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nomenclature and writing formulasSome textbook authors describe chemistry as the science of matter and change.
Many students find the changes more interesting than the matter but in order to
understand the changes we need to know something about the substances as they
are at the "beginning" and as they are at the "end".
Matter is something that occupies space and has mass. That's a basic definition from
some elementary science course in your past. Chemists go a little further and
subdivide matter into two categories:pure substances and mixtures.
Pure substances may be either elements (all atoms alike) or compounds(differentatoms combined in molecules of definite proportions and inseparable by physical
means).
Mixtures contain at least two pure substances and can--at least in theory--be
separated by physical means. Mixtures of the same substances will not necessarily
contain the same proportions of those substances.
Elements and compounds have constant properties which can be described as both
physical and chemical. Most of the substances we will work with and study during the
course are compounds.
Compounds vary a lot. Even compounds containing the same elements like water
(H2O) and hydrogen peroxide (H2O2) have very different chemical and physical
properties. So it can't be just the kinds of elements present in compounds that
determines their properties.
Compounds can be roughly divided into two very large categories based on the way
the elements are put together in them.
Ionic compounds are composed of positive and negative ions (atoms with extra or
fewer electrons than their neutral counterparts). In such compounds the total charge
of the positive and negative ions always adds up to zero. The attractions of theopposite charges holds the compounds together.
Molecular compounds are composed of neutral atoms which are held together by
sharing some of their electrons in common. No separate charges are involved.
How can you distinguish one kind of compound from the other?
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In general ionic compounds are composed of metals and non-metals. Molecular
compounds typically consist of only non-metals. In water solutions ionic compounds
are electrolytes. Molecular compounds generally are not.
For example, magnesium chloride, MgCl2, contains one metal (magnesium) and one
non-metal (chlorine). It is ionic. Nitrogen dioxide, NO2, contains only non-metals so itis molecular.
You will notice these compounds are named a little differently. The rules of chemical
nomenclature begin simply enough but eventually become pretty messy. Our goal
here is to keep things simple. The basic "bottom line" rules are:
in ionic compounds metals always come first and numerical prefixes
arenever used to indicate the number of a type of atom in the formula
in molecular compounds numerical prefixes must be used to indicate the
number of a type of atom in the formula
Examples
More examples
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Still more examples.....
Writing chemical formulas from names turns the process around. Molecular formulas
stand out because they will contain no metals and probably have at least one
numerical prefix. The words are simply translated into symbols in the same order in
which they appear in the name.
Ionic formulas have to be constructed more carefully since the total charge of the
ions must add up to zero. So subscripts must often be inserted into the formula to make
the math come out right.
Examples
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More examples
Still more examples.....
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Just to be sure the picture is not too simplistic, we should realize that many
compounds do not follow these simple rules. The compounds we have looked at are
mainly inorganic. That means they do not contain carbon and hydrogen.
Organic compounds have an entirely different system of nomenclature--much more
complicated--which we will not attempt to learn this year.
Organic compounds are typically molecular. These include compounds like ethanol
(C2H6O) and sucrose (C12H22O11).
There are also some compounds--like water--which have developed trivial names over
the years that are not systematic. No one calls water "dihydrogen monoxide". A small
list of these is included in your study guide. Learn them.
carbon monoxide CO
ammonia NH3
methane CH4
hydrogen sulfide H2S
Finally, the nomenclature of acids is more complicated than is worth learning in an
introductory course. So rather than struggle with a lot of rules you will never
use, learn these five common acids (also listed in your study guide):
hydrochloric acid HCl
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nitric acid HNO3
sulfuric acid H2SO4
phosphoric acid H3PO4
acetic acid CH3COOH
MOLE CONCEPT
One of the fundamental ideas of modern chemistry is that all matter ismade up of atoms. The atoms themselves are composed of smaller
particles. For our purposes these particles include the following three:
protons
neutrons
electrons
Most of the mass of an atom comes from the first two which together
coexist in the center of the atom called the nucleus. It is the number of
protons which is called the atomic number(Z, from the GermanZahlor"number") and differentiates one element from another. The atomic
numbers are the sequential integers on the periodic table which currently
number the elements from 1 to 116.
Neutral atoms (as opposed to ions) have a number of electrons equal to the
number of protons in the nucleus. Electrons are found in the space
surrounding the nucleus rather than inside it.
And while all neutral atoms of say, oxygen, always contain 8 protons and 8
electrons they do not all necessarily have the same number of neutrons.
Atoms of the same element with different numbers of neutrons are knownas isotopes.
Some oxygen atoms may have 8 neutrons while others may have 7 or 9.
Chemically these atoms are essentially the same and often come mixed in
nature. Thus the average atomic masses (which are the other numbers on
the periodic table) are weighted averages of the mass numbers (A, from the
GermanAtomgewichte or "atomic weight") of the different isotopes that
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occur naturally. The mass numbers themselves are simply the sum of the
masses of the protons and neutrons on the atomic mass scale---a scale on
which each particle is counted as 1.
For example, the oxygen isotope with a mass number of 16 contains eight
protons and eight neutrons (as well as eight electrons). Another oxygenisotope with a mass number of 18 contains.....how many neutrons???
Since all oxygen atoms have 8 protons (the atomic number orZ), 18 - 8 = 10
neutrons. The general relationship would be:
A =Z+ #n
Complete isotopic symbols contain this information.
The simple atomic mass scale (amu) gives only relative masses. Individual
atoms are far too small to be massed with instruments we typically use in
the lab.
Historically this was quite a problem. How to mass something you can't see?
