Download Atoms and Ions

People asked the question – for thousands of years: What is matter made of? The ancient
philosopher Democritus believed that the universe consisted of an infinite number of
indestructible atoms moving in empty space. He thought that atoms differed in size, weight
and shape, but that all atoms were made from the same material. Most ancient thinkers
disagreed with him and sided with Aristotle, who believed that matter was continuous,
infinitely divisible, and made up of four elements: earth, air, fire, and water.
Things changed in the 18th and 19th century when experimentation allowed scientists to
accumulate data that could explain atomic theory.
1803
Dalton’s Atomic Theory:
1. All matter consists of atoms: which are tiny indivisible particles of an element that cannot
be created or destroyed. Democritus, 2000 years previous to Dalton came up with the
idea of atoms)
2. Atoms of one element cannot be converted into atoms of another element by chemical
means. In chemical reactions, atoms from elements recombine to form different
substances.
C5H12
(l) +
8O2 (g)
5 CO2 (g) +
6 H2O (l)
Carbon atoms and hydrogen atoms from the molecule pentane combine with
oxygen atoms in the molecule oxygen to form the molecules carbon dioxide and
water.
3. Atoms of an element are identical in mass and other properties and are different from
atoms of another element.
4. Compounds result from the chemical combination of specific ratios of atoms of different
elements.
SO3 / SO2 ; CO / CO2 ; FeCl2 / FeCl3 ; UF3 / UF4 / UF6
Dalton’s Atomic Theory parallels the mass laws: all matter consists of indivisible atoms
of a fixed and unique mass, mass remains constant during a chemical reaction because
the atoms form new combinations, each compound has a fixed mass due to the fact that a
compound is composed of a fixed number of each type of atom, different compounds
made up of the same elements exhibit multiple proportions because they exist in whole
number ratios.
1897
J. J. Thomson: the atom as plum pudding
In the 19th century, various investigators were using discharge tubes to study the conduction of
electricity in gasses.
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The discharge tube is a sealed glass tube that contains two metal plates called electrodes. One
electrode, the anode, is given a positive charge. The other electrode is the cathode, and is given
a negative charge. When the difference in charge between the electrodes was sufficiently large
enough, a beam of radiation, called a cathode ray, is emitted by the cathode. Under vacuum,
the ray can be detected when it impacts a fluorescent screen. It was noticed, that if a positively
charged object was placed near the beam that the beam bended towards the positive plate.
Thus, since opposites attract, the beam must be negatively charged!
J. J. Thomson had discovered electrons (although he called them “corpuscles”). Thomson found
that when different cathodes were used, the electron beam generated was the same, thus he
concluded that electrons are a fundamental and universal constituent of matter. He was unable
to determine the mass or charge of an electron separately, but he found the ratio of the
electron’s charge to its mass:
e
= 1.759 x 108 C/g
m
e = magnitude of the electron charge in Coulombs (C)
m = mass of the electron in grams
Thomson did not know how positive charge was dispersed in an atom, but since atoms were
found to be neutral species, the negatively charged electrons were counteracted by positively
charged species. Thomson believed the atom to appear something like:
+9 uniformly distributed positive
charge
-9 “corpuscles”
of negative
charge
_
_
_
_
_
_
_
_
_
Controlled electron beams are now used in TV picture tubes (see additional material at
the end of these notes), x-ray tubes, oscilloscopes, computer monitors, and electron
microscopes.
2
Cathode ray tube as beam bends towards a positive plate
Cathode ray tube as with beam uninterrupted
1909:
Robert Millikan: the oil-drop experiment
Millikan determined the charge of an electron. He used an apparatus, as shown below, to
produce tiny oil droplets. Very fine oil droplets were sprayed into a chamber and then were
allowed to fall between two charged plates where they were then observed, visually. The air
inside the chamber was exposed to x-rays, which displace electrons from air molecules resulting
in both negatively charged electrons and positively charged air molecules. These species come
in contact with the oil droplets. Some of the oil droplets pick up the electrons, thus acquiring a
net negative charge; other droplets pick up the positively charged fragments. Each charged
droplet is under the influence of two forces: gravity, and the force exerted by the charged
plates. By turning on the electric field, Millikan could control the rate of the droplet falling. He
could make the drop fall more slowly, rise, or even suspend it in midair! (Notice the lower plate
is negatively charged, and if the droplet is negatively charged, the plate will repel that oil
droplet). Millikan chose a particular oil droplet and measured its rate of falling with the electric
field turned on. Then, with the electric field turned off, the rat of falling of an oil droplet was
measured again. A comparison of the two rates allowed Millikan to calculate the droplet’s
charge.
