Chemistry 20 | Mrs. Sample


Balancing Chemical Reactions

  • Every chemical reaction involves the rearrangement of atoms into different combinations. However, during these reactions, the total number of atoms of each type of element is the same after the reaction as it was before the reaction.
  • Chemical reactions have to be properly balanced in order to clearly obey the Law of Conservation of Matter.
  • Every chemical reaction must first be written so that each reactant and product has the correct chemical formula and state of matter.
  • Coefficients are then used in order to balance the various atoms. All coefficients must be in the simplest whole number ratio possible.
  • Hints for balancing equations:
  1. Write out an unbalanced equation, making sure that each reactant and product formula is written correctly.
  2. Use coefficients to balance any atoms that are not already balanced.
  3. If a polyatomic ion is remaining intact, it may be easiest to balance it as a group.
  4. Always balance any elements last.
  5. Make sure that the coefficients have the simplest whole number ratio possible.
  • EX: Re-write the following word equations as balanced chemical equations.

(a)sodium + chlorine → sodium chloride

(b)aluminium chlorate → aluminium + chlorine + oxygen

(c)butane(C4H10(g)) + oxygen → carbon dioxide + water vapour

(d)scandium + copper(II) sulfate → copper + scandium sulfate

(e)hydrochloric acid + barium hydroxide → water + barium chloride

Formation Reactions

  • A formation reaction is a reaction in which two or more elements react together to form a compound.

X + Y → XY

N2(g) + 3 H2(g) → 2 NH3(g)

  • EX: Write a balanced chemical equation for the formation of glucose (C6H12O6(s)).
  • EX: Write a balanced chemical equation for the formation of ammonium benzoate.

Decomposition Reactions

  • A decomposition reaction is a reaction in which a compound reacts and breaks down into its component elements.

XY → X + Y

2 H2O(l) → 2 H2(g) + O2(g)

  • EX: Write a balanced chemical equation for the decomposition of diphosphorousheptaoxide.
  • EX: Write a balanced chemical equation for the decomposition of sodium sulfate.

Combustion Reactions

  • A complete combustion reaction is a reaction in which a hydrocarbon burns in a plentiful supply of oxygen and the only products are carbon dioxide and water vapour.

CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)

  • EX: Write a balanced chemical equation for the complete combustion of hexane (C6H14(l)).
  • EX: Write a balanced chemical equation for the complete combustion of methanol (CH3OH(l)).

Practice: Balancing Chemical Reactions

  1. Balance the following chemical equations by correctly filling in the coefficients.

a)_____ Ni (s) + _____ Br2 (l) → _____ NiBr3 (aq)

b)_____ Al2(CO3)3 (s) → _____ Al2O3 (s) + _____ CO2 (g)

c)_____ C5H12 (aq) + _____ O2 (g) → _____ CO2 (g) + _____ H2O (g)

d)_____ CrCl2 (aq) + _____ Sn (s) → _____ SnCl4 (aq) + _____ Cr (s)

e)_____ ZnCl2 (aq) + _____ (NH4)3PO4 (aq) → _____ NH4Cl (aq) + _____ Zn3(PO4)2 (s)

  1. Write a balanced chemical equation for each of the following.

a)The complete combustion of dodecane (C12H26(l))

b)The formation of calcium dihydrogen phosphate

c)The decomposition of titanium(IV) cyanide

d)The formation of sulfur trioxide

e)The complete combustion of benzoic acid.

Single Replacement Reactions

  • A single-replacement reaction involves an atom of one element taking the place of an atom of another element.
  • These reactions typically involve an ionic compound reacting with an element.
  • Either the positive ion or the negative ion of the compound can be replaced.

A + BX AX + B

AX + Y AY + X

Cu(s) + 2 AgNO3(aq) 2 Ag(s) + Cu(NO3)2(aq)

  • EX: Write a balanced chemical equation for the reaction between zinc and hydrochloric acid.
  • EX: Write a balanced chemical equation for the reaction between aluminium chloride and bromine.

