Atoms rarely exist as separate entities – they are usually combined with other atoms making a
compound or a molecule. Some elements do occur as free atoms in nature – and you should
know them – specifically the noble gases (Group VIIIA). Other atoms will combine with
themselves, such as O2 (probably heard of that!), N2, H2, Cl2, F2, Br2, I2, as well as other
interesting ones like S8, and P4.
By in large, atoms like to hang out together and form
molecules.
A molecule is a group of two or more atoms held together in a definite spatial arrangement by a
force called a bond. It is electrically neutral, and the atoms exist in whole number ratios in the
compound (thus we do not have half a hydrogen present in a molecule!!!) In order to form the
bond the electrons of the elements interact with one another, generally in one of two ways:
1. sharing the electrons forming a covalent bond (sharing can be equal or unequal)
2. transferring the electrons from one atom to another to form an ionic bond (this is an
electrostatic attraction between oppositely charged particles)
Chemical Formulas:
1. empirical formulas: the simplest formula, the most numerically reduced formula one can
write for a compound. It lists all the elements present and indicates the smallest integral
(whole number) ratio in which the atoms are combined. Thus the empirical formula does
not tell you the actual number of atoms present. The empirical formula is written with
the element capitalized and the number of each atom as the subscript. By convention, the
number 1 is not ever written as a subscript. Thus C1H2 is written as CH2.
Example: a compound with an empirical formula of CH just indicates that only
carbon and hydrogen are present in the molecule and their ratio is 1:1. Many
compounds fit this empirical formula, for example benzene (C6H6) and ethene
(C2H2) are obviously NOT the same molecule but they share the same lowest
ration empirical formula.
2. molecular formula: shows how many atoms of each element are actually present in the
molecule. It provides a more accurate picture of the molecule itself. One can see the
difference in the molecules when looking at the molecular formula compared to the
empirical formula. Thus, C6H6 and C2H2 are molecular formulas which correspond to
benzene and ethene respectively. The molecular formula is written with the element
capitalized and the number of each atom as the subscript.
3. structural formula: shows the number of atoms and how they are arranged in space.
Below are the structural formulas of cyclohexane (C6H12) and ethanol (C2H6O). As you
can see by comparing the molecular formula to the structural formula, the spatial
arrangement IS important and is not really indicated in the molecular formula.
1
H H
H C
C
H C
H H
H
C H
H
C
H
C H
H H
H C C O
H
H H
H
cyclohexane
ethanol
Ionic compounds:
Remember that the noble gases are also called the inert gases, inert meaning no reactions. So
they will not adopt a charge and for all practical purposes they are a reference point for us on
the periodic table – the ideal. Consider them charge zero. Also, consider the number of
electrons that they contain to be special – an “ideal” number to be achieved.
Bonding between a metal and a nonmetal or between two charged species (e.g. polyatomic ions)
will be considered ionic. We can go back to the periodic table, remember the one that I said
has most of the answers on it – to determine the charges that particular atoms will have. We are
looking at the metals
1
2
3
4
Specifically we are looking at the metals in hot pink in groups 1-4 (notice that hydrogen is not
hot pink – because it is not a metal. You will oftentimes find hydrogen elemental boxes on the
2
periodic table floating above the table, or also over top of Fluorine. This is because while
hydrogen is a non-metal, you can find hydrogen as a +1 ion just like the rest of that column, or
you can find it as a -1 ion just like the fluorine column (info coming next!)). And just like the
metal column numbers indicates –that is the CHARGE that the element will adopt when
forming an ion. The charge comes from the loss of an electron (or electrons). When a metal
loses an electron (or 2 or 3) it then has more protons (positively charged species) than electrons
(negatively charged species). Group 1A ALL turn into +1 ions – they each have 1 less electron
than proton. Group IIA ALL turn into +2 ions – they each have 2 less electrons than protons.
Group IIIA turn into +3 ions – they each have 3 less electrons than protons. And so on. Group
IVA is a little trickier, but right now, let us just say they are +4 ions and we’ll go from there . . . 
