Honors Chemistry Unit 7: Chemical Reactions

Evidence of Chemical Reactions

When a chemical reaction occurs, new substances are formed from existing atoms. The substances that are present at the start are rearranged to form new substances that are made of the same number and types of atoms as the original substances. In other words, the atoms that are present at the start are also present at the end. (Atoms cannot be created or destroyed by a chemical reaction.) Chemical reaction are often accompanied by a visual signal.

Examples:

Combination / Chemical Reaction (yes/no) / Evidence
potassium iodide solution + lead nitrate solution
battery in water
ammonium chloride + water
magnesium + oxygen
calcium chloride + water
zinc + hydrochloric acid
hydrogen + oxygen

Representing a Chemical Reaction

A chemical reaction occurs when solutions of potassium iodide and lead (II) nitrate are mixed to produce a solid precipitate of lead (II) iodide and a solution of potassium nitrate.

The chemical reaction described above can be easily represented with a chemical equation. A chemical equation uses symbols to represent what happens during a chemical reaction. The equation for the reaction described above is:

Parts of a Chemical Equation:

  1. Reactants
  2. Products
  3. Subscripts
  4. Coefficients
  5. State symbols

Often a catalyst is required to make a reaction occur. A catalyst is a substance that speeds up a chemical reaction without being used up. Since it is not technically a part of the reaction, the catalyst is written above the “yields” arrow in the equation.

In biological systems catalysts are called enzymes

Diatomic Elements:

Certain elements exist in nature as diatomic molecules. It is important to remember these when writing and balancing equations.

Balancing Equations:

Since atoms cannot be created or destroyed in chemical reactions, it is important that we balance chemical equations. A balanced equation show the ratio of substances that are needed to ensure that the same number and types of atoms are present before and after the reaction occurs.

Equations are balanced by using coefficients. NEVER change a subscript when balancing an equation! Changing a subscript changes the identity of the substance.

Example:

H2 (g) + O2 (g)  H2O(g)

Correct:______

Incorrect:______

Guided Practice 1: Balancing Equations

  1. K(s) + H2O(l)  H2(g) + KOH(aq)
  2. C2H5OH(l) + O2(g)  CO2(g) + H2O(g)
  1. Solid potassium reacts with liquid water to form gaseous hydrogen and potassium hydroxide that dissolves in the water.
  1. Under appropriate conditions at 1000oC, ammonia gas reacts with oxygen gas to produce gaseous nitrogen monoxide (common name, nitric oxide) and gaseous water.

Independent Practice 1

Propane, C3H8, a liquid at 25oC under high pressure, is often used for gas grills and as a fuel in rural areas where there is no natural gas pipeline. When liquid propane is released from its storage tank, it changes to propane gas that reacts with oxygen gas (it burns) to give gaseous carbon dioxide and gaseous water. Write a balanced equation for this reaction.

Guided Practice 2

Glass is sometimes decorated by etching patterns on its surface. Etching occurs when hydrofluoric acid (an aqueous solution of HF) reacts with the silicon dioxide in the glass to form gaseous silicon tetrafluoride and liquid water. Write and balance an equation for this reaction.

Independent Practice 2: Write balanced chemical equations for the following reactions.

  1. When solid ammonium nitrate is heated, it produces nitrogen gas and water vapor.
  1. Gaseous nitrogen monoxide (common name nitric oxide) decomposes to produce dinitrogen monoxide gas (common name, nitrous oxide) and nitrogen dioxide gas.
  1. Liquid nitric acid decomposes to reddish-brown nitrogen dioxide gas, liquid water, and oxygen gas. (This is why bottles of nitric acid become yellow upon standing.)

Homework Practice 1: Complete on notebook paper.

Write complete balanced equations for the reactions described below.

  1. When iron wire is heated in the presence of sulfur, the iron soon begins to glow, and a chunky, blue-black mass of iron (II) sulfide is formed.
  2. Pennies in the United States now consist of a zinc disk that is coated with a thin layer of copper. If a penny is scratched and then soaked in hydrochloric acid, it is possible to dissolve the zinc disk, leaving only a thin, hollow shell of copper. Write the equation to illustrate the reaction of zinc metal with hydrochloric acid, which produces dissolved zinc chloride and evolves gaseous hydrogen.
  3. When small amounts of ammonia gas are needed, they can be generated by the reaction of an ammonium salt with a strong base. For example, if ammonium chloride is heated with sodium hydroxide, ammonia gas, water vapor, and sodium chloride are produced. Write the equation for this reaction.
  4. Acetylene gas (C2H2) often used by plumbers, welders, and glass blowers because it burns in oxygen with an intensely hot flame. The products of the combustion of acetylene are carbon dioxide and water vapor. Write the balanced equation for this reaction.
  5. Pure silicon, which is needed in the manufacturing of electronic components, may be prepared by heating silicon dioxide (sand) with carbon at high temperatures, releasing carbon monoxide gas. Write the equation for this reaction.

