Chemical Equations and Reactions

An interactive review for your upcoming exam!

***Any text written in pink links to additional information and activities.***

***Any text written in yellow requires you to make use of your internet handout.***




A chemical reaction is the process in which atoms present in the starting substances rearrange to give new chemical combinations present in the substances formed by the reaction.  These starting substances of a chemical reaction are called the reactants, and the new substances that result are called the products. 


Chemical reactions are all around, as-well-as inside of us.  From the fireworks we see on the 4th of July, to the digestion of this morning’s breakfast, chemical reactions are everywhere!  So how can we tell when a chemical reaction takes place?


Observable indications of a chemical reaction can include:

·       Evolution of heat or light

·       Production of gas

·       Formation of precipitate




The only way to be positively sure that a chemical reaction has occurred

 is through chemical analysis of the products!  We might not be able to observe one of the above indicators…

***Remember our “invisible” acid/base reaction???***







A chemical reaction can be accurately summarized with

 a properly written chemical equation.

A properly written chemical equation provides information about reactants, products, molecule-to-molecule relationships between reactants and products, and mole-to-mole relationships

between reactants and products.


***What are three pieces of information that a chemical equation

DOES NOT provide?***


Since there’s still so much information to be gained from a chemical equation,

we need to become experts at writing them out!

But just what is a “properly written chemical equation”?


A properly written equation is a chemical equation that…

1.  represents known facts

·        All reactants and products are identified

·        Identification through chemical analysis in lab or from reported experimental results

2.  contains correct formulas for reactants and products.



3.  and satisfies the law of conservation of mass. (Atoms can neither be created nor destroyed in ordinary chemical reactions.)

·        Same number of atoms of each element appear on each side of chemical equation


To be sure that a chemical equation obeys the law of conservation of mass,

you have to be sure that the  equation is balanced.  Practice your balancing skills using the CHEMBALANCER. 

Write out your answers on the accompanying internet handout.


***How did you do???***




Ok, well now that we’ve reviewed chemical equations, let’s see if we can

 apply these concepts in the MINI REVIEW!!!

Directions:  Write out and balance the correct chemical equations for the following

 reactions on the accompanying handout. (Don’t forget to indicate the state

 in parentheses.)

1.  When solid sodium nitrate is heated, it decomposes to give solid sodium nitrite and oxygen gas.

2.  When solid calcium phosphate and aqueous sulfuric acid solution react, aqueous phosphoric acid and solid calcium sulfate are produced.

3.  An aqueous solution of ammonium chloride and barium hydroxide is heated, and the compounds react to give off ammonia gas.  Barium chloride solution and water are also products.

4.  Gaseous ammonia reacts with oxygen gas to yield nitrogen gas and water.

5.  Copper(II) oxide is boiled in an aqueous solution of  sulfuric acid to form copper(II) sulfate and water.      








We can categorize chemical reactions according to the way in which the atoms or

molecules of the reactants form new groupings.  Many chemical reactions can be classified

as belonging to one of five main groups.



Also known as a composition reaction, a synthesis reaction is a reaction in which two or more substances combine to form a new compound.

It is represented by the general equation:

A + X à AX

 where A and X can be elements or compounds and AX is a compound.

Examples of synthesis reactions include:

·       Synthesis of sodium chloride

Na(s) + Cl(g) à

·        Synthesis of magnesium oxide

Mg(s) + O2(g) à

·        Synthesis of water

H2(g) + O2(g) à




A decomposition reaction is the reverse of a synthesis reaction and is a reaction in which a single compound undergoes a reaction that produces two or more simpler substances.

It is represented by the general equation:

AX à A + X

        where AX is the compound and and A and X can be elements or compounds.

               Most decomposition reactions take place only when energy in the form of

     electricity or heat is added.             

               Examples of decomposition reactions include:

·        Decomposition of water

H2O(l) à

·        Decomposition of lead (II) carbonate

PbCO3(s) à

·        Decomposition of sodium iodide

NaI(s) à




A single-replacement reaction, or displacement reaction, is a reaction in which one element replaces a similar element in a compound. 

It is represented by the general equation:

A + BX à AX + B


Y + BX à BY + X

where A, B, X, and Y are elements and AX, BX, and BY are compounds.

Many single-replacement reactions take place in water, and in comparison to both synthesis and decomposition reactions, the amount of energy required for a single-replacement reaction is smaller .

Examples of single-replacement reactions include:

·        Thermite reaction

Al(s) + Fe2O3(s) à

·        Reaction of solid calcium and water

Ca(s) + H2O(l) à

·        Reaction of solid lithium and chlorine gas

Li(s) + Cl2(g) à




A double replacement reaction is a reaction in which the ions of two compounds exchange places in an aqueous solution to form two new compounds. 

One of the new compounds that form are usually a precipitate, an insoluble gas that bubbles out of solution, or a molecular compound (usually water). 

The other compound which forms is usually soluble and remains in solution.

Double-replacement reactions are represented by the general equation:

AX + BY à AY + BX

where A, X, B, and Y in the reactants represent ions and products AY and BX represent ionic or molecular compounds.