Please note – ions are only formed by changing the NUMBER OF ELECTRONS. If you try and
achieve a charge by changing the number of protons, you just committed alchemy! Remember
the number of protons gives us the identity of the element (or ion). If you change the protons
instead, you change the identity of the chemical – you are NOT creating the ion you think you
are. Sadly, alchemy is not possible – so ONLY change the number of electrons, NEVER protons.
Let’s examine these charges in terms of the metal and the noble gases.
Sodium metal – has 11 electrons. The nearest numeric noble gas (remember, the noble gases
have “ideal” numbers of electrons, has 10 electrons. How can sodium get to 10 electrons also?
By losing 1 electron. And if it loses 1 electron it now has fewer negative particles than positive,
and therefore has a charge of +1. Examine the rest of column 1 . . . what pattern do you notice?
K – 19 electrons
Rb – 37 electrons
Cs – 55 electrons
Ar – 18 electrons
Kr – 36 electrons
Xe – 54 electrons
The difference in the number of electrons in column 1 and the ideal number of electrons in the
noble gas column is 1. If those elements in column 1 simply lose 1 electron, then they have the
ideal number. And notice that the charge over top of column 1 is +1!
Consider the same process for the metals in column 2.
So what about the negative species? Looking back at our periodic table, if the metals are the
positive species then the nonmetals that they bond with (remember ionic bonds are a metal
pairing with a nonmetal) MUST be the negative species.
Remember, the noble gases are the “ideal” which means all the other elements strive to “look”
like them. By look like them I mean have the same number of??? Electrons. Remember that
isotopes have different numbers of neutrons but the same number of electrons, and since
chemical properties depend on the number of electrons those are the most important subatomic
particles. So if Na loses an electron, it then has the same number of electrons that Ne has – Na
lost 1 of its 11 electrons, it still has 11 protons but now it is isoelectronic with Ne – which has 10
electrons (and 10 protons).
3
So, working backwards away from the noble gases the periodic table looks like this:
0
4
3
2
1
So how can the non-metals look like (become isoelectronic with) the noble gases? Well, think
about it. If F were to start losing electrons it would have to go from having 9 electrons alllllll the
way down to having 2 electrons!! Or Cl would have to go from having 17 electrons allllllll the
way down to 10 electrons. It would be much easier and take a WHOLE lot less energy to just
GAIN an electron! Group VIIA ALL turn into -1 ions – they each have 1 more electron than
protons. Group VIA ALL turn into -2 ions – they each have 2 more electrons than protons.
Group VA turn into -3 ions – they each have 3 more electrons than protons. And so on. Group
IVA lies in the middle of becoming positive or negative . . . sooo they can do either.
We will define ionic bonding when we see a metal in the compound. While there are some
exceptions to this – at this point in time, when you see a metal – we will presume the bonding
to be ionic.
The simplest type of ionic compound is a binary ionic compound, which is formed from TWO
elements (bi = two). Generally speaking we are looking at a metal bonding with a nonmetal. As
in above, the metal will lose its electron(s) and become positive – called a cation. The nonmetal
will gain the electron(s) given up from the metal – become negative and is called an anion.
Metal+
Cation+
NonmetalAnion-
When forming an ionic compound you MUST end up with a neutral species!! This means that
the charge from the cation must equal the charge on the anion. The strength of the attraction of
the ions can be expressed by Coulomb’s Law: the energy of attraction or repulsion between two
4
particles is directly proportional to the product of the charges and inversely proportional to the
distance between them.
In English: ions with greater charges (e.g. +2, +3, +4 etc. . . ) attract or repel each other more
strongly than ions with lower charges (e.g. +1 compared to +2, +2 compared to +3). Likewise
smaller ions attract or repel each other more so than larger ions because the charges can get
closer together.
No molecules exist in an ionic compound. Although the overall charge is neutral, a sample of
an ionic compound contains no molecules (neutral species), only ions (charged species)!!