Double Replacement Reactions

  • A double-replacement reaction involves two ionic compounds reacting together. During the reaction, the positive ions switch places.
  • Generally, the ionic compounds are dissolved (if soluble). These reactions often result in the production of a precipitate (an insoluble ionic compound).
  • Neutralization reactions represent a special case of double-replacement in which one of the compounds is an acid and the other compound is a base.

AX + BY AY + BX

CuSO4(aq) + 2 NaOH(aq) Na2SO4(aq) + Cu(OH)2(s)

  • EX: Write a balanced chemical equation for the reaction between aluminium nitrate and ammonium phosphate.
  • EX: Write a balanced chemical equation for the reaction between hydrosulfuric acid and potassium hydroxide.

Calculating Amounts

The coefficients in a balanced chemical equation represent the relative amounts of the chemicals that react with each other.

The actual amount of a chemical can be found by using one of the following equations:

/ use when the mass of a pure substance is given
/ use when the volume and molar concentration of a solution is given
/ use when the pressure, volume, and temperature of a gas is given

EX: Calculate the amount of potassium carbonate in a 33.5 g sample.

EX: Calculate the amount of zinc nitrate that must be dissolved to prepare 750 mL of a solution that has a concentration of 15.8 mmol/L.

EX: Calculate the amount of helium in a 8.00 L balloon if it has a pressure of 103 kPa at 22.0oC.

Practice: Balancing Chemical Reactions

  1. Change each of these word equations into a balanced chemical equation. Classify each of these reactions as one of the five common types of reactions.
  1. carbon + hydrogen + oxygen benzoic acid
  1. ethanol(C2H5OH(l)) + oxygen carbon dioxide + water
  1. lead(II) chlorate + germanium germanium chlorate + lead
  1. cobalt(III) hydrogen sulfate cobalt + hydrogen + sulfur + oxygen
  1. chromium(III) sulfate + calcium nitrate

calcium sulfate + chromium(III) nitrate

  1. Write a balanced chemical equation for each of the following reactions:
  1. The decomposition of nickel(II) dichromate.
  1. The combustion of C6H13OH(l).
  1. The formation of titanium(IV) thiosulfate.
  1. The reaction between aluminium and iron(III) nitrate.
  1. The reaction between ammonium sulfide and nickel(III) chlorate.

Net Ionic Equations

  • In replacement reactions, many of the reactants and products exist as dissociated ions.
  • Some of these dissociated ions remain unchanged in any way throughout the reaction.
  • For example, in the reaction between solutions of silver nitrate and sodium chloride:

AgNO3(aq) + NaCl(aq) NaNO3(aq) + AgCl(s)

  • The Ag+ and Cl- react to form a solid precipitate.
  • However, the Na+ and NO3- remain dissolved in both the reactants and the products.
  • They are classified as spectator ions, because they are left unchanged by the reaction itself.
  • DEMO: Mix together a test tube of potassium iodide with a test tube of lead(II) nitrate.
  • The complete balanced equation for this reaction would be

2 KI(aq) + Pb(NO3)2(aq) 2 KNO3(aq) + PbI2(s)

  • An ionic equation shows all soluble ionic compounds as dissociated ions. Strong acids are also shown as being completely ionized.
  • The ionic equation for the previous reaction would be

2 K+(aq) + 2 I-(aq) + Pb2+(aq) + 2 NO3-(aq) 2 K+(aq) + 2 NO3-(aq) + PbI2(s)

  • The K+(aq) and NO3-(aq) are the spectator ions. They remain unchanged throughout the reaction.
  • The net ionic equation omits the spectator ions.
  • The net ionic equation for this reaction is:

Pb2+(aq) + 2 I-(aq) PbI2(s)

  • EX: Write a complete balanced equation, ionic equation, and net ionic equation for the reaction between barium hydroxide and hydrochloric acid.
  • Ionic and net ionic equations can also be written for single replacement reactions.
  • EX: Write a complete balanced equation, ionic equation, and net ionic equation for the reaction between aluminium and copper(II) nitrate.