Quiz 7.1

Reactions in Aqueous Solutions:

In order of chemical reactions to occur, reactants must come into contact with one another. This usually requires that one or more of the reactants be in a fluid state. That is at least one reactant must be a gas, a liquid, or an aqueous solution.

What “Drives” a Reaction to Occur?

  1. Formation of a solid
  2. Formation of water
  3. Transfer of electrons
  4. Formation of a gas

Formation of a Solid (Precipitation Reactions)

Reactions in which a solid is formed from the combination of aqueous solutions are called precipitation reactions. The solid that “falls” out of the solution is called the precipitate.

A precipitate is an insoluble salt that is formed from two soluble salts. When a salt (ionic compound) is soluble, it breaks apart in solution to form ions.

It is important to understand which ions will be present when a particular salt is dissolved in water.

Examples:

NaCl(aq) contains Na+ and Cl- ions

NaNO3(aq) contains Na+ and NO3- ions

Independent Practice 3

Identify the ions present in aqueous solutions of the following salts:

  1. CaF2(aq)
  2. K2SO4(aq)
  3. NH4Cl(aq)
  4. LiOH (aq)
  5. Na3PO4(aq)

When two aqueous solutions are mixed, we must examine the possible combinations of ions to determine if an insoluble salt (precipitate) is likely to form.

Example: Solutions of potassium chromate and barium nitrate are mixed

To determine if a salt is soluble, we need to refer to the solubility rules. Soluble salts are those that readily dissolve in water. Insoluble salts are those that will not dissolve at all or will only dissolve slightly in water.

Solubility Guidelines

Soluble Compounds / Exceptions
Alkali Metals / None
Ammonium / None
Halides / Those containing: Hg, Ag, Pb
Sulfates / Those containing: Ca, Ba, Sr, Pb, Hg
Nitrates / None
Insoluble Compounds / Exceptions
Sulfides / Those containing: Alkali metals or ammonium, or Ca, Ba, Sr
Carbonates / Those containing: alkali metals or ammonium
Phosphates / Those containing: alkali metals or ammonium
Hydroxides / Those containing: Alkali metals or ammonium, or Ca, Ba, Sr

Guided Practice 4

Determine the solid precipitate that forms in the following reactions:

  1. Silver nitrate solution mixes with potassium chloride solution
  1. Potassium nitrate solution mixes with barium chloride solution
  1. Sodium sulfate solution mixes with lead (II) nitrate solution
  1. Potassium hydroxide solution mixes with iron (III) nitrate solution

Independent Practice 4

Determine the solid precipitate that forms in the following reactions:

  1. Barium nitrate solution mixes with sodium chloride solution
  1. Sodium sulfide solution mixes with copper (II) nitrate solution
  1. Ammonium chloride solution mixes with lead (II) nitrate solution

Ionic Equations:

A molecular equation shows the complete formulas for all reactants and products in the chemical reaction. So far, the equations that we have written have been molecular equations. When working with precipitation reactions, however, it is sometimes useful to use ionic equations.

An ionic equation represents the substances involved in a chemical reaction as separate ions in solution. In a complete ionic equation all substances that are strong electrolytes are represented as ions (not molecules). Strong electrolytes include all soluble ionic compounds, and strong acids and bases (more on those later).

For example: When aqueous solutions of potassium chromate and barium nitrate are mixed a reaction occurs to form solid barium chromate and dissolved potassium nitrate.

Molecular Equation:

Complete Ionic Equation:

From the complete ionic equation in the example above, we can see that not all ions actually participate in the reaction. Some ions are present in solution before and after the reaction. These are called spectator ions. Spectator ions can be cancelled out to give the net ionic equation:

Net Ionic Equation:

Guided Practice 5

Write molecular, complete ionic, and net ionic equations for the following reactions:

  1. Aqueous solutions of lead (II) nitrate and sodium sulfate are mixed.
  1. Aqueous sodium chloride is added to aqueous silver nitrate.
  1. Aqueous potassium hydroxide is mixed with aqueous iron (III) nitrate.