Covalent Compounds:
Unlike ionic bonds, which can be thought of as 100% unequal sharing of electrons (more like a
give and take relationship), covalent bonds are involved in sharing the electrons. No one gives
up and no one takes, each atom in the bond shares, not always equally, but the electrons are
shared nonetheless. Some examples of “perfect” (or as perfect as one can get) sharing between
atoms are any species that exists as a diatomic (X2), such as H2, O2, N2, and the like. Atoms of
different elements can also share their electrons to form molecules: H 2O, NH3, and CH4.
Generally speaking, covalent bonds occur between NONMETALS. (remember hydrogen is not
a metal). In this course, most of our covalent bonding will occur between non-metals. When
there are only non-metals in a compound we will define this as covalent bonding.
Polyatomic Ions:
A polyatomic ion consists of two or more atoms that are covalently bonded together but have a
net positive or negative charge. For example, the carbonate ion CO 3-2 consists of covalently
bonded carbon to oxygen but the molecule has a net negative 2 charge.
To this point all molecules have had an overall charge of zero, however, this is not the case with
polyatomic ions.
o poly = many
o atomic = atoms
o ion = charged particle
o So…polyatomic ion means a molecule of multiple atoms that have a charge.
Polyatomic ions are stable. The ions are found in solutions or solids where the overall charge of
the solid is zero or the solution is zero. What this means is that you can not have a test tube full
of Na+ and no negative charges like a chloride (Cl-1) to cancel out the overall positive charge. IT
CAN NOT BE DONE, a law of physics, and you can’t break the laws of physics.
5
o the list of polyatomic ions you must commit to memory. (note only those in
bold/highlighted need to be memorized, the rest be familiar with)
Formula
NO3NO2CrO42Cr2O72CNSCNMnO4OHO22NH2SO42SO32PO33PO43HPO42H2PO4C2H3O2CH3COO-
Name
nitrate
nitrite
chromate
dichromate
cyanide
thiocyanate
permanganate
hydroxide
peroxide
amide
sulfate
sulfite
phosphite
phosphate
hydrogen phosphate
dihydrogen phosphate
acetate
acetate
Formula
ClO4ClO3ClO2ClOIO4IO3IO2IOBrO4BrO3BrO2BrOCO32HCO3HSO4HSO3HSNH4+
Name
perchlorate
chlorate
chlorite
hypochlorite
periodate
iodate
iodite
hypoiodite
perbromate
bromate
bromite
hypobromite
carbonate
hydrogen carbonate or bicarbonate
hydrogen sulfate
hydrogen sulfite
hydrogen sulfide
ammonium
Monoatomic Ions to commit to memory (from the Periodic Table)
6
Names and Formulas for Ionic Compounds:
Compounds from Monoatomic ions:
1.) the name of the cation is exactly the same as the metal! Na is the same name as
Na+1
2.) the name of the anion keeps the root portion (first portion) of the name and then
adds – ide. Cl = chlorine drop the ine and add –ide so that Cl-1 = chloride
3.) the charges must balance – meaning they must cancel out
if the cation has a +2 charge and it is paired with an anion that has a -1
charge you need (2) -1 anions to balance out (1) positive cation.
Ca+2 paired with Cl-1: if we wrote CaCl then we would have (1) net
positive charge left over, so we need (2) Cl-1 species. We write the
formulas with the number of atoms needed as a subscript. Thus Ca+2
paired with Cl-1 becomes CaCl2.
This has often been termed “cross the charges”
Ca+2
Cl-1
Ca1Cl2
write the 1
→ except by convention we do not
This method of crossing the charges can also be used to determine
what the charge of the ion would be if you separated the compound
into its ions.
AlI3
Answer: if we “uncrossed” the charges shows us that the Al would
have a +3 charge and the I would have a -1 charge – and see – we did
not even need the periodic table (but that would confirm our
uncrossing method since Al is a metal and in Group IIIA and I is a
nonmetal and in Group VIIA!! – check for yourself!)
