetal oxide (No Redox – NO ions)
• A solid sample of calcium sulfate is
heated
• CaSO4 
• CaSO4  CaO + SO3
Decomposition # 2
2. Base  metal oxide + water (No
Redox – NO ions)
• Calcium hydroxide is decomposed
• Ca(OH)2 
• Ca(OH)2  CaO + H2O
Single Replacement # 1
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
Single Replacement # 2
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-
Single Replacement # 3
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
Double Replacement # 1
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)
Double Replacement # 2
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
Double Replacement # 3
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-
Combustion
1. Hydrocarbon + O2 CO2+ H2O (No
ions)
• Combustion of methane
• CH4 + O2 CO2+ H2O
• CH4 + 2O2 CO2+ 2H2O