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After studying many droplets, Millikan confirmed that the charges on the droplets were always
some whole number multiple of a minimum charge. Different oil droplets are able to pick up
different numbers of electrons, so the minimum charge must be due to a single electron. He
determined that no charge existed that was less than 1.602 x 10-19 C. Thus, that must be the
charge of a single electron.
Using J.J. Thomson’s equation :
e
= 1.759 x 108 C/g
m
where e now equals: 1.602 x 10-19 C, Millikan could solve for the mass of a single electron:
9.109x 10-28 g.
Since electrons are negatively charged species we can place negative charges in front of
(-)1.759 x 108 C/g and (-) 1.602 x 10-19 C.
1910
Ernest Rutherford: the gold foil experiment
Thomson’s work with cathode rays and the discovery of positive ion beams led scientists to
believe that every atom contains electrons and that the removal of those electrons leaves behind
a positive ion that contains most of the atom’s mass. The inner structure of that atom was still a
mystery. J.J. Thomson has his theory (the plum pudding model), and soon, Rutherford, with
the help of his co-worker Hans Geiger, came up with their own. Early in the 20th century
scientists had discovered radioactivity. Rutherford used one of these radioactive particles, the
alpha particle, in his experiments.
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They performed the experiment where a very thin piece of gold foil was bombarded with alpha
particles, and then they observed the path that the alpha particles took when they emerged from
the foil. Rutherford found that most of the particles went straight through the foil undeflected,
or deflected only slightly. These results were consistent with Thomson’s model of the atom. If
the atoms in the foil consisted of electrons embedded in a homogenous sphere of positive
electricity, then the electric field acting on the positively charged alpha particles would be
evenly balanced, and the particles would pass through the foil with very little deflection. The
embedded electrons were too small to cause deflections, it would be like a ping pong ball trying
to deflect a baseball thrown by Randy Johnson!
If Rutherford and Geiger had stopped there, they would have drawn the same conclusion about
the structure of the atom that J.J Thomson had – but they persevered, and several experiments
later they began to see alpha particles deflected backwards!
It seemed that in these experiments, the alpha particles were being repelled by something small,
dense, and positive within the gold atoms. Rutherford therefore concluded that there was a
region of positive charge within the atom (the nucleus), and as such, only when alpha particles
encounter this mass will they be deflected. And since most alpha particles passed through the
atom without being deflected, the size of this region must be small with respect to the overall
size of the atom. And since J.J. Thomson and Millikan had already shown that the electrons
could easily be removed from the atom, those electrons must have been separate from the
positive mass in the atom. Thus, Thomson’s model had to be abandoned and a new model, the
nuclear atom model was adopted.
In the nuclear atom model, the atom consisted of a small, but massive, positively charged
nucleus surrounded by negative electrons. Why massive? Almost all of the mass of the atom is
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retained in the nucleus. He called the positive particles protons, and predicted nuclear charge
with remarkable accuracy. Rutherford’s model explained the charged nature of matter, but it
did not account for the atom’s mass. It would be another 22 years before James Chadwick
discovered another piece of the atomic puzzle, the neutron – an uncharged species that also
resides in the nucleus.
_
_
_
_
_
++
+ +
+
_
The atom is electrically neutral, spherical species that contains a positively charged nucleus
surrounded by one or more negatively charged electrons. The electrons move rapidly around
the nucleus and are held there in space by attraction with the positively charged nucleus (much
like the planets orbiting the sun). The nucleus, which takes up 1 ten-trillionth of the volume but
makes up 99.97% of the atom’s mass is incredibly dense. The atom’s total diameter is about
10,000 times the diameter of the nucleus!
The nucleus is composed of neutrons and protons. The protons are positively charged and the
neutrons have no charge at all, they just contribute to the overall mass of the atom. The
magnitude of the charge of a proton is equal to that of an electron, but the electron is negatively
charged. An atom is neutral because the number of protons ALWAYS equals the number of
electrons.
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Atomic Notation: 3 Types
A = Atomic Mass = protons + neutrons
Z = Atomic Number = protons
X = Atomic Symbol = the elements letter designation
Type 1
Type 2
Type 3
A
Z
A
X
12
6
X
12
C
This type came first and gives the most
information
C
This type came second as the first type is
redundant. You do not need to indicate that
carbon has an atomic number of 6, all
carbon atomic numbers are 6.