4.) Some metals can form more than one positive ion, meaning they can give up
different numbers of electrons and form different charges. Typically, the
elements that do this are known as the inner transition metals – found on the
periodic table labeled with a “B” instead of the “A” that we have been examining
thus far. For reasons we’ll talk about later . . . these elements can form different
ions. For now, examine the periodic table above and memorize them 
5.) There is an antiquated form of naming these ions, with an ic or ous, therefore I
am requesting that you not use this method . . . just be aware that it IS out there
so just be familiar with it! We will use the roman numeral approach to naming
these ions. In this case, the roman numeral actually indicates the charge for you!
Roman numeral II indicates a +2 charge while Roman numeral III indicates a +3
charge. Using the roman numeral method of naming Fe+2 would be iron(II) and
Fe+3 would be named iron(III).
7
Compounds from Polyatomic ions:
1.) the polyatomic ion is a group that STAYS TOGETHER!!! It is a unit, an entity unto itself
– do NOT split it up into its individual atoms!! It is no longer a polyatomic atom if you
do that!!
2.) Name the cation first (same as above) and then name the polyatomic atom (from the list
given to you!)
3.) Again you MUST make sure that the overall charge on the compound is neutral once you
combine the cation (anion in the case of NH4+1) with the polyatomic ion. Since these ions
are UNITS/GROUPS/PARTNERS you must indicate that you want more of the total ion
grouping, so we use parentheses with the number needed as a subscript.
Example: Pairing Al+3 with NO3-1
Al has a +3 charge and the charge on the total NO3 species is a -1. Therefore
we need (3) units of NO3-1 in order to have a neutral compound. Thus the
molecular formula for aluminum nitrate will be Al(NO3)3. Again, we are
using the “cross the charges” method! Only now we need 3 sets of NO3 If
you just wrote NO33 it obviously looks like a nitrogen atom bonded to 33
oxygens!! BAD!!
4.) On the polyatomic ion list there are several ions that have the same root name and only
differed in the number of oxygens present (e.g. BrO4-1, BrO3-1, BrO2-1, BrO-1). The charge
is the same for EACH ion, the only difference is the number of oxygens present. Even
their names are veeeerrrry similar. These are the oxoanions. And there are families of
oxoanions that are apparent on the list above. It is often easies to learn ONE of the
oxoanions and remember that the names change in a systematic way as the number of
oxygens increases or decreases.
When there are only 2 oxoanions in the family (SO4-2 and SO3-2), the ion
with the MOST oxygen atoms gets the –ate ending (sulfate) while the ion
with the LEAST oxygen atoms gets the –ite ending (sulfite).
8
When there are 4 oxoanions in the family (BrO4-1, BrO3-1, BrO2-1, BrO-1), the
ion with the MOST oxygen atoms gets a prefix – per – and a suffix – ate :
BrO4-1 would be perbromate. The ion with the next most oxygen atoms
drops the per and becomes bromate. The ion with the next most oxygen
atoms becomes bromite (parallels when there are only 2 oxoanions), and
the ion with the least oxygen atoms gets the prefix – hypo – and the suffix ite
and is named hypobromite.
5.) Ionic compounds can do one more thing, adding to their name. They can pick up some
water – termed a hydrate. For ionic hydrates we follow allll of the rules given up to this
point and then indicate the number of water molecule using prefixes given below
followed by hydrate:
Number
1
2
3
4
5
6
7
8
9
10
Prefix
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
Thus: Ba(OH)2∙8H2O would be: barium hydroxide octahydrate
Ba is the metal (cation) which is named barium it has a +2 charge
(notice its in group 2A!!) while OH has a -1 charge, thus we need 2
OH-1 to cancel out the +2 charge of Ba
OH from the polyatomic ion list is hydroxide
Octa from the prefix list stands for 8
Hydrate indicates the presence of water
And: copper(II) nitrate trihydrate would be: Cu(NO3)2∙3H2O
Nitrate is a -1 charge and copper(II) is a +2 charge therefore when the
charges are “crossed” we have 1 Cu and 2 (NO3) groups
Tri indicates that there are 3 of something
Hydrate indicates water
9
Naming Acids: Specifically looking at H containing acids (defined as Arrhenius acids).
