Reading
Notes for Chapter 3
These are Dr. Bodwin's reading notes for Chapter 3 of "Chemistry 2e" from OpenStax. I am using a local .pdf
copy that was downloaded in May 2020.
Chapter
Summary:
This chapter builds upon Chapter 2 by developing the tools that we will use
to make macroscopic observations of the microscopic world of atoms, ions,
and molecules. Keep on reviewing chemical formulas and know your polyatomic
ions - if you don't have those right, formula masses and moles are going to
be much more challenging.
Formula
Mass:
In many ways, formula mass is just another way of expressing a chemical
formula, and because of that it is essential to have a correctly balanced
chemical formula if you want to have a correct formula mass.
There are a few terms that are commonly used to describe this same idea:
"formula mass", "formula weight", "molecular weight", "molecular mass",
"molar mass"... there are some subtle differences in exactly what these
terms mean, but they pretty much refer to the same idea.
Don't make this harder than it is. If we were doing this with grocery
store items instead of atoms and molecules, you'd be able to figure it out
without a second thought. Each type of atom in a formula has a mass listed
on the Periodic Table. Add up the masses of all the pieces in the chemical
formula, and you have the formula mass. Carbon dioxide has 1 carbon and 2
oxygens. A carbon has an atomic mass of 12.011amu and each oxygen has an
atomic mass of 15.999amu. Carbon dioxide has a formula mass of
(12.011+15.999+15.999) = 44.009amu.
Formula Masses
The Mole:
Another example of "don't make it harder than it is". A mole is a grouping
unit. When we're counting atoms in a macroscopic sample, the number of atoms
is HUGE, so to make it more manageable to discuss that amount, we use a
great big number that we call a "mole". One of my favorite cookies come in
boxes of 36 cookies. That box is a grouping unit. When I go to the store, I
don't think "hey, I should pick up 36 cookies", I think "hey, I should pick
up a box of cookies". Moles work the same way. I don't say "Here's
602214179000000000000000 sodium atoms", I say "Here's 1 mole of sodium
atoms"
Back to formula mass... I used "atomic mass units" or "amu" in the example
above. But that's not a really useful macroscopic
mass unit. Here comes the mole... when we're looking at an atom, amu is a
convenient mass unit; when we're macroscopically looking at a mole of some
substance, grams makes more sense. Because some scientists were smart about
the way they defined the mole, it turns out that the "mass" listed on the
Periodic Table has units of amu when you're looking at atoms, but has units
of grams when you're looking at moles
of atoms. An average sodium atom has a mass of 22.990amu... a mole
of sodium atoms has a mass of 22.990g!
Because we will almost always be making macroscopic observations of the
microscopic world, we're going to be using moles a LOT. Again, a mole is
just a grouping unit, so we can talk about moles of anything... atoms, ions,
molecules...
And just to try and get a better idea of just how big a mole is...think of
something small, like a blueberry. If we say a blueberry is 0.5cm in
diameter and we pile them in a simple cubic packing arrangement (http://www.ntci.on.ca/chem/sch4u/crystalpacking.pdf),
we would have a cube of blueberries a bit over 260 miles wide. That's a lot
of blueberries.
Moles from Grams
Empirical
Formulas and Molecular Formulas:
Percent composition and finding empirical formulas are the same process in
opposite directions. Start with percent composition:
Sodium thiocyanate (NaSCN) has a formula mass of
(22.990+32.066+12.011+14.007) = 81.074g/mol. Its percent composition is:
%Na = (22.990g/mol / 81.074g/mol)*100 =
28.357% Na
%S = (32.066g/mol / 81.074g/mol)*100 = 39.552% S
%C = (12.011g/mol / 81.074g/mol)*100 = 14.815% C
%N = (14.007g/mol / 81.074g/mol)*100 = 17.277% N
{NOTE: Those don't add up to exactly 100% because of rounding.}
OK, let's go the other way... start with percents:
A substance is found to be 36.112% calcium and 63.888% chlorine. What is its
formula?
The problems almost always start with
"assume you have 100g of sample" because it makes the math easier, but I
think students sometimes get confused because the math was made easier!
Let's assume that we have some amount of substance... let's try 9.37g.
From the percent composition data, we know that if we have9.37g of
substance, then (9.37*0.36112) = 3.38g of the sample is calcium and
(9.37*0.63888) = 5.99g of the sample is chlorine.That means there are
(3.38g Ca / 40.078g/mol) = 0.0843mols of Ca in the sample and (5.99g Cl /
35.453g/mol) = 0.169mols of Cl in the sample. We can take a ratio of those
moles (with the smaller number on the bottom...) and simplify that ratio
to see that this samplecontains 2 moles of chlorine for every mole of
calcium. The empirical formula is CaCl 2. {Which is what we
would predict based upon expected charges from the Periodic Table.}
An "empirical formula" represents the smallest whole-number ratio of
elements. For larger molecules, there are often multiple empirical formula
units, so the "molecular formula" might be a multiple of the empirical
formula. Without additional information, we cannot determine if the
molecular formula contains more than one empirical formula unit.
Solutions
and Concentration Units:
Solutions are homogeneous mixtures of 2 or more substances.
The major component of a solution is called the "solvent".
The minor component(s) is called the "solute(s)".
We use "concentration units" to describe the amount of solute relative to
the amount of solvent or solution.
There are a whole bunch of fraction-based concentration units (fraction,
mole fraction, percent, ppm, ppb, etc) that are used in practical
settings. In general, don't make them harder than they are just because
you're studying chemistry and you should be fine. If a solution contains
10.0grams of glucose and 990.0grams of water, then it's a 1% glucose
solution. If a solution contains 2.58mg of chloride ions in 3.0L of water,
the solution is 2.58ppm. {There are a couple assumptions embedded in that
calculation... can you find them?} Concentration (ppm)
The most important concentration unit for us to use in this class is
Molarity, moles of solute per liter of solution, designated with an "M". Moles from
Concentration and Volume
That capital "M" is important! If you use a lower case "m", it means
something completely different.
The other concentration units mentioned in this chapter show up in a number
of different places and contexts, but we won't worry about them at this
point. We will revisit solutions at the beginning of Gen Chem II.
When diluting a solution, you can just about always use C1V1=C2V2
to determine the new concentration. This simple formula is your friend when
you use it for dilutions! Dilutions
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