C.W. BAKER HIGH SCHOOL

CORE AREA III REVIEW

Moles – Stoichiometry

Prepared by M. Foster

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1.      Compounds – a substance made up of two or more different substances that are chemically combined in a fixed proportion.

a.       Formulas – the chemical composition of a compound is indicated by its formula.

                                                   i.      Subscripts – the small numbers which come after the symbols for the different elements in the formula for a compound.  They indicate the relative numbers of atoms of the different elements in a compound:

Example:  H2O indicates two hydrogen atoms are present for each oxygen atom.  Where no subscript is written, the subscript is assumed to be one.

 

                                                 ii.      Coefficients - the numbers which appear in front of the formulas in a balanced chemical equation.   The coefficients give the ratio of molecules, or moles of molecules, that are involved in a reaction.

Example:  2 H2    +     O2      --->     2 H2O

Where no coefficient is written. the coefficient is assumed to be one.

 

b.      Types of Chemical Formulas

 

                                                   i.      Molecular Formulas – indicate the type and number of atoms which make up a molecule of a substance.  Example: the formula for ethane is C2H6 because each ethane molecule contains two carbon atoms and six hydrogen atoms.

 

                                                 ii.      Empirical Formulas – a formula that tells the simplest whole number ratio of the atoms of the elements in a compound.  The empirical formula can be determined from the molecular formula by reducing it to simplest terms.  Example: the empirical formula of ethane is CH3.  It is important to note that sometimes the empirical formula and the molecular formula are the same.  Examples of this would be water (H2O) and carbon dioxide (CO2).

 

                                                iii.      Structural Formulas – a formula which indicates the type and number of atoms that make up a molecule, but also indicate how the atoms are arranged.  The structural formula for ethane is

 

c.       Formula Mass

 

                                                   i.      The sum of the atomic masses of the atoms that make up the formula.  The formula mass is numerically equal to the molar mass. 

 

                                                 ii.      The molar mass or gram formula mass (sometimes referred to as the molecular mass or molecular weight) is the mass of one mole of that substance.  The molar mass is determined by adding up the gram atomic masses (from the Periodic Table)  for the atoms in one mole of the substance. 

Example:  one mole of  water (H2O) contains two moles of hydrogen ( 2 x 1 gram/mole) + one mole of oxygen (1 mole x 16grams). 

Therefore one mole of water has a mass of 18 grams (2g + 16g).

 

                                                iii.      The Molar mass can be used to calculate the number of moles of a substance.  The formula used for this is found in Reference Table T.

 

 

 

 

 

d.      Per cent composition – the percent composition of a compound can be determined experimentally or theoretically (from the formula).  In both cases, the formula to be used is found in Reference Table T.

 

 

                                                  i.      Calculating % composition from the formula. 

Example:  What is the percent of carbon in propane (C3H8)?

Assume that you have 1 mole of the substance.  One mole of propane contains 3 moles of carbon (3moles  x 12grams/mole = 36 grams) and 8 moles of hydrogen (8moles x 1 gram/mole = 8 grams).  The total mass of the whole compound then is 44 grams.  Plugging those numbers in the formula we get:

% of carbon by mass = 36g/44g x100 = 81.8 % ~ 82%

 

                                                 ii.      Calculating % composition from experimental data.

Example: A sample of a compound is found to consist of 24 grams of carbon and 8 grams of hydrogen, what is the percent of hydrogen in the compound?

The answer can be found by plugging the numbers into the formula:

% of hydrogen by mass = 8g/32g x100 = 25%

 

2.      Balancing Chemical Equations - a shorthand way of describing what happens in chemical a chemical reaction.  The reactants are the substances that are present at the beginning of the reaction and are written on the left side, while the products are the new substances that are formed as a result of the reaction and are found on the right side.  The  products and reactants are separated by an arrow indicating the direction the reaction is going in.  A double arrow indicates an equilibrium reaction.

 

a.      In chemical reactions, three things are conserved (mass, charge, and energy).  In order to demonstrate that these things are conserved, it is necessary to write a balanced chemical equation.

 

                                             iii.      Conservation of mass - several hundred years ago, scientists discovered that the mass of the reactants was always equal to the mass of the products in a chemical reaction. Because of this, every atom that is present on the reactant side must also be present on the product side of the equation.  In order to insure that an equation is balanced, follow these two steps:

 

1.      Write the correct formulas for each of the reactants and products.  In doing this be sure to pay attention to the diatomic elements and utilize the charges/oxidation states in the Periodic Table and in the Table of Polyatomic ions to make sure that the formulas are correctly written.  Most incorrectly balanced equations result from failing to do this step correctly.

 

2.      After all formulas are correctly written, adjust the coefficients (the numbers in front of the formulas) to make sure that there are the same number of each type of atom is present on the reactant and product sides.

 

                                                     iv.      Conservation of charge - the sum of the charges on the reactant side of an equation must equal the sum of the charges on the product side.  The equation can be balanced by adjusting the coefficients to make sure that the charges add up.