Typically they are in water, thus in solution. When naming them, we consider the hydrogen to
be a hydro group as the cation bonded to an anion. Specifically you should be familiar with the
binary (2 atoms) acids involving H and the elements in Group VIIA.
Binary Acids: HF, HCl, HBr, HI: use hydro in the name
Named: prefix + nonmetal (drop –ide adding) + ic + acid
HF = hydro
fluoride
ic acid = hydrofluoric acid
HCl = hydro chloride
ic acid = hydrochloric acid
HBr = hydrobromic acid
HI = hydroiodic acid
Oxoacid: from the polyatomic ion sheet, the anions that contain oxygen can be paired
with H thus making an acid (e.g. H2SO4, H2SO3). These names do NOT include hydro!
There is a “fun” saying to remember how to name oxoacids:
-ates become ic’s and -ites become ous’es
So H2SO4 from SO4-2 (sulfate) becomes sulfuric acid (notice no hydro!!!)
H2SO3 from SO3-2 (sulfite) becomes sulfurous acid (notice no hydro!!!)
In situations where per-ate and hypo-ite are used we keep the prefixes but change the
suffixes as above
HBrO4 from BrO4-1 (perbromate) becomes perbromic acid
Binary covalent compounds: formed when 2 elements – specifically nonmetals combine
together to share their electrons. This does not mean 1:1 ratio of the elements, the molecular
formula could be X10Y8 which indicates 18 total atoms – but the bonding is occurring between
two elements – X and Y.
Rules for naming binary covalent compounds:
1. The first element is named as the element. If there is more than one of that
element we must indicate how many using the same prefix system that we use for
the naming of hydrates. If there is only one of the first element, by convention we
do not use the prefix mono.
2. The second element is named as its root with the suffix –ide.
3. The second element will always have a numerical prefix.
Examples: PCl5, CO, CO2 H2O
PCl5 = phosphorus pentachloride
CO = carbon monoxide (notice it is not monocarbon monoxide!!)
CO2 = carbon dioxide
H2O = dihydrogen monoxide (which is the “real” name for water!)
10
Rules 4 and 5 are less commonly used, but you should be aware of them.
4. the element with the lower group number on the periodic table is the first word in
the name. The element with the higher group number is the second word in the
name. [Exception!!!!: when a halogen (group VIIA) is covalently bonded to
oxygen, the halogen is named first and the oxygen is named second.
5. If both elements are in the same group, the one with the higher period number is
named first (remember period is the row)
Number
1
2
3
4
5
6
7
8
9
10
Prefix
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
Calculating molecular weights of compounds and formula weights of ionic compounds:
Using the number of atoms present in the compound and the molecular weight of the
individual atoms or ions, we can calculate an overall molecular weight for the entire compound.
Molecular weight of the compound = sum of the atomic mass of the elements present in the
compound.
Example: Molecular weight of C2H6
2C atoms = 2 * 12.01g/mole = 24.02g/mole
6H atoms = 6* 1.008g/mole = 6.048g/mole
Then the individual atomic masses are summed 24.02+6.048 = 30.07g/mole
Example: Molecular weight of Al(NO3)3
1 Al = 1* 26.98g/mole = 26.98g/mole
3N = 3* 14.01g/mole = 42.03g/mole
9O (remember 3 sets of NO3 means 9 oxygens total!!) =
9*15.99g/mole = 143.91g/mole
Then take the sum: 26.98+42.03+143.91 = 212.92g/mole
Note – significant figure rules apply when calculating molar masses – remember we have
multiplication rules AND addition rules to pay attention to (total number of sig figs and
decimal places)
11
“the monkey wrench”
Another relationship exists that you should be aware of. It is DIFFERENT than the relationship
between grams and moles and the two should NOT be confused. Avogadro’s Number (NA) is
used to calculate the number of entities – e.g. atoms, molecules, formula units, ions, etc in 1
mole of substance. NOTICE it is not used for grams!! Grams to moles and moles to grams
only uses the molecular weight!!