Independent Practice 5

Write molecular, complete ionic, and net ionic equations for the following reactions:

  1. Aqueous sodium sulfide is mixed with aqueous copper (II) nitrate.
  1. Aqueous ammonium chloride and aqueous lead (II) nitrate react.

Lab 7.1

Homework Practice 2: Complete on notebook paper

Write molecular, complete ionic, and net ionic equations for the reactions that occur when aqueous solutions of the given compounds are mixed.

  1. Silver nitrate and potassium chloride
  2. Nickel (II) sulfate and barium chloride
  3. Ammonium phosphate and calcium chloride
  4. Calcium chloride and ammonium sulfate
  5. Lead (II) nitrate and barium chloride

Quiz 7.2

Reactions that Form Water (Acid-Base Reactions):

While there are several different accepted definitions of acids, for our purposes we will focus on the Arrhenius definition: An acid is a substance that produces hydrogen ions (H+) when dissolved in water.

Strong acids (like soluble ionic compounds) are strong electrolytes and will completely dissociate into ions when placed in water. There are 7 strong acids.

Arrhenius defined a base as a substance that produces hydroxide ions (OH-) in water. The strong bases are then the soluble metal hydroxides.

Example: Reaction of hydrochloric acid with sodium hydroxide

When solutions of strong acids and bases are mixed, the products are a salt and water.

The net ionic equation for the reaction of a strong acid and strong base:

Guided Practice 6

Write the molecular, complete ionic and net ionic equation for the reaction between aqueous solutions of nitric acid and potassium hydroxide.

Reactions of Metals with Nonmetals (Reduction-Oxidation Reactions)

A reaction that involves a transfer of electrons is called a reduction-oxidation (or redox) reaction.

The substance that loses electrons is oxidized while the substance that gains electrons is reduced. (OIL RIG)

Example: Consider the reaction of magnesium metal with oxygen gas.

Another example is the reaction of aluminum with solid iron (III) oxide. (This reaction is often called the thermite reaction.)

Guided Practice 7

For each of the reactions below show how electrons are gained or lost and identify the substance that is oxidized and the substance that is reduced.

  1. 2Al(s) + 3I2(s)  2AlI3(s)
  1. 2Cs(s) + F2(g)  2CsF(s)

Independent Practice 7

For each of the reactions below show how electrons are gained or lost and identify the substance that is oxidized and the substance that is reduced.

  1. 2 Na(s) + Br2(l)  2NaBr(s)
  1. 2Ca(s) + O2(g)  2CaO (s)

Classification of Reactions

Most reactions can be classified as one or more type of reaction based on the reactants and products.

Synthesis Reactions

  • Two or more reactants combine to form one product
  • General equation is: A + B  AB

Decomposition Reactions

  • One reactant breaks apart to form two or more products
  • Often requires a catalyst
  • General equation is: AB  A + B

Single Replacement Reactions

  • A single element replaces another element from a compound
  • These reactions are also redox reactions
  • General equation is: A + BX  B + AX

Double Replacement Reactions

  • Two ions exchange places
  • General equation is: AY + BX  AX + BY

Combustion Reactions

  • Reactions that involve oxygen and produce heat so quickly that a flame results
  • Common combustion reactions involve hydrocarbons
  • Combustion of hydrocarbons produce CO2 and H2O

Important “Things” to Remember

When carbonic acid is produced in a reaction, it immediately breaks into carbon dioxide and water.

Metal carbonates decompose to produce the metal oxide and carbon dioxide.

Guided Practice 8

Classify the following reactions. Some may be classified more than one way.