If you can't measure mass reliably then the roots of modern chemistry
wither away because the most fundamental laws of chemical combination
are based on combining masses of substances and mass conservation
overall during chemical changes.
A number of approaches have been used:
using the lightest element (H) as the mass standard set at 1
using oxygen (with which most elements combine) set at 16
using the isotope carbon-12
The third method is the current mass standard. All atomic masses
are relative to the mass of an atom of carbon-12. For example, the average
atomic mass of hydrogen is very close to 1. That means an atom of
hydrogen is 1/12 the mass of an atom of carbon-12. Similarly, the average
atomic mass of magnesium is about 24. So a magnesium atom is about
twice as heavy as an atom of carbon-12.
Because the atomic mass scale is relative, the actual units (amu, g, tonnes,etc.) don't really matter. Two atoms of magnesium will still be about twice
as heavy as two atoms of carbon-12. In that simple idea lies the modern
solution to the dilemma of massing what you can't see.
Consider the following comparisons:
1 atom H : 1 atom C-12
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10 atoms H : 10 atoms C-12
100 atoms H : 100 atoms C-12
(you get the idea...)
What is the mass ratio in each case? One to twelve
Tell us something we don't know! Well, the problem with 100 or even 1000
atoms is that you still can't mass them. They are just too small. So chemists
have approached this problem from the opposite end: pick reasonable and
convenient masses in the ratio of 1:12--how about grams?
1 gram H : 12 grams C-12
What do these masses have in common????
They each contain the same number of atoms. How many atoms? A lot.
Modern determinations of the number of atoms in 12 g of carbon-12 put it
somewhere around 6.02 x 1023 atoms.
This is a BIG number.
This number is known as Avogadro's number (NA). The quantity is called
themole (L. "heap or pile"). The relative atomic mass of an element in grams
is commonly called the molar mass of the element. For carbon that would
be about 12.0 g/mol.
Moral #1: the average atomic mass of an element, when expressed in
grams, is one mole of that kind of atom.
This concept enables chemists to essentially "count" atoms by massing
them. For example, if 12.0 g of carbon contains 6.02 x 1023 atoms then 6.00
g of carbon must contain half that amount or 3.01 x 1023 atoms.
Even simpler, we could say that if 12.0 g of carbon is 1 mole of atoms, then
6.00 g of carbon must be 0.500 mole of atoms.
Moral #2: the ability to change back and forth between grams and moles is
survival skill No. 1 in chemistry
Converting between grams and moles can be thought of in terms of simple
proportions or set up using unit analysis as the following examples will
show.
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It is also possible to determine the number of atoms (or molecules) in a
sample in a similar fashion, although this is very seldom done.
Finally, the molar masses of compounds can be determined by adding the
individual atomic masses of the constituent elements.
Examples
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More examples
Still more examples..
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% composition and empirical formulas
With what we know today about ion charges, nomenclature rules and so on,
writing chemical formulas is fairly routine. But where did formulas come
frombefore all of that was figured out?
Early chemists spent a lot of time developing techniques to analyze
substances, eventually arriving at data such as the number of grams of
hydrogen and oxygen in a given mass of water. In percent form thisinformation is called "percent composition by mass" or simply percent
composition.
The calculation of % composition can be based on what we know today is
the correct formula for water, H2O. The molar mass of water is 18.0 g/mol.
Of that 18.0 g, 2.0 g are hydrogen and 16.0 g are oxygen. So the mass
percents of each element could be expressed as:
EXAMPLE
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In some respects this method puts the cart before the horse. But the
examples shown can be used in some situations to help answer parts of
more complicated questions.
Generally experimental mass data is used to determine a chemical formula
in a laboratory. There are also more advanced instrumental methods
available today but they won't help you understand how grams, moles and
chemical formulas work together!
Consider the data obtained by analyzing a compound. A chemist determines
that the substance contains potassium, chlorine and oxygen. Here is some
sample data:
elementmass,
g
potassi
um0.479
chlorine 0.434
oxygen 0.588
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What is the chemical formula of this compound? It might be tempting to say
K0.479Cl0.434O0.588 but of course that is not correct.
The numbers in chemical formulas are always integers since they represent
actual whole atoms which combine. The fact that they represent
the numbers of atoms and not their masses should suggest a direction to
follow.
n a formula like H2O, two atoms of hydrogen are combined with one atom of
oxygen. We could also say that two moles of hydrogen atoms are combined
with one mole of oxygen atoms. So if we knew the moles of K, Cl and O we
would be that much closer to knowing the formula.
elementmole
s
potassi
um
0.01
23
chlorine 0.0122
oxygen0.03
68
This is a little disappointing as we are no closer to integers than we were
before. A closer look at those decimals will show that there is a nearly
integer ratio hidden in the data. One way to get it out is to divide all of the
values by the smallest value. That makes the smallest number 1 and
hopefully all the other values larger integers (or also 1).
So the formula for this compound is KClO3. Such a formula is often called
theempirical formula. It gives the smallest integer ratio which represents
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the proportions in which the atoms combine in the compound. It may also
represent the actual or molecular formula.
To illustrate the difference, consider the empirical formula for hydrogen
peroxide: HO. The known molar mass of hydrogen peroxide is 34.0 g/mol.
But the empirical formula mass is only 17.0 g/mol. That means the actualmolecular formula for the compound must be H2O2 or twice the empirical
formula.
Molecular formulas are either the same as the empirical formula or some
integer multiple. To know whether an empirical formula is also the
molecular formula you need to know the molar mass of the compound.
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