Chemistry Chapter 10

Formula Based Calculations

• a mole is 6.02 x 1023 particles like a dozen is 12 particles

• it is a large number, because we are counting particles that are so small we have to have enough of them for use in the lab

• 6.02 x 1023 is called The Avogadro Number• one mole of a substance has a mass

equivalent to its atomic mass in grams – examples: one mole of carbon atoms has a

mass of 12 grams; one mole of sulfur atoms has a mass of 32 grams (their atomic masses from the periodic chart in grams instead of amu)

Formula Based Calculations

• Molecular and Formula Mass– a list of atomic masses is in the back of

the book, remember the unit for atomic mass is the atomic mass unit and it is represented by u

– if we take the atomic masses for all of the elements in a compound, and add them together, we come up with the mass of that compound

– this mass is referred to as the molecular mass or the formula mass

• for molecules (non-ionic) substances we use the term molecular mass to indicate the sum of the atomic masses of all of the elements in the molecule

• for ionic compounds, we use the term formula mass to describe the sum of the atomic masses for all of the atoms in the formula (the simplest ratio of ions)

• the procedure for finding these masses is the same for either type

• to find the molecular/formula mass:1. write down all of the atoms in the

compound2. write the number of each type of atom in

the compound3. look up the atomic mass for each atom in

the compound4. multiply the number of atoms for each

element in the compound by its atomic mass

5. add all of the resulting atomic masses together

• remember, one mole of a substance has a mass in grams equivalent to that of one particle of the substance in atomic mass units

• because of this relationship, we may use another set of units in place of amu: grams/mole

• the mass of one mole of molecules, atoms, ions, or formula units is called the molar mass of that substance

• Converting between Grams and Moles and Atoms

– knowing the relationship that one mole of a substance has a mass in grams equivalent to that of one particle of the substance in atomic mass units, we can use our factor label method to convert between grams and moles

1. procedure for grams/moles conversions

2. determine the mass of the compound

3. set up the conversion factor by placing the “formula” mass in grams over 1 mole

4. follow the procedure for the factor label method shown previously to assure the cancellation of the unwanted units resulting in the appropriate units

• knowing the relationship that one mole of a substance contains 6.02 x 1023 particles allows us to convert between moles and particles

– procedure for grams to particle conversions

1. use the relationship of 6.02 x 1023 particles in one mole

2. Factor label method

• STP stands for standard temperature and pressure (101.3kpa & 0 C)⁰

• 1 mole of any gas at STP occupies a volume of 22.4L

• This referred to as the molar volume

• We can use this relationship in the factor label method to convert between moles and L

• Percent Composition– tells us the relative mass that each element

contributes to the total mass of the compound– the percentages of all the elements present

should add up to 100%– Procedure

1. find the “formula” mass of the compound (amu)

2. find the mass for the appropriate number of each element in the compound individually

3. divide each individual mass by the total mass and multiply by 100% to get the percent composition for each individual element present

• Empirical Formula– the elements in a compound combine in

specific ratios of whole numbers • (ex. 1 to 1, 1 to 3), the moles of the atoms for each

element are in this same whole number ratio

– to find the empirical formula from grams

1. find the grams of each element in the compound

2. calculate the number of moles of each element in the compound

3. then find the simplest ratio of the number of moles of each element (this will tell you the relative numbers of each element in the formula)

• to find the empirical formula from percent composition

1. when you are given the percent of each element in the compound, assume that you have 100 grams of that compound and your grams will be equal to the percentage for that element in the compound

– (ex. If you r compound is 51% O, then you would have 51g of O)

2. then follow the procedure from grams

• Molecular Formula

– formulas for ionic compounds show the simplest whole number ratio, so their empirical formula is their correct formula

– formulas for molecular compounds on the other hand, can be different from their empirical formulas

– some molecular compounds do have the same molecular formula as their empirical formula; but for most, the molecular formula is a whole number multiple of its empirical formula

• we can determine the molecular formula of a molecule if we are given its empirical formula and its molar mass

– Procedure:

1. find the molar mass of the empirical formula

2. divide the given molar mass for the substance by the molar of the empirical formula

3. multiply the empirical formula by the number found in step 2

• Hydrates– some compounds have their ions attached to

water molecules, they are called hydrates– these compounds usually contain a specific

ratio of water to compound– they are named the same way we learned,

then add the name for the number of waters present

• ex. CuSO4 x 5H2O is called, copper (II) sulfate pentahydrate

– the prefixes for the numbers of water are: mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

• Procedure for finding the formula for Hydrates

1. given the mass of the hydrate and the mass of the dry sample, find the difference of these 2 masses and that gives you the mass of the water.

2. then convert the mass of the water and the dry sample into moles

3. find the simplest ratio (just like in the empirical formula calculations), this tells you how many moles of compound versus water and we can now write the formula

Chemical Equations

• a chemical reaction is the process by which 1 or more substances are changed into 1 or more different substances

• they are represented by an equation using symbols and formulas

• they show what change takes place, and the relative amounts of the various elements present in the compounds involved

• The starting substances are called the reactants, and the ending substances are called the products

• the arrow means “yields” or makes or produces

• Equations are like recipes they tell you what ingredients you have and how much of each ingredient

• they can also show special instructions – letters in parentheses indicate the state of

that substance• ex’s. (g) = gas, (l) = liquid, (cr) = solid,

(aq) = aqueous which means the substance is dissolved in water

– they can show if heat, energy or catalysts are needed

– we don’t use these all of the time, only if they are given

Writing and Balancing Equations• write a word equation first, then change it

to the chemical equation • steps

– determine the reactants and the products

– assemble the parts of the equation• ? + ? ?• the symbols and formulas must be

correct – balance the equation

To balance an equation• make sure you have equal numbers of all

elements on both sides of the arrow, because mass can not be lost or gained in a reaction

• don’t confuse subscripts with coefficients• subscripts determine the substance,

coefficients tell us how much of that substance we have

• we don’t change the subscripts, or we change the substance, rather we change or add coefficients to adjust amounts

• Polyatomic ions can be balanced as a unit

Types of Reactions

• if we understand the basic types of reactions, we can better predict the products that will form from certain reactants

• we will look at the 5 basic reaction types here, there are others but we are not concerned with them

• The 5 types are: single displacement, double displacement, decomposition, synthesis, and combustion

displacement – an element replaces another in a compound

• single displacement – one element is replaced by another

– a single element reacts with a compound and displaces one of the elements from the compound and takes its place in the compound

– We can use an Activity Series to helps us predict which elements will replace which

Activity Series

– elements are arranged in order of activity with the most active on top.

– An element can displace those listed below it, but not those above it

• double displacement – the positive and negative ions of 2 compounds switch

– parts of 2 compounds switch places

– water is sometimes written HOH

Decomposition - “break down”

• a substance breaks down into simpler substance

• only one reactant

Synthesis – “put together”

• 2 or more substances form together to make one new substance

• only one product (binary or ternary)

Combustion

• Usually;

– a hydrocarbon + oxygen - carbon dioxide + water

• can be alcohols also

• the substances burn in air

• usually used to generate energy

• to determine reaction type

– Look at the products and the reactants :

• one product = synthesis

• one reactant = decomposition

• hydrocarbon + oxygen forms carbon dioxide and water = combustion

• 1 element + 1 compound forming a different element and a compound = single displacement

• 2 compounds reacting to form 2 new compounds = double displacement