Ionic bonds form when two atoms have a large difference in electronegativity.
(Electronegativity is the quantitative representation of an atom’s ability to attract an
electron to itself). Although scientists do not have an exact value to signal an ionic
bond, the amount is generally accepted as 1.7 and over to qualify a bond as ionic.
Ionic bonds often occur between metals and salts; chloride is often the bonding salt.
Compounds displaying ionic bonds form ionic crystals in which ions of positive and
negative charges hover near each other, but there is not always a direct 1-1
correlation between positive and negative ions. Ionic bonds can typically be broken
through hydrogenation, or the addition of water to a compound.
Covalent bonds form when two atoms have a very small (nearly insignificant)
difference in electronegativity. The value of difference in electronegativity between
two atoms in a covalent bond is less than 1.7. Covalent bonds often form between
similar atoms, nonmetal to nonmetal or metal to metal. Covalent bonding signals a
complete sharing of electrons. There is usually a direct correlation between positive
and negative ions, meaning that because they share electrons, the atoms balance.
Covalent bonds are usually strong because of this direct bonding.
Polar covalent bonds fall between ionic and covalent bonds. They result when
two elements bond with a moderate difference in electronegativity moderately to
greatly, but they do not surpass 1.7 in electronegativity difference. Although polar
covalent bonds are classified as covalent, they do have significant ionic properties.
They also induce dipole-dipole interactions, where one atom becomes slightly
negative and the other atom becomes slightly positive. However, the slight change in
charge is not large enough to classify it entirely as an ion; they are simply considered
slightly positive or slightly negative. Polar covalent bonds often indicate polar
molecules, which are likely to bond with other polar molecules but are unlikely to
bond with non-polar molecules.
Hydrogen bonds only form between hydrogen and oxygen (O), nitrogen (N) or
fluorine (F). Hydrogen bonds are very specific and lead to certain molecules having
special properties due to these types of bonds. Hydrogen bonding sometimes results
in the element that is not hydrogen (oxygen, for example) having a lone pair of
electrons on the atom, making it polar. Lone pairs of electrons are non-bonding
electrons that sit in twos (pairs) on the central atom of the compound. Water, for
example, exhibits hydrogen bonding and polarity as a result of the bonding. This is
shown in the diagram below.
Because of this polarity, the oxygen end of the molecule would repel negative
becomes slightly positive, would repel positive atoms (like other hydrogen atoms)
and attract negative atoms (such as oxygen atoms). This positive and negative
attraction system helps water molecules stick together, which is what makes the
boiling point of water high (as it takes more energy to break these bonds between
In addition to the four types of chemical bonds, there are also three categories
bonds fit into: single, double, and triple. Single bonds involve one pair of shared
electrons between two atoms. Double bonds involve two pairs of shared electrons
between two atoms, and triple bonds involve three pairs of shared electrons between
two atoms. These bonds take on different natures due to the differing amounts of
electrons needed and able to be given up.
Now, let’s look at determining what types of bonds we see in different
compounds. We’ve already looked at the bonds in H
O, which we determined to be
(also known as Nitric acid)
look like; first we can decide whether the bonds are covalent, polar covalent, ionic, or
hydrogen. Then, we can determine if the bonds are single, double, or triple.
In order to decide whether the bonds are covalent, polar covalent, ionic or
hydrogen, we need to look at the types of elements seen and the electronegativity
values. We look at the elements and see hydrogen, nitrogen, and oxygen—no metals.
This rules out ionic bonding as a type of bond seen in the compound. Then, we would
look at electronegativity values for nitrogen and oxygen. Oftentimes, this information
can be found on a periodic table, in a book index, or an educational online resource.
The electronegativity value for oxygen is 3.5 and the electronegativity value for
nitrogen is 3.0. The way to determine the bond type is by taking the difference
between the two numbers (subtraction). 3.5 – 3.0 = 0.5, so we can determine that the
bond between nitrogen and oxygen is a covalent bond. We can also determine, from
past knowledge, that the bond between oxygen and hydrogen is a hydrogen bond as it
was in water.