Example:     

Cu     +     Ag+     --->    Cu+2      +   Ag

This equation is NOT balanced because the sum of the charges on the left (0 +1) does NOT equal the sum of the charges on the right side (+2 + 0).

Adjust the coefficients to

Cu     +     2 Ag+     --->    Cu+2      +  2 Ag

Now the charges on each side are equal (both sides equal +2).

 

 

 

                                                   v.         Conservation of Energy - In chemical reactions, the potential energy of the reactants and products usually change.

 

1.      When high energy reactants go to lower energy products energy is given off.  This type of reaction is called an exothermic reaction. 

Example:  

high energy reactants ----> low energy products  + energy

(Note: the energy term is on the product side)

 

2.      When low energy reactants go to higher energy products energy must be supplied.  This type of reaction is called an endothermic reaction.

Example:  

low energy reactants + energy---->high energy products

(Note the energy term is on the reactant side)

 

3.      Using balanced chemical equations

a.      A chemical recipe - balanced chemical equations are much like recipes in cooking.  A cooking recipe tells what ingredients must be used and also the ratios of the ingredients.  A simple recipe for cooking rice would be:

 

2 cups water  +    1 uncooked  cup rice  ---> 2 cups of cooked rice

 

It is important to realize that the recipe doesn't just tell us how to make 2 cups of rice, but by utilizing the ratios, almost any amount!  For instance, to make 12 cups of cooked rice you would need 6 cups of uncooked rice.

 

b.      The coefficients provide the mole ratios - in a balanced equation the coefficients indicate the mole ratios of all of the products or reactants.  This is extremely useful, because if we know how many moles we have of any product or reactant, we can find the number of moles of all of the rest of the reactants.

Example:   How many moles of oxygen are required to react with 12 moles of hydrogen when water is formed according to the equation:

 

2 H2    +    O2   ------->    2 H2O

 

The ratio of hydrogen to oxygen in any case must be the same as the ratio given by the equation in the balanced equation.  Since the ratio of the coefficients is 2 to 1, the ratio of the reactants in the specific case must be the same.  To achieve the required 2:1 ratio, 6 moles of oxygen will be required to react (12:6 as 2:1)

 

c.       Titrations – a very useful application of the stoichiometry of balanced equation is the titration.  In this lab procedure, a measured volume of an acid of known concentration is added to a known volume of a base whose concentration is not known (or vice versa).  The concentration of the unknown base can then be calculated.  The device used to accurately measure out the volume of the solution being added is called a buret.

 

                                                  i.      Molarity – molarity is a unit of concentration that indicates the number of moles of solute which are present in one liter of solution.  Molarity can be found by in Table T.

 

                                                ii.      Titrations – in reactions involving acids and bases that are in a 1:1 ratio, the equation found in Table T may be used to simplify the math.


4.      Types of chemical reactions

 

a.      Synthesis - a reaction in which two simple substances come together to form a more complicated substance.  Sometimes this reaction is represented as

A + B ----->  AB

Examples:   

2Na    +   Cl2   ----->  2 NaCl

C   +   O2   ------->  CO2

 

b.      Decomposition - a reaction in which a more complex substance breaks down into simpler substances.  (The opposite of synthesis.)

AB  ------>   A  +    B

Example:

2 H2O   ------->  2 H2   +   O2

 

c.       Single replacement - a reaction in which a metal in a compound is replaced by a more active metal (see Table J to determine activities).

A  +  BC  ------>  AC   +  B

Examples:

Zn  +  CuSO4  ---->   ZnSO4   +  Cu

Mg  + 2  HCl   ----->  MgCl2    +    H2

 

Also, sometimes a more active non-metal may replace the non-metal in the compound.

Example:                   

Cl2   +   2 NaBr   ----> 2  NaCl   +   Br2

 

 

d.      Double replacement - a reaction in which two ionic substances switch their "partners" (sometimes compared to two couples dancing and then switching partners).

AB   +   CD  ---->  AD    +    CD

Examples:

NaOH    +    HCl   ------>   NaCl    +   H-OH

 

e.      Combustion - a reaction in which something reacts with oxygen to produce energy (exothermic reaction).

 

                                                  i.      Complete combustion - sufficient oxygen is available to completely oxidize the reactants.  Common products are carbon dioxide and water.

Example:   

Methane +  oxygen ------> carbon dioxide   + water

(CH4    + 2 O2  ---->   CO2   +  2 H2O )

 

                                                ii.      Incomplete combustion -  less oxygen is available and carbon monoxide is produced with water.

Example:   

Methane  +  oxygen  -----> carbon monoxide  + water

(2 CH4   +  3  O2 ------>  2 CO  +   4 H2O )

 

Notice the only difference between complete and incomplete is the ratio of methane and oxygen (1:2 for complete vs. 2:3 for incomplete).  The ratio determines which type of combustion will occur.