Avogadro’s Number is used to convert:
Moles to atoms
Moles to ions
Moles to molecules
Moles to formula units
Avogadro’s Number = 6.022 x 1023
entities X
moles X
Given: 4 moles of carbon want to know ? atoms
4 moles Carbon x
6.022x10 23 atoms of carbon
= 2.409 x 1024 atoms of carbon
1 mole of carbon
Given: 16 grams of carbon want to know ? atoms
Step 1: convert grams to moles (using MW)
Step 2: convert moles to atoms (using NA)
16.0 grams C x
1 mole C
= 1.33 moles of Carbon
12.01grams C
1.33 moles C x
6.022x10 23 atoms of carbon
= 8.02 x 1023 atoms of carbon
1 mole of carbon
Note – as you become more proficient in your calculations, you should attempt to do this
mathematical operation in one step
16.0 grams C x
6.022x1023 atoms of carbon
1 mole C
x
= 8.02 x 1023 atoms of carbon
1 mole of carbon
12.01grams C
12
Given: ethylene which consists of C2H4 molecules. How many atoms of C are
in 2.50 moles of C2H4 and how many atoms of H are in 2.50 moles of C2H4?
Note: 1 mole of C2H4 contains 2 moles of C and 4 moles of H
2.50 moles C2H4 x
2 mole C
6.022x10 23 atoms of carbon
x
= 3.01 x 1024 C atoms
1 mole C 2 H 4
1 mole of carbon
2.50 moles C2H4 x
4 moles H
6.022x10 23 atoms of H
x
= 6.02 x 1024 H atoms
1 mole C 2 H 4
1 mole of Hydrogen
Given: How many Ca ions are there in 4.56 grams of calcium chloride? How
many chloride ions are there in 4.56 grams of calcium chloride?
Step 1: write molecular formula for calcium chloride
Step 2: calculate FW (formula weight) for calcium chloride
Step 3: convert grams to moles using answer from Step 2
Step 4: convert moles to atoms using NA
Calcium chloride = Ca+2 + Cl-1 = CaCl2
CaCl2 = Ca: 1 x 40.078 = 40.078
Cl: 2 x 35.453 = 70.906
FW = 40.078 + 70.906 = 110.984 grams/mole
4.56 grams CaCl2 x
1 mole CaCl 2
1 mole Ca
6.022x10 23 Ca 2 ions
x
x
= 2.47 x 1022 Ca+2 ions
1 mole of Ca
110.984 grams CaCl 2 1 mole CaCl 2
4.56 grams CaCl2 x
1 mole CaCl 2
2 mole Cl
6.022x10 23 Cl -1 ions
x
x
= 4.95 x 1022 Cl-1 ions
1 mole of Cl
110.984 grams CaCl 2 1 mole CaCl 2
Given: 4.56 g of calcium chloride ? formula units do you have
6.022x10 23 formula units of CaCl 2
1 mole CaCl 2
4.56 grams CaCl2 x
x
= 2.47 x1022 formula
1 mole of CaCl 2
110.984 grams CaCl 2
units of CaCl2
Each element in a compound makes up its own percentage of that compound. A certain % of
the mass of CaCl2 is attributed to the calcium and a certain % of the mass of CaCl 2 is attributed
to the chlorine (chloride). As with any % - the sum of the parts should add up to 100!
13
The mass % of a particular element in the compound can be calculated by the following
equation:
moles X in formula x molar mass of X (grams
mole)
Mass % of element X =
mass (grams) of 1 mole of compound
Given: A compound with the molecular formula C3H4O is known as acrolein. ? mass
percent of each element in acrolein. ? grams of oxygen in 144.5 mg of acrolein
There are 3 moles of C for every 4 moles of H for every 1 mole of O in acrolein.