Examples of double-replacement reactions include:

·        Reactions of silver nitrate with sodium chloride and silver nitrate with sodium iodide

AgNO3(aq) + NaCl(aq) à

AgNO3(aq) + NaI(aq) à

·        Reaction of sulfuric acid and sodium hydroxide

H2SO4(aq) + NaOH(aq) à

***There are two special cases which have specific names given to them.***

1.     When the reaction is between two ionic compounds and they form a precipitate, this reaction is also called a precipitation. The first examples given above show precipitation reactions.

2.     When the reaction is between an acid (any compound that forms hydrogen ions) and a base (any compound that forms hydroxide ions), water is formed as one of the products. This is called neutralization.  An example of a neutralization reaction is seen in the reaction above between sulfuric acid and sodium hydroxide.



A combustion reaction is a reaction in which a substance combines with oxygen and releases a large amount of energy in the form of heat and light.

Examples of combustion reactions include:

·        Three balloons of H2 and one of H2/O2 mixture

H2(g) + O2(g) à

·        Burning of methane (a hydrocarbon)

CH4(g) + O2(g) à

·        Burning of hexane (a hydrocarbon)

C6H14(l) + O2(g) à 

***Typically, combustion reactions occur when hydrocarbons react with oxygen to produce carbon dioxide and water, and will have the general equation 

CxHy + O2 à CO2+ H2O

where x represents the number of carbon atoms and y the number of

hydrogen atoms in the hydrocarbon

Hydrocarbons are a class of compounds that primarily consist of hydrogen and carbon.***






Another section reviewed… Time for another MINI REVIEW!!!

Directions:  Write out or complete and balance each of the following equations and

identify each as synthesis, decomposition, single-replacement, double-replacement, or combustion

1.  C3H8(g) + O2(g)  à

2.  Solid sodium reacted with water produces aqueous sodium hydroxide and hydrogen


3.  CaCO3(s)  à

4.  AgNO3(aq) + KI(aq) à

5.  Hydrogen gas reacts with iodine gas to form hydrogen iodide gas.








It is possible to arrange metals in the order of their chemical reactivity’s and thereby establish an activity series of metals.  The activity series helps to predict whether or not single replacement reactions of metals with metal ions and of metals with water and acids will occur.

Activity Series of the elements

Activity of Metals



K                 Can react with cold H2O

Ba              and acids, replacing

Sr               hydrogen





Mn             Can react with steam

Zn              and acids, replacing

Cr              hydrogen.




Ni               Can react with acids,

Sn              replacing hydrogen.



Sb             React with oxygen,

Bi              forming oxides.



Ag            Fairly unreactive.

Pt             Form oxides only

Au            indirectly.

Activity of Halogen Nonmetals








Some general trends in an activity series are listed below.  ***You have this as a handout***

  • An element will replace from a compound in aqueous solution any of those elements below it in the activity series.  The larger the interval between the elements in the activity series, the greater the tendency for the replacement reaction to occur.
  • Reactivity of metals towards other elements decreases as you go down the series.  The stability of their compounds also decreases.
  • Group I metals in the periodic table are the most reactive followed by Group II elements.
  • Any metal above magnesium replaces hydrogen from water
  • Any metal above cobalt replaces hydrogen from steam.
  • Any metal above hydrogen reacts with acids, replacing hydrogen.  

The synthesis reactions of metals with oxygen also occur more readily, the higher a metal is placed in the activity series.

  • Any metal above silver reacts with oxygen, forming oxides; those near the top react rapidly.
  • Any metal below mercury forms oxides only indirectly (i.e. not by reaction with O2).

The more active a metal, the more strongly it holds onto oxygen in an oxide and therefore, the more strongly the oxide resists decomposition into its elements upon heating.

  • Oxides of metals below copper decompose with heat alone.
  • Oxides of metals below chromium yield metals when heated with hydrogen.
  • Oxides of metals above iron resist conversion to the free metal when heated with hydrogen.

The most active metals are not likely to remain uncombined with other substances for very long.  Some are so active that they must be isolated from air when they are stored. 

·        Elements near the top of the series are never found free in nature.

  • Elements near the bottom of the series are often found free in nature.

Thus the activity series is useful because it indicates the possibility of reaction of a given metal with water, acids, oxygen, sulphur, halogens and compounds of other metals.  It also provides a good indication of the relative stability of compounds of any metal. 



Final section reviewed… one last MINI REVIEW!!!

Directions:  Using the activity series, predict whether each of the following reactions will

occur, and write the balanced equations that you predict will occur.  For those that you predict won’t occur, simply write NR (for no reaction).

1.  Pb(s) + ZnCl(s) à

2.  Cl2(g) + KBr(aq) à

3.  Al(s) + Pb(NO3)2(aq) à

4.  Cu(s) + FeSO4(aq) à

5.  Ni(s) + O2(g) à




Well, that concludes the internet review!

The accompanying internet review handout must be handed in on the day of the exam!

And don’t forget to study hard for the exam … Go over your quiz, your homework assignments, and the review handout!  And come to class with any questions

you might still have!