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

• We will be discussing 5 major types of reactions

1. Synthesis2. Decomposition3. Single Replacement4. Double Replacement5. 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

• 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