Knowing the moles of each element, calculate grams
3 moles C x
12.000 g C
= 36.000 grams C in 1 mole
1 moles C
4 moles H x
1.0078 g H
= 4.0312 grams H in 1 mole
1 moles H
1 mole O x
15.999 g O
= 15.999 grams O in 1 mole
1 moles O
Total mass of acrolein = 36.000+4.0312+15.999 = 56.030 grams/1mole
Mass % of carbon =
36.000 g C
x100 = 64.251% C
56.030g acrolein
Mass % of hydrogen =
Mass % of oxygen =
4.0312 g H
x 100 = 7.195% H
56.030g acrolein
15.999 g O
x 100 = 28.554% O
56.030g acrolein
Math check: 64.251%+7.195%+28.554% = 100.000%
Sample of acrolein 144.5 mg ? grams of O:
1 grams
= 0.1445 grams
1000 mg
Acrolein is composed of 28.554% O, so 28.554% of 0.1445 grams is oxygen
0.1445 grams x 0.28554 = 0.04126 or 4.126 x 10-2 grams of oxygen in 144.5 mg of
acrolein
Convert 144.5mg to grams: 144.5 mg x
14
Given: A 28.64 mg sample of vitamin A, a compound consisting of C, H, and O, is
burned to form 88.02 mg of carbon dioxide and 27.03 mg of water. Calculate the % of
each element in vitamin A
To find the % of each element in vitamin A we need to find the mass of each element in
the original sample of vitamin A. After combustion, all the carbon is CO 2 and all the
hydrogen is water.
Thus:
So:
the mass of carbon in vitamin A = mass of carbon in CO2
The mass of hydrogen in vitamin A = mass of hydrogen in H2O
The mass of oxygen in vitamin A is tied up in BOTH water and carbon
dioxide
the mass of carbon in 28.64 mg vit A = mass of carbon in 88.02 mg of CO2
the mass of hydrogen in 28.64 mg of vit A = mass of H in 27.03 mg of H2O
the mass of oxygen in 28.64 mg of vit A = 28.64 mg – mg C – mg H
calculate oxygen last!!
1.)
2.)
3.)
4.)
convert all masses to grams: calculate MW’s
calculate mass of C in CO2
calculate mass of H in H2O
calculate mass of O from difference
1 grams
1 grams
= 0.08802 grams & 27.03 mg H2O x
= 0.02703 grams
1000 mg
1000 mg
1 grams
28.64 mg vitamin A x
= 0.02864 grams of vitamin A
1000 mg
88.02 mg CO2 x
CO2: C: 1 x 12.000 = 12.000
O: 2 x 15.999 = 31.998
43.998 grams in 1 mole of CO2
H2O: H: 2 x 1.0078 = 2.0156
O: 1 x 15.999 = 15.999
18.015 grams in 1 mole of H2O
___________________________________________________________________________
12.000 grams C
0.08802 grams CO2 x
= 0.02401 grams C
43.998 grams CO 2
0.02703 grams H2O x
2.0156 grams H
= 3.024 x 10-3 grams H
18.015 grams H 2 O
grams of O = 0.02864 g vitA – (0.02401 gC + 3.024 x 10-3 g H)
grams of O = 0.02864 g vitamin A – 0.02703
grams of O = 1.61 x 10-3 grams O
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Calculating mass % of each element in vitamin A
0.02401 grams C
x 100 = 83.83% carbon
0.02864 grams vitamin A
3.024 x 10 -3 grams H
x 100 = 10.56 % hydrogen
0.02864 grams vitamin A
1.61 x 10 -3 grams O
x 100 = 5.62 % oxygen
0.02864 grams vitamin A
Math check: 83.83% + 10.56% + 5.62% = 100.01 (close enough!!!!)
