Nomenclature and Writing Formulas

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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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Nomenclature and Writing Formulas

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