Now, we need to count the electrons and draw the diagram for HNO
. For more
drawing the Lewis structures, please see the page on Lewis Structures. This process
combines both of these in order to determine the structure and shape of a molecule of
First, we determine that N follows the octet rule, so it needs eight surrounding
electrons. This is important to keep in mind as we move forward. Next we count up
how many valence electrons the compound has as a whole. H gives us 1, N gives us
5, and each O gives us 6. We can discern this from looking at the tops of the columns
in the periodic table (see above). We then add these numbers together (3 x 6 = 18, + 1
= 19, + 5 = 24), and we get 24 electrons that we need to distribute throughout the
molecule. First, we need to draw the molecule to see how many initial bonds we’ll be
putting in. Our preliminary structure looks like this:
Now, we can count how many electrons we have used by counting 2 electrons
electrons. 24 – 8 = 16 electrons that we need to distribute. In order to correctly place
the rest of the electrons, we need to determine how many electrons each atom needs
to be stable.
The central atom, N, has three bonds attached (equivalent of 6 electrons) so it
needs 2 more electrons to be stable. The O to the right has one bond (two electrons)
so it needs 6 more to be stable. The O above the N has one bond (two electrons) so it
also needs 6 electrons to be stable. The O to the left of the N is bonded both to N and
to H, so it has two bonds (4 electrons); therefore, it needs 4 more electrons to be
stable. We add up the total amount of electrons needed, 2 + 6 + 6 + 4 = 18, and see
that we need 18 electrons to stabilize the compound. We know this is not possible,
since we only have 16 available electrons. When this happens, we need to insert a
double bond in order to resolve the problem of lack of electrons. This is because,
although we count each bond as 2 electrons, the elements joined together in the bond
are actually sharing the electrons. Therefore, when we count out the bonds, we are
counting some electrons twice because they are shared. This is normal and expected,
and resolves not having enough valence electrons. Now, we need to decide where to
put the double bond in this compound. We know that the double bond cannot go
between O and H, because H does not have enough room to accept another electron.
Therefore, we know we must place the bond between N and O. You might be
thinking, how do I decide where to put the bond? In this particular example, we can
place the bond either between the top O and N, or the right O and N. This is because
We are going to keep the bond between N and the right O in our example.
After we add in the bond, we subtract two more electrons from our available
electrons (16) and are left with 14 electrons to distribute. Now we need to make sure
we have the correct number of electrons. After placing in the double bond, N is now
stable because it has 4 bonds (8 electrons) surrounding it. It does not need any
additional electrons. The top O (above N) needs 6 electrons, the right O now only
needs 4 electrons (because it has a double bond now, which is 4 electrons), and the
left O still needs 4 electrons to become stable. We add these numbers together, 6 + 4
+ 4 = 14, and we see that 14 is the number of electrons we have, so we can go ahead
and distribute them, like this:
Now, our compound is stable with appropriately distributed valence electrons.
double bond (N==O). 
The early innovation of Periodic law by a Russian chemist, Dmitry I.
Mendeleev in the mid-19th century, has been of great value in the growth of
chemistry. In chemistry, Periodic law is the arrangement of theelements in order of
increasing atomic number, that is, the total number of protons in the atomic nucleus
and their physical and chemical properties recur in a pattern with the increasing
atomic number. In the periodic table, the horizontal rows, known as the “periods”
exemplify these relations. In fact, until the second decade of the 20th century, it was
not documented that the order ofelements in the periodic system is that of their
Classification became necessary due to the increase in the number of elements
discovered. In 1817, J.W. Dobereiner, a German chemist , explained that the atomic
weight of strontium rests in the middle between that of calcium and barium, and few
years later he demonstrated that other such “triads” such as chlorine, bromine, and
iodine and lithium, sodium, and potassium are also present. Other scientists,
during1827 and 1858 developed analogous relationships which extended more than
the triads of elements, fluorine being added to the halogens and magnesium to the
alkaline-earth metals, while oxygen, sulfur, selenium, and tellurium were classed as
one family and nitrogen, phosphorus, arsenic, antimony, and bismuth as one more
element family. De Chancourtois, a French scientist, in 1862 anticipated a
classification of the elements based on the new values of atomic weights. It was
finally the Russian chemist, Mendeleev, who proposed the periodic law, which
indicated that the elements arranged in the increasing order of atomic weights show a
periodic change of properties. In 1869, the first periodictable was tabulated by
Mendeleev. It had 17 columns, with two nearly complete periods of elements,
frompotassium to bromine and rubidium to iodine, lead by two fractional periods of
seven elements each (lithium to fluorine and sodium to chlorine), and three
What is the importance of the Periodic law?