When given the chemical formula for a compound, you can see how easy it is to determine
mass percent of each element present in the compound. What if only the mass percents were
given? Could you use that information in reverse to predict the molecular formula?
We are going to determine the subscripts (the number of moles) of each atom present in the
compound. Chemists will take a sample compound, break it down into its components,
calculate the number of grams, convert that to moles, and then convert the number of moles into
a whole number ratio. This is a valid representation of the formula, however, this is the
empirical formula.
The “problem” with the empirical formula is that it is the smallest whole number ratio of the
elements present in the compound, and therefore it may not be an accurate representation of
the compound itself. (C2H2 is NOT C6H6 even though they have the same empirical formula!!)
But nonetheless, they are useful and can be used to help determine the molecular formula.
Calculating the Empirical Formula:
1. If given %, simply cancel out the % unit and change it to the gram unit (e.g. 75.6% C
would become 75.6 grams C)
2. Using the grams of the element, convert grams to moles using molar mass
3. Divide all answers in number 2 by the lowest number of moles – DO NOT SKIP THIS
STEP!
4. These numbers are the subscripts, but remember – they must be whole numbers (no
partial atoms exist!!) you therefore MAY have to multiply through by the smallest integer
that will turn all numbers into whole numbers. What you do to one number – you must
must must must must do to all the others.
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For example: if you calculated that you have 2 moles of A, 2.5 moles of B, and 1 mole of C,
you will notice that 2.5 is not a whole number. We need to make it so, therefore we
multiply by the SMALLEST number that will make it a whole number. Multiplying by 1
gets us nowhere. What about 2? Yes, that works. BUT we must multiply everything by 2!
Therefore we will have 4 moles of A, 5 moles of B, and 2 moles of C.
Some common decimal values that CANNOT be rounded to the nearest whole number and
MUST be multiplied to make a whole number value. The X signifies that it could be and
whole number integer such as 1.20 or 2.20, or 4.20.
X.20; X.25; X.33; X.50; X.66; X.75
Concept Test
Determine the empirical formula of 2.45 grams of Si combined with 12.4 grams of Cl
Concept Test
Determine the empirical formula of a compound that is 27.3% C and 72.7% oxygen
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How do empirical formulas tie into molecular formulas?? They each contain the same atoms
(e.g. if we have an empirical formula of CH, the molecular formula will also have CH in it – and
nothing else!!) but the difference is, molecular formulas are some whole number multiple of the
empirical formula: CH could mean C2H2 or C4H4 or C6H6. Thus, if we know (are given) the
molar mass of the compound, we can compare our empirical formula mass to the given molar
mass. If they are the same, then we have an empirical formula that is the molecular formula. If
they are different, then we must find the whole number multiple!
Determining the Whole Number Multiple:
grams
(given)
mole
whole-number multiple =
grams
empirical formula mass in
mole
molar mass in
Remember that the whole number multiple applies to ALL the elements in the compound!
Concept Test
A chloride of silicon contains 79.1% by mass Cl. Considering that the element silicon is also
present, what is the % by mass of silicon present in the compound?
What is the empirical formula of the chloride?
If the molar mass is 268.9 g/mole what is the molecular formula?
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Another type of chemical reaction to be aware of is the combustion reaction. Combustion is
used to measure the amount of hydrogen and carbon in an organic substance as it burns in the
presence of oxygen. All of the hydrogen is converted or H2O, while all of the carbon is
converted to CO2. If the organic compound contains a halogen, nitrogen, or oxygen, simple
math (!!) can be used to determine its mass. We will know the mass of the components at the
end of the combustion reaction and we know the mass of the starting compounds.
Conservation of mass tells us that the mass must be conserved. Using addition/subtraction, the
mass of the missing component can be found.
Concept Test
Menthol (MM = 153.6 g/mole) is a compound that contains C, H, and O. When
0.1595 grams of menthol was subjected to combustion analysis, it produced
0.449 g of CO2, 0.184 g of H2O. What is its molecular formula?
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