INTRODUCTION  TO    CHEMICAL  REACTIONS  

Balancing  Chemical  Equations  

Introduction  to  Chemical  Reactions  

§ A chemical reaction takes place when two or more chemicals mix and a new compound is formed.

§ Burning a piece of wood shows us an example of this.

Introduction  to  Chemical  Reactions  

Wood  that  is  burned  turns  into  ash,  soot,  water  vapor,  and  carbon  dioxide  

Introduction  to  Chemical  Reactions  

§ A scientist by the name of Antoine Lavoisier discovered that matter cannot be created, nor can it be destroyed.

§ He  named  this  the  law  of  conservation  of  mass.  

Introduction  to  Chemical  Reactions  

§ Basically what this means is that the amount of material remaining after a reaction takes place must equal the amount that was present before the reaction.

Introduction  to  Chemical  Reactions  

Starting  material  

Ending  material  

Introduction  to  Chemical  Reactions  

§  This law is not easily observed in our everyday lives.

§  For example – a log that has burned seems to have “lost” mass (the leftover ash is a lot smaller than the original log).

§  This is because much of the material escaped into the air.

Introduction  to  Chemical  Reactions  

§  If all the smoke, ash, and water vapor could be collected; what we would find is that it would equal the original weight of the unburned log.

Writing  Chemical  Equations  

§  In chemistry a reaction is represented by a chemical equation.

§  The basic formula for a chemical equation is made of two parts.

§  The REACTANT is the starting material. §  The PRODUCT is the ending material. §  An ARROW between the two represents

the reaction taking place.

Writing  Chemical  Equations  

Reactant   Product  

The  arrow  stands  for  “yields”  or  “produces”  

Writing  Chemical  Equations  

§ Here is an example of a chemical reaction:

H2  +  O2     H2  O    

§ However…what  do  you  notice  if  you  count  up  the  number  of  atoms  on  each  side  of  the  arrow?  

An  oxygen  atom  seems  to  have  been  destroyed!  

Writing  Chemical  Equations  

§  In the real situation, no atoms were lost. The equation was just written incorrectly.

§  There are three rules to writing chemical equations in order to make sure all the atoms are balanced correctly.

Balancing  Chemical  Equations  

§ Rule 1 – the subscripts of molecules cannot be changed (otherwise you make a completely new molecule)

§ Example: CO2 and CO are both made of carbon and oxygen, but one is a relatively harmless byproduct of your body and the other is a potentially lethal gas!

Balancing  Chemical  Equations  

§ Rule 2 – You can only add more atoms by putting coefficients in front of the entire molecule.

§ Example: In order to get four oxygen atoms out of CO2, you cannot write CO4.

§  Instead, you must write 2CO2.

Balancing  Chemical  Equations  

§ Writing 2CO2 means you have two carbon dioxide molecules:

CO2 + CO2 § A total of two carbon atoms and four

oxygen atoms.

Balancing  Chemical  Equations  

§ Rule 3 – Continue adding coefficients until all the atoms are balanced.

Balancing  Chemical  Equations  

§  Lets go back to our first equation:

H2  +  O2     H2  O    

Balancing  Chemical  Equations  

§  Try another one:

N2  +  H2     NH3      

Practice  Problems  

H2  +  Cl2     HCl      

Practice  Problems  

Zn  +  HCl     ZnCl2  +  H2      

Practice  Problems  

Na  +  OH     Na2O  +  H2