  1. 2K(s) + Cl2(g)  2KCl(s)
  2. Fe2O3(s) + 2Al(s)  Al2O3(s) + 2Fe(s)
  3. 2Mg(s) + O2(g)  2MgO(s)
  4. HNO3(aq) + NaOH(aq)  H2O(l) + NaNO3(aq)
  5. KBr(aq) + AgNO3(aq)  AgBr(s) + KNO3(aq)
  6. PbO2(s)  Pb(s) + O2(g)

Independent Practice 8

  1. 4NH3(g) + 5O2(g)  4NO(g) + 6H2O(g)
  2. S8(s) + 8O2(g)  8SO2(g)
  3. 2Al(s) + 3Cl2(g)  2AlCl3(s)
  4. 2AlN(s)  2Al(s) + N2(g)
  5. BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2NaCl(aq)
  6. 2Cs(s) + Br2(l)  2CsBr(s)
  7. KOH(aq) + HCl(aq)  H2O(l) + KCl(aq)
  8. 2C2H2(g) + 5O2(g)  4CO2(g) + 2H2O(l)

Homework Practice 3: Complete on notebook paper

Write balanced equations for the following reactions. Classify each in as many different ways as possible.

  1. Solid tetraiodine nonoxide produces solid diiodine hexoxide, solid iodine, and oxygen gas.
  2. Solid magnesium reacts with a solution of silver nitrate to produce a solution of magnesium nitrate and solid silver.
  3. Liquid silicon tetrachloride reacts with solid magnesium to produce solid magnesium chloride and solid silicon.
  4. Solutions of copper (II) chloride and silver nitrate react to produce solid silver chloride and a solution of copper (II) nitrate.
  5. Solid aluminum reacts with liquid bromine to produce solid aluminum bromide.

Quiz 7.3

Lab 7.2

Test Review Homework: Complete on notebook paper

  1. The element carbon undergoes many inorganic reactions, as well as being the basis for the field of organic chemistry. Write balanced chemical equations for the reactions of carbon described below.
  2. Carbon burns in an excess of oxygen (for example in the air) to produce carbon dioxide.
  3. If the supply of oxygen is limited, carbon will still burn, but will produce carbon monoxide rather than carbon dioxide.
  4. If molten lithium metal is treated with carbon, lithium carbide, Li2C2, is produced.
  5. Iron (II) oxide reacts with carbon above temperatures of about 700oC to produce carbon monoxide gas and molten elemental iron.
  6. Carbon reacts with fluorine gas at high temperatures to make carbon tetrafluoride.
  7. Balance the following equations:
  8. C2H6(g) + O2(g)  CO2(g) + H2O(g)
  9. C5H10(g) + O2(g)  CO2(g) + H2O(g)
  10. Cl2O7(g) + H2O(l)  HClO4(aq)
  11. H2S(g) + Cl2(g)  S8(s) + HCl(g)
  12. HCl, HNO3, and H2SO4 are strong acids while NaOH and KOH are strong bases. Write complete molecular equations for the neutralization reaction between each of these acids with each of these bases.
  13. Balance the following equations and classify each in as many ways as possible:
  14. FeO(s) + HNO3(aq)  Fe(NO3)2(aq) + H2O(l)
  15. Mg(s) + CO2(g) + O2(g)  MgCO3(s)
  16. HI(aq) + KOH(aq)  KI(aq) + H2O(l)
  17. C12H22O11(s)  C(s) + H2O(g)
  18. B(s) + O2(g)  B2O3(s)
  19. Write balanced net ionic equations for reactions that occur when aqueous solutions of the following substances are mixed. If not reaction occurs, indicate so.
  20. Barium nitrate and hydrochloric acid
  21. Barium nitrate and sulfuric acid
  22. Silver nitrate and hydrochloric acid
  23. Lead (II) nitrate and sulfuric acid
  24. Iron (II) sulfate and sodium hydroxide
  25. Nickel (II) chloride and ammonium sulfide
  26. Magnesium chloride and sodium carbonate
  27. Lead (II) nitrate and barium nitrate
  28. Complete and balance the following equations:
  29. Pb(NO3)2(aq) + Na2S(aq) 
  30. AgNO3 (aq) + HCl(aq) 
  31. Mg(s) + O2(g) 
  32. H2SO4(aq) + KOH(aq) 
  33. BaCl2(aq) + H2SO4(aq) 
  34. Na3PO3(aq) + CaCl2(aq) 
  35. C4H10(l) + O2(g) 
  36. The solutions below are found on a shelf in a chemical supply room.

AgNO3 / NaCl / HC2H3O2 / HNO3 / H2SO4 / K2CrO4 / Ba(NO3)2 / H3PO4 / HCl / Pb(NO3)2 / NaOH / Na2CO3

Determine if it is possible to prepare the following substances using solutions on the shelf. If so, describe how.

  1. BaCrO4(s)
  2. NaC2H3O2(s)
  3. AgCl(s)

Unit 7 Test