Practice: Ionic and Net Ionic Equations

  1. Write a complete balanced equation for the following reactions:
  1. The formation of manganese(II) phosphate
  1. The decomposition of potassium hydrogen oxalate
  1. The complete combustion of benzene (C6H6(l))
  1. Write an ionic equation for each of the following replacement reactions:
  1. The reaction between aluminium and iron(III) nitrate.
  1. The reaction between ammonium sulfide and nickel(III) chlorate.
  1. Write a net ionic equation for each of the following replacement reactions:
  1. The reaction between chromium(II) sulfate and sodium hydroxide.
  1. The reaction between magnesium and titanium(IV) chloride.

Qualitative vs. Quantitative Analysis

  • Qualitative analysis involves determining by experiment whether a certain substance is present in a sample.
  • Quantitative analysis involves determining how much of a certain substance is present in a sample.
  • For aqueous solutions, typical qualitative analytic techniques include observing solution colour, flame tests and precipitation reactions.
  • When certain ions are dissolved in water, they give the solution a distinct color.

Ions / Symbol / Colour
Cations / chromium(II)
copper(II) / Cr2+(aq)
Cu2+(aq) / blue
chromium(III)
copper(I)
iron(II)
nickel(II) / Cr3+(aq)
Cu+(aq)
Fe2+(aq)
Ni2+(aq) / green
iron(III) / Fe3+(aq) / pale yellow
cobalt(II)
manganese(II) / Co2+(aq)
Mn2+(aq) / pink
Anions / chromate / CrO42-(aq) / yellow
dichromate / Cr2O72-(aq) / orange
permanganate / MnO4-(aq) / purple
  • The more concentrated the ion is in the solution, the more evident this characteristic colour will appear.
  • Many metal ions produce a distinct colour of flame when they are heated.
  • One way to test for the presence of metal ions in solution is to heat a drop of the solution in a hot flame and observe the colour. This is called a flame test.
  • The different colours that fireworks can have are due to the explosive ignition of different metals and metal salts.

Ions / Symbol / Colour
lithium / Li+(aq) / red
sodium / Na+(aq) / yellow
potassium / K+(aq) / violet
calcium / Ca2+(aq) / yellowish-red
strontium / Sr2+(aq) / red
barium / Ba2+(aq) / yellowish-green
copper(II) / Cu2+(aq) / bluish-green

Demo: Flame Test of Various Solutions

Lab Design:

  • A variety of unidentified solutions will be provided.
  • A metal loop will be dipped into a solution and then placed into the blue flame of a Bunsen burner.
  • The colour of the flame will be observed.
  • The metal loop will be cleaned by dipping it into hydrochloric acid and then being placed into the flame. This will be repeated 3 times.
  • The remaining solutions will then be tested, with the metal loop being cleaned in between each test.

Practice: Net Ionic Equations

For each of the following chemical reactions, write a complete balanced equation and a net ionic equation.

  1. The reaction between zirconium and aluminium chlorate.
  1. The reaction between cobalt(II) sulfate and ammonium hydroxide.
  1. The reaction between gallium acetate and potassium.
  1. The reaction between magnesium bromide and sodium phosphate.

Selective Precipitation

  • Precipitation reactions can determine whether an ion is present in solution or not.
  • A precipitation reaction is another term for a double-replacement reaction in solution that produces a solid product (the precipitate).
  • Chemists add dissolved substances to unknown solutions and observe whether a precipitate forms.

  • At each stage, the colour of the solution is observed and the resulting precipitate is removed.
  • Flame tests can be used to identify the precipitates.
  • EX: You are given an unidentified solution and are told that it may or may not contain sulfide ions. How could you confirm or deny the presence of S2-(aq) in this solution?
  • EX: You are given an unidentified solution and are told that it may or may not contain zinc ions. How could you confirm or deny the presence of Zn2+(aq) in this solution?
  • If a solution contains more than one dissolved ion, it is essential to design the technique carefully so that only one precipitate is formed.
  • EX: You are given an unidentified solution and are told that it may or may not contain acetate ions and/or sulfate ions. How could you confirm or deny the presence of CH3COO-(aq) and/or SO42-(aq) in this solution?
  • EX: You are given an unidentified solution and are told that it may or may not contain lead(II) ions and/or aluminium ions. How could you confirm or deny the presence of Pb2+(aq) and/or Al3+(aq) in this solution?