- Innovation of New Elements: Through Mendeleyev’s efforts in 1871, the
grand significance of the periodic law was made apparent in predicting that the
properties of the 17 elements could be interrelated with those of other elements by
relocating the 17 to new positions from those shown by their atomic weights. The
subsistence of many of the properties of uninvented elements of that time like eka-
boron, eka-aluminum, and eka-silicon, now identified with the elements scandium,
gallium, and germanium, was also predicted by Mendeleev. Likewise, following the
discovery of helium and argon, the periodic law allowed the discovery of the
subsistence of neon, krypton, xenon, and radon. Besides, the absence of element 72
was anticipated, from its position in the periodic system, to be alike to zirconium in
its properties rather than to the rare earths, was shown by Niels Bohr, a Danish
physicist. In 1922, other scientists also examined zirconium ores following Bohr’s
- Implication of atomic numbers: A number of of the elements in the
Mendeleyev periodic tables were necessary to arrange elements according to their
atomic weight . For example, in the pair’s of argon and potassium, cobalt and nickel,
and tellurium and iodine, the first element had the previous place in the periodic
system but with a bigger atomic weight. When the structure of the atom was studied,
the explanation to this complexity was solved. Research done by Ernest Rutherford
on the scattering of alpha particles by the nuclei of heavy atoms, in1910, helped the
prediction of the nuclear electrical charge. Approximately, the ratio of the nuclear
charge to that of the electron was noted to be one-half the atomic weight. Another
scientist, in 1911 recommended that this quantity, the atomic number, might be
recognized with the ordinal number of the element in the periodic table. In 1913, this
proposal was vividly established by the English physicist, H.G.J. Moseley’s
measurements of the wavelengths of the characteristic X-ray spectral lines of many
elements, which showed that the wavelengths relied in a usual way on the atomic
numbers indistinguishable with the ordinal numbers of the periodic table elements.
How did the periodic law evolve ?
Starting in 1913, thorough knowledge of the the elements and their properties
had improved. Followed by the exclusion principle by the Austrian theoretical
physicist, Wolfgang Pauli in 1925, the invention of the spin of the electron by George
E. Uhlenbeck and Samuel Goudsmit in 1925, and the development of quantum
mechanics by Werner Heisenberg were the significant advanced inventions which
developed the periodic law. The development of the electronic theory of valence and
molecular structure, beginning with the postulate of the shared electron pair by the
chemist, Gilbert N. Lewis in 1916, also played a vital role in elucidating the periodic
"If all the elements are arranged in the order of their atomic weights, a periodic
Dmitri Mendeleev, Principles of Chemistry, Vol. 2, 1902, P. F. Collier, p17.
The periodic table we use today is based on the one devised and published by
Dmitri Mendeleev in 1869.
Mendeleev found he could arrange the 65 elements then known in a grid or
table so that each element had:
1. A higher atomic weight than the one on its left. For example, magnesium
(atomic weight 24.3) is placed to the right of sodium (atomic weight 23.0):
2. Similar chemical properties to other elements in the same column - in other
chemistry. And more than that, Mendeleev saw that his table was incomplete - there
were spaces where elements should be, but no-one had discovered them.
Just as Adams and Le Verrier could be said to have discovered the planet
Neptune on paper, Mendeleev could be said to have discovered germanium on paper.
He called this new element eka-silicon, after observing a gap in the periodic table
between silicon and tin:
Similarly, Mendeleev discovered gallium (eka-aluminum) and scandium (eka-
Although Mendeleev had made a crucial breakthrough, he made little further
progress. With the benefit of hindsight, we know that Mendeleev's periodic table was
underpinned by false reasoning. Mendeleev believed, incorrectly, that chemical
properties were determined by atomic weight. Of course, this was perfectly
reasonable when we consider scientific knowledge in 1869.
In 1869 the electron itself had not been discovered - that happened 27 years
later, in 1896.
In fact, it took 44 years for the correct explanation of the regular patterns in
Mendeleev's periodic table to be found. 
The law of conservation of mass or principle of mass conservation states that
for any system closed to all transfers of matter andenergy, the mass of the system
must remain constant over time, as system mass cannot change quantity if it is not
added or removed. Hence, the quantity of mass is "conserved" over time. The law
implies that mass can neither be created nor destroyed, although it may be rearranged
in space, or the entities associated with it may be changed in form, as for example
when light or physical work is transformed into particles that contribute the same
mass to the system as the light or work had contributed. The law implies (requires)
that during anychemical reaction, nuclear reaction, or radioactive decay in an isolated
system, the total mass of the reactants or starting materials must be equal to the mass
of the products.
