Transcript ChemChpt 10 2014

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
compound
2. write the number of each type of atom in
the compound
3. look up the atomic mass for each atom in
the compound
4. 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 & 0⁰C)
• 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