X – A C -12
This type came last and is the easiest to type,
and still relays all the info you need. This
symbol is spoken, “carbon twelve.”
All atoms of a particular element have the same atomic number, and different elements have
different atomic numbers. On the periodic table, the atomic number is indicated over the top of
the element.
The total number of protons and neutrons together is the atomic mass of an element. On the
periodic table, the atomic mass is found beneath the element. Thus, given a particular atom’s
mass and knowing that the number of protons remains constant, the number of neutrons can
then be calculated.
All atoms of a particular element have the same number of protons (atomic number) but do not
have the same mass. All carbon atoms have 6 protons, but only 98.89% of carbon atoms have 6
neutrons. A very small percentage of carbon atoms have 7 neutrons.
Isotopes of an element are atoms that have the same number of protons and electrons but differ
in the number of neutrons.
Carbon isotopes include 12C, 13C, 14C. Each of these species have 6 protons and 6 electrons, they
are neutral species, but the number of neutrons changes.
Chemical properties, if you remember, depend on the electrons. Since each isotope has the
same number of electrons, their chemical properties are nearly identical, even though they differ
in mass.
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Particle
p
n
e
Charge
+
n/a
-
Location
nucleus
nucleus
orbit-
AMU
1
1
0
Even though there are only 109 elements on the periodic table, the existence of isotopes has
given rise to more that 1450 different atomic species!
Most elements have one isotope that more predominant that the others, (e.g. carbon). But
others have significant percentages of their isotopes existing in nature: Mg consists of 78.7%
24Mg, 10.13% 25Mg, and 11.17% 26Mg. Other elements, such as Xe, and Sn, have 9 and 10 natural
isotopes, respectively.
Hydrogen is the only element whose isotopes have been given special names. Most abundant is
hydrogen, whose nucleus consists of a single proton. The other isotopes are deuterium, with 1
neutron and 1 proton, and tritium, which has 2 neutrons and 1 proton.
Using a mass spectrometer, the isotopic make-up of an element can be determined, as well as
the relative abundance of each isotope. Each isotope contributes a certain percentage of its
mass to generate the overall atomic weight of a particular element. It is recognized that each
isotope contributes to the average mass in proportion to its abundance. Atomic masses
obtained in this way are said to be weighted averages.
For example: Gallium has two naturally occurring isotopes, 69Ga which has a mass of
68.9257 Da, and 71Ga which has a mass of 70.9249. The percent abundance of 69Ga is
60.27% while 71Ga makes up 39.73% of the element. Calculate the chemical atomic
weight of Ga.
Isotope
69Ga
71Ga
Isotopic Mass (Da)
68.9257
70.9249
*
*
Fractional Abundance
0.6027
0.3973
=
=
Product (Da)
41.54
28.18
________________
Sum =
Examining the periodic table we see that this is the atomic mass of Ga!
8
69.72 Da
Reassessing Modern Atomic Theory:
1. all matter is composed of atoms: but they are divisible and are composed of even smaller
units; protons, electrons, and neutrons, but the atom is the smallest body that retains the
unique identity of the element.
2. atoms of one element cannot be converted into atoms of another element by chemical
means: this can occur in nuclear reactions but not in a chemical reaction!
3. all atoms of an element have the same number of protons and electrons, they only vary in
their number of neutrons. It is the electrons that determine the chemical properties of the
element. We treat a sample of the element as having an average mass but still being
composed of the same species.
4. compounds are formed by the chemical combination of two or more elements in a
specific ratio.
The periodic table is your answer key. Many of the things that we will talk about in class can be
found directly on the periodic table! The atomic mass of an element, the number of protons, the
number of electrons, and thus, through some simple math, the number of neutrons. Given the
talk about weighted averages, when you calculate the weighted average mass of an element,
you can check your answer by looking on the periodic table!
1. each element has a box that contains its atomic number, atomic symbol, and atomic mass.
The elements are arranged by increasing atomic number (increasing number of protons)
and with few exceptions this translates into being arranged by increasing atomic mass
(notice Co is heavier than Ni, Th is heavier than Pa, U is heavier than Np)
2. the elements are arranged in periods (by rows) and also by groups (columns). Each row,
or period, is numbered 1-7 and each group, or column, is numbered from 1-8 and
includes either the letter A or B.
3. the groups with the letter A are considered to be the main group, or representative
elements. The B group elements are known as the transition or the inner transition
elements. The inner transition elements run from Ce to Lu and Th to Lr.