The concept of mass conservation is widely used in many fields such as
chemistry, mechanics, and fluid dynamics. Historically, mass conservation was
discovered in chemical reactions by Antoine Lavoisier in the late 18th century, and
was of crucial importance in the progress from alchemy to the modern natural science
The closely related concept of matter conservation was found to hold good in
chemistry to such high approximation that it failed only for the high energies treated
by the later refinements of relativity theory, but otherwise remains useful and
sufficiently accurate for most chemical calculations, even in modern practice.
In special relativity, needed for accuracy when large energy transfers between
systems is involved, the difference between thermodynamically closed and isolated
systems becomes important, since conservation of mass is strictly and perfectly
upheld only for so-called isolated systems, i.e. those completely isolated from all
exchanges with the environment. In this circumstance, the mass–energy equivalence
theorem states that mass conservation is equivalent to total energy conservation,
which is the first law of thermodynamics. By contrast, for a thermodynamically
closed system (i.e., one which is closed to exchanges of matter, but open to
exchanges of non-material energy, such as heat and work, with the surroundings)
mass is (usually) only approximately conserved. The input or output of non-material
energy must change the mass of the system in relativity theory, although the change
is usually small, since relatively large amounts of such energy (by comparison with
ordinary experience) carry only a small amount of mass (again by ordinary standards
In special relativity, mass is not converted to energy, since mass and energy
cannot be destroyed, and energy in all of its forms always retains its equivalent
amount of mass throughout any transformation to a different type of energy within a
system (or translocation into or out of a system). Certain types of matter (a different
concept) may be created or destroyed, but in all of these processes, the energy and
mass associated with such matter remains unchanged in quantity (although type of
energy associated with the matter may change form).
In general relativity, mass (and energy) conservation in expanding volumes of
space is a complex concept, subject to different definitions, and neither mass nor
energy is as strictly and simply conserved as is the case in special relativity and in
Minkowski space. For a discussion, see mass in general relativity.
An important idea in ancient Greek philosophy was that "Nothing comes from
nothing", so that what exists now has always existed: no new matter can come into
existence where there was none before. An explicit statement of this, along with the
further principle that nothing can pass away into nothing, is found in Empedocles
(approx. 490–430 BC): "For it is impossible for anything to come to be from what is
not, and it cannot be brought about or heard of that what is should be utterly
A further principle of conservation was stated by Epicurus (341–270 BC) who,
describing the nature of the Universe, wrote that "the totality of things was always
such as it is now, and always will be".
Jain philosophy, a non-creationist philosophy based on the teachings of
Mahavira (6th century BC),
matter cannot be destroyed or created. The Jain text Tattvarthasutra (2nd century AD)
states that a substance is permanent, but its modes are characterised by creation and
destruction. A principle of the conservation of matter was also stated by
It only changes its form, condition, composition, color and other properties and turns
into a different complex or elementary matter".
Mass conservation in chemistry
The principle of conservation of mass was first outlined by Mikhail
Lomonosov (1711–1765) in 1748. He proved it by experiments—though this is
1774. Others whose ideas pre-dated the work of Lavoisier include Joseph Black
(1728–1799), Henry Cavendish(1731–1810), and Jean Rey (1583–1645).
The conservation of mass was obscure for millennia because of the buoyancy
effect of the Earth's atmosphere on the weight of gases. For example, a piece of wood
weighs less after burning; this seemed to suggest that some of its mass disappears, or
is transformed or lost. This was not disproved until careful experiments were
performed in which chemical reactions such as rusting were allowed to take place in
sealed glass ampoules; it was found that the chemical reaction did not change the
weight of the sealed container and its contents. The vacuum pump also enabled the
weighing of gases using scales.
Once understood, the conservation of mass was of great importance in
progressing from alchemy to modern chemistry. Once early chemists realized that
chemical substances never disappeared but were only transformed into other
substances with the same weight, these scientists could for the first time embark on
quantitative studies of the transformations of substances. The idea of mass
conservation plus a surmise that certain "elemental substances" also could not be
transformed into others by chemical reactions, in turn led to an understanding of
chemical elements, as well as the idea that all chemical processes and transformations
(such as burning and metabolic reactions) are reactions between invariant amounts or
weights of these chemical elements.
Following the pioneering work of Lavoisier the prolonged and exhaustive
experiments of Jean Stas supported the strict accuracy of this law in chemical
reactions, even though they were carried out with other intentions. His research
indicated that in certain reactions the loss or gain could not have been more than from
2 to 4 parts in 100,000. The difference in the accuracy aimed at and attained by
Lavoisier on the one hand, and by Morley and Stas on the other, is enormous.