Practice: Selective Precipitation

  1. What amount of calcium nitrate would be found in a 300 g sample?
  1. What volume would 650 g of dinitrogen tetroxide gas occupy in order to exert a pressure of 650 mm Hg at 212oC?
  1. Write a complete balanced equation, an ionic equation and a net ionic equation for the reaction between silver nitrate and titanium(IV) acetate.
  1. You have been given a clear, colourless solution. You do several qualitative experimental tests on the solution and get the following results: (1) the flame test shows a bit of a red colour; and (2) a precipitate is formed when ammonium sulfate is added. What cation(s) might be present in this solution? What additional test could you do that would help you to be more certain as to which cation is actually present?
  1. You have been given a solution that contains chloride ions and/or iodate ions. Design a selective precipitation procedure that would allow you to determine which of these ions (if either at all) are present.
  1. You have been given a solution that contains lead(IV) ions and/or zinc ions. Design a selective precipitation procedure that would allow you to determine which of these ions (if either at all) are present.

Stoichiometry

  • Balanced chemical equations are essential to doing calculations and making predictions related to quantities in a chemical reaction.
  • The balancing coefficients in a chemical equation illustrate the relative number of particles of each chemical involved.
  • For example, the production of nitrogen dioxide has the following balanced chemical equation:

2 NO(g) + O2(g) 2 NO2(g)

  • For every 2 molecules of NO2(g) that are produced, 2 molecules of NO(g) and 1 molecule of O2(g) have been consumed.
  • Because of the large numbers of molecules involved in any chemical reaction, it is more convenient to compare the amount of the reactants and products.
  • For every 2 mol of NO2(g) that are produced, 2 mol of NO(g) and 1 mol of O2(g) have been consumed.
  • In a reaction, the actual amounts involved may vary but this ratio will always be observed.
  • EX: Use the following reaction to determine the missing amounts.

C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O (g)

A / 1.28 mol
B / 0.38 mol
C / 6.2x10-3 mol
D / 758.3 mol
  • EX: What amount of zinc will be produced by the decomposition of 0.40 mol of zinc phosphate?
  • EX: During the formation of benzoic acid, 0.366 mol of carbon is consumed. What amount of hydrogen will have been consumed during this reaction?

Practice: Stoichiometry

For each of the following questions, write a balanced chemical equation and then use the mole ratio to answer the question.

  1. A student mixes together a solution of silver nitrate with a solution of sodium chromate and a precipitate forms. What amount of precipitate will form if the student has reacted 0.314 mol of silver nitrate?
  1. When ammonia (NH3(g)) is mixed with carbon dioxide, the products are water vapour and urea (NH2CONH2(s)). What amount of water vapour is formed when 6.00 mol of carbon dioxide has reacted?
  1. Ammonia can react with sulfuric acid to produce ammonium sulfate. What amount of ammonia is required to react fully with 3.28 mol of sulfuric acid?
  1. The decomposition of titanium(III) sulfate produces 0.914 mol of sulfur. What amount of the compound has reacted?
  1. The complete combustion of lauric acid, CH3(CH2)10COOH(s), produces carbon dioxide and water vapour. Calculate the amount of oxygen that must react if this combustion reaction produces 4.55 mol of carbon dioxide.
  1. What amount of aluminium will be able to react completely with 2.19 mol of zinc nitrate?

Gravimetric Stoichiometry

  • Gravimetric stoichiometry is the analysis of the various masses of reactants and/or products involved in a chemical reaction.
  • However, the coefficient ratio can only be used to compare amounts of chemicals.
  • For example, in the formation of carbon dioxide gas,

C(s) + O2(g) CO2(g)

it would be correct to say that 1 mol of carbon reacts with 1 mol of oxygen, but it would be incorrect to say that 1 g of carbon reacts with 1 g of oxygen.