There is a clear marking on the periodic table which indicates the division between the metals
and non-metals. This staircase line that runs from boron (B) down to polonium (Po) divides the
metals, which are on the left hand side of the line, from the non-metals, which are to the right of
the line. All elements which share a SIDE with the line (NOT a corner!!) are known as the
metalloids, which retain properties of both metals and non-metals.
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Additionally, groups have been termed families, and just like real people families, elements in a
family share characteristics. These elements have similar chemical properties while the
elements in a period (row) have very different chemical properties.
Group IA – except for H: alkali metals
Group IIA – alkaline earth metals
Group VIIA – halogens
Group VIIIA – noble or inert gasses
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 gasses (Group VIIIA). Other atoms will combine with
themselves, such as O2 (probably heard of that!), N2, H2, S8, Cl2, F2, Br2, I2. 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
2. transferring the electrons from one atom to another to form an ionic bond
Chemical Formulas:
1. empirical formulas: the simplest 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.
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 molecular
formula.
2. molecular formula: shows how many atoms of each element are 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.
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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.
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:
Bonding between a metal and a nonmetal or between two charged species (e.g. polyatomic
ions).
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. And just like their 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. When a metal loses an electron (or 2 or 3) it then has
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more protons (positively charged species) than electrons (negatively charged species). Group
1A ALL turn into +1 ions – they each have 1 more proton then electron. Group IIA ALL turn
into +2 ions – they each have 2 more protons than electrons. Group IIIA turn into +3 ions – they
each have 3 more protons than electrons. 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 . . . 
So what about the negative species? 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
we are not going to pay them too much attention when talking about ion formation. They are a
reference point for us on the periodic table – the ideal.
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.
So, working backwards away from the noble gases the periodic table looks like this:
But why? Remember, the noble gasses 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).
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
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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.
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
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, only ions!!
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 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: H2O, NH3, CH4. Generally speaking, covalent bonds occur
between NONMETALS. (remember hydrogen is not a metal, sometimes it just gets put over
group IA since it has 1 electron to lose too!).
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.
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o
o
o
o
poly = many
atomic = atoms
ion = charged particle
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.
o the list of polyatomic ions you must commit to memory. (memorize the ones in
bold, the rest you might encounter on homeworks or in 141, but do not need to
memorize)
Formula
Name
Formula
Name
NO3
CNMnO4OHSO42PO43C2H3O2- or
CH3COONH4+
NO2CrO42Cr2O72SCNSO32PO33HPO42H2PO4-
nitrate
cyanide
permanganate
hydroxide
sulfate
phosphate
acetate
acetate
ammonium
nitrite
chromate
dichromate
thiocyanate
sulfite
phosphite
hydrogen phosphate
dihydrogen phosphate
CO3
ClO4ClO3ClO2ClOIO4IO3IO2IOBrO4BrO3BrO2BrOHCO3HSO4HSO3HS-
carbonate
perchlorate
chlorate
chlorite
hypochlorite
periodate
iodate
iodite
hypoiodite
perbromate
bromate
bromite
hypobromite
hydrogen carbonate
hydrogen sulfate
hydrogen sulfite
hydrogen sulfide
-
2-
Monoatomic Ions to commit to memory (from the Periodic Table)
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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→ CaCl2 (1’s are always omitted for clarity)
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 not
requiring you to learn this method . . . just be aware that it IS out there and might be
helpful in the future so just be familiar with it! Instead, 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.
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For the ous and ic naming: the higher charged ion for a particular element
gets the ic ending while the lower charged ion for a particular element gets
the ous ending
Fe+2 vs. Fe+3
Fe+2 has the lower charge and would be named with the ous ending while
Fe+3 ion has the higher charge and would be named with the ic ending.
The other component to the name is the root. And unfortunately it is not
ironous and ironic. (Of course not!!) These special ions use the Latin base
names given to them.
Using the roman numeral method of naming Fe+2 would be iron(II) and
Fe+3 would be named iron(III).
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 NO 3 If
you just wrote NO33 it obviously looks like a nitrogen atom bonded to 33
oxygens!! BAD!!
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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).
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
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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
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.
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Rules for naming binary covalent compounds:
1. 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.
2. If both elements are in the same group, the one with the higher period number is
named first (remember period is the row)
3. 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.
Number
1
2
3
4
5
6
7
8
9
10
Prefix
mono
di
tri
tetra
penta
hexa
hepta
octa
nona
deca
4. The second element is named as its root with the suffix –ide.
5. The suffix will 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!)
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