In special relativity, the conservation of mass does not apply if the system is
open and energy escapes. However, it does continue to apply to totally closed
(isolated) systems. If energy cannot escape a system, its mass cannot decrease. In
relativity theory, so long as any type of energy is retained within a system, this
energy exhibits mass.
Also, mass must be differentiated from matter (see below), since matter may
not be perfectly conserved in isolated systems, even though mass is always conserved
in such systems. However, matter is so nearly conserved in chemistry that violations
of matter conservation were not measured until the nuclear age, and the assumption
of matter conservation remains an important practical concept in most systems in
chemistry and other studies that do not involve the high energies typical of
radioactivity and nuclear reactions.
The mass associated with chemical amounts of energy is too small to measure
The change in mass of certain kinds of open systems where atoms or massive
particles are not allowed to escape, but other types of energy (such as light or heat)
are allowed to enter or escape, went unnoticed during the 19th century, because the
change in mass associated with addition or loss of small quantities of thermal or
radiant energy in chemical reactions is very small. (In theory, mass would not change
at all for experiments conducted in isolated systems where heat and work were not
allowed in or out.)
The theoretical association of all energy with mass was made by Albert
Einstein in 1905. However Max Planck pointed out that the change in mass of
systems as a result of extraction or addition of chemical energy, as predicted by
Einstein's theory, is so small that it could not be measured with available instruments,
for example as a test of Einstein's theory. Einstein speculated that the energies
associated with newly discovered radioactivity were significant enough, compared
with the mass of systems producing them, to enable their mass-change to be
measured, once the energy of the reaction had been removed from the system. This
later indeed proved to be possible, although it was eventually to be the first artificial
nuclear transmutation reaction in 1932, demonstrated by Cockcroft and Walton, that
proved the first successful test of Einstein's theory regarding mass-loss with energy-
The conservation of relativistic mass implies the viewpoint of a single observer
(or the view from a single inertial frame) since changing inertial frames may result in
a change of the total energy (relativistic energy) for systems, and this quantity
determines the relativistic mass.
The principle that the mass of a system of particles must be equal to the sum of
their rest masses, even though true in classical physics, may be false in special
relativity. The reason that rest masses cannot be simply added is that this does not
take into account other forms of energy, such as kinetic and potential energy, and
massless particles such as photons, all of which may (or may not) affect the total
mass of systems.
For moving massive particles in a system, examining the rest masses of the
various particles also amounts to introducing many different inertial observation
frames (which is prohibited if total system energy and momentum are to be
conserved), and also when in the rest frame of one particle, this procedure ignores the
momenta of other particles, which affect the system mass if the other particles are in
motion in this frame.
For the special type of mass called invariant mass, changing the inertial frame
of observation for a whole closed system has no effect on the measure of invariant
mass of the system, which remains both conserved and invariant (unchanging), even
for different observers who view the entire system. Invariant mass is a system
combination of energy and momentum, which is invariant for any observer, because
in any inertial frame, the energies and momenta of the various particles always add to
the same quantity (the momentum may be negative, so the addition amounts to a
subtraction). The invariant mass is the relativistic mass of the system when viewed in
the center of momentum frame. It is the minimum mass which a system may exhibit,
as viewed from all possible inertial frames.
The conservation of both relativistic and invariant mass applies even to
systems of particles created by pair production, where energy for new particles may
come from kinetic energy of other particles, or from one or more photons as part of a
system that includes other particles besides a photon. Again, neither the relativistic
nor the invariant mass of totally closed (that is, isolated) systems changes when new
particles are created. However, different inertial observers will disagree on the value
of this conserved mass, if it is the relativistic mass (i.e., relativistic mass is conserved
by not invariant). However, all observers agree on the value of the conserved mass if
the mass being measured is the invariant mass (i.e., invariant mass is both conserved
The mass-energy equivalence formula gives a different prediction in non-
isolated systems, since if energy is allowed to escape a system, both relativistic mass
and invariant mass will escape also. In this case, the mass-energy equivalence
formula predicts that the change in mass of a system is associated
with the change in its energy due to energy being added or subtracted:
This form involving changes was the form in which this famous equation was
explained simply if the mass of the energy added or removed from the system, are
taken into account.
The formula implies that bound systems have an invariant mass (rest mass for
the system) less than the sum of their parts, if the binding energy has been allowed to
escape the system after the system has been bound. This may happen by converting
system