  • Stoichiometric Process:
  • Write a balanced chemical equation.
  1. Using the information given, calculate the amount of the given substance (ngiven) by the following equation:
  1. Calculate the amount of the required substance (nrequired) using the mole ratio from the balanced equation and the ngiven you calculated in step 2.
  1. Calculate the mass of the required substance using the nrequired you calculated in step 3 by the following equation:
  • EX: What mass of oxygen must be available in order to burn 120 g of ethane (C2H6(g))?
  • EX: Solid lithium hydroxide reacts with carbon dioxide gas to produce lithium carbonate and water vapour. What mass of lithium carbonate will be produced when 5.00 g of lithium hydroxide is used up?
  • EX: Solutions of sodium bromide and lead(II) acetate are mixed together. The precipitate is filtered, dried, and found to have a mass of 2.17 g. What minimum mass of lead(II) acetate was dissolved in the original solution?

Practice: Mass Stoichiometry

  1. Write a balanced chemical equation for the formation of iron(II) thiocyanate.

a)What mass of the compound can be produced when 50.0 g of iron completely reacts?

b)What mass of the compound can be produced when 50.0 g of nitrogen completely reacts?

c)What mass of sulfur must have reacted in order to produce 75 g of the compound?

  1. Write a balanced chemical equation for the combustion of butane (C4H10).

a)What mass of oxygen is required to completely react with 30.0 g of butane?

b)What mass of carbon dioxide will be produced when 75.00 g of butane is burned?

c)What mass of water vapor will be produced during a reaction in which 3.85 g of carbon dioxide is also produced?

  1. Write a balanced chemical equation for the reaction between zinc and silver nitrate.

a)What mass of zinc is required to completely react with 14 g of silver nitrate?

b)What mass of silver can be produced when 40.0 g of zinc is completely reacted?

c)What mass of zinc nitrate will be produced if 11.0 g of silver nitrate is completely consumed?

Solution Stoichiometry

  • Reactions taking place in aqueous environments are typically between solute particles. Generally, the water solvent molecules are not involved in the reaction itself.
  • Solution stoichiometry follows that same general process as gravimetric stoichiometry except that molar concentrations and volumes can be used to calculate amounts.
  • These questions typically include the use of the following equation:
  • EX: What volume of 0.214 mol/L sodium hydroxide would be required to completely neutralize 500 mL of 0.0104 mol/L hydrochloric acid?
  • EX: A 100 mL portion of hydrochloric acid is able to react with 5.00 g of zinc. What is the concentration of the hydrochloric acid solution?
  • EX: A student mixes 225 mL of 0.078 mol/L cobalt(II) nitrate with an excess volume of sodium hydroxide. Predict the mass of precipitate that should be made.

Practice: Solution Stoichiometry

  1. Ethanoic acid, CH3COOH(l), is produced according to the following chemical equation:

CH3OH(l) + CO(g) → CH3COOH(l)

Calculate the mass of ethanoic acid that would be produced by the reaction of 6.0 x 106 g of CO(g) with sufficient CH3OH(l)

  1. Sulfuric acid can be neutralized by reacting it with potassium hydroxide. What volume of 0.676 mol/L sulfuric acid can be neutralized by 41.7 mL of 0.442 mol/L potassium hydroxide?
  1. When solutions of lead(II) nitrate and sodium iodide are mixed, a bright yellow precipitate appears.
  1. What volume of 0.125 mol/L sodium iodide is necessary to precipitate all the aqueous lead(II) ions in 25.0 mL of 0.100 mol/L lead(II) nitrate?
  1. What mass of precipitate is formed in this reaction?
  1. Calculate the mass of the compound produced when 323 g of Cl2(g) reacts completely during the formation of phosphorus trichloride.
  1. When heated, calcium carbonate decomposes into calcium oxide and carbon dioxide. What mass of calcium carbonate must be decomposed in order to produce 500 kg of calcium oxide?
  1. What minimum volume of 0.50 mol/L aqueous magnesium chloride to you need to add to 60 mL of 0.30 mol/L aqueous silver nitrate to remove all of the chloride ions?

Gas Stoichiometry