The Law of Conservation of Mass:
The total mass of the substances does not change during a chemical reaction. The
number of substances might change, their properties might change, but the total amount
of matter must remain constant. The moral is: matter can be neither created nor
destroyed. The law of conservation of mass can be explained as : atoms are neither
created nor destroyed during a chemical reaction, they just rearrange to form a new
substance.
C5H12
(l) +
8O2 (g)
5 CO2 (g) +
6 H2O (l)
We took a hydrocarbon (C5H12) known as pentane, added some oxygen to it and with an
input of energy we generate CO2 gas and water.
The Law of Definite Proportion/Constant Composition:
Constant composition refers to the elemental composition by mass of a given compound
being the same for all samples of that compound. This means that every, every, every
time water is decomposed, there are 88.8 grams of oxygen present for every 11.2 g of
hydrogen. And a 100-gram sample of NaCl always contains 39.3 grams of Na and 60.7 g
of Cl.
How does this work?
NaCl is made up of 1 Na and 1 Cl.
The molar mass of Na = 22.9
The molar mass of Cl = 35.4
Thus Na =
22.9
22.9  35.4
*
100 = 39.3 grams
Constant composition/definite proportion is explained by assuming that atoms combine
in fixed ratios when they form a compound. Thus, if one oxygen atom combines with 2
hydrogen atoms to form water, then all samples of water must have the same
composition
The Law of Multiple Proportions:
In different compounds containing the same elements, the masses of one element
combined with a fixed mass of the other element are in the ratio of the smallest whole
numbers.
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Consider SO2 and SO3. Both compounds contains only sulfur and oxygen atoms. The
formulas suggest that for a given mass of sulfur, the masses of oxygen in SO 3 compared
to SO2 will be in a 3/2 ratio. Other examples include CO and CO2, FeCl2 and FeCl3, UF3
UF4 and UF6.
How does this work?
SO2
molar mass of S = 32.06
molar mass of O = 15.99
grams of S in a 100 g sample = 50.1
grams of O in a 100 g sample = 49.9
gramsO
=1
gramsS
SO3
molar mass of S = 32.06
molar mass of O = 15.99
grams of S in a 100 g sample = 40.1
grams of O in a 100 gram sample = 59.9
gramsO
= 1.5
gramsS
Since we do not want decimals (remember we are looking for the ratio of the smallest
whole numbers!!) for ratios we must get rid of the 0.5 and turn it into a whole number. If
we multiply a 0.5 by 2 we get 1 (for 0.33 multiply by 3 to get the new whole number, for
0.25 multiply by 4, for 0.75 multiply by 4 etc . . .)
So – comparing the O in SO3 to the O in SO2 for a given 1 atom of sulfur results in 3 O
atoms (1.5*2) to 2 O atoms (1*2). REMEMBER!! Whatever we do to one ratio number we
must mathematically do to the other. Therefore a 1.5:1 ratio is equal to 3:2 ratio if we
want the answer in only whole numbers (which we DO!!)
It is very important to begin thinking in terms of moles when examining chemical reactions as it
clarifies the amount of each substance present. Comparing molar masses does not clarify the
ratio of each substance present. Thinking in terms of masses does nothing but show us that the
mass during the chemical reaction will be conserved (the grams at the beginning will be the
same number of grams at the end of the chemical reaction). For example, consider the following
relationships:
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2.016 g of H2 + 38.00 g F2  40.02 g HF
1 mole of H2 + 1 mole F2  2 mole HF
This information shows that equal amounts of H2 and F2 molecules combine to form twice as a
large a population of HF molecules.
This chemical equation is made up of chemical formulas that express the identity and quantity
of substance that undergo a chemical or physical change. These chemical equations are like
sentences, which means you need to KNOW the names of the chemical species that are coming
together and undergoing a reaction. The left side of the equation shows the chemical species
that are coming together and doing the reacting, thus they are called the reactants. The
chemical species that we end up with at the end are called the products. The reactant side
shows the amount of substance present before the change and the right side shows the amount
of material produced. For the equation to accurately depict the amounts, it must be balanced.
That means the same atoms must appear on both sides of the equation and there must be the
same number of moles of each atom on both sides of the equation.
Rules for Balancing Chemical Reactions
1.) Write a skeletal equation: this will include the reactants and the products
2.) Balance the number of atoms: the number of atoms on the left side of the
arrow should be the same as the number of atoms on the right side of the arrow.
3.) Make the coefficients whole numbers:
even though we can accurately
balance a chemical equation using fractions, we will want all the numbers as
whole numbers. This means multiplying through the entire equation.
4.) Indicate the state of matter: are the reactants solids, liquids or gases? Wh
What about the products? Identify the state of matter!
Keep in mind that when you are balancing the equations that the coefficient operates on all the
atoms in the formula. Just like the parentheses when we write a chemical formula applies to all
atoms inside the parentheses, the coefficient operates on all atoms in the formula. You cannot
change the chemical formulas to balance the number of atoms. The formulas are set by the type
of compound it is. For example, in the reactions above, you cannot change Na 3PO4 to NaPO4 in
order to “balance” the number of sodium ions! You also cannot add new elements to chemical
reactions in order to balance the reaction. You do not balance the reaction by adding 2 Na +1
ions to the reactant side so that you have 3 sodium ions on each side. Use the reactants and
products given and balance by using coefficients. A balanced chemical equation will stay
balanced no matter if you multiply or divide each coefficient by the same value. You can
multiply the system by 2 or 10 or 1000, but the atoms/ions will still be balanced. However,
generally speaking, a balanced equation should contain the lowest whole number ratio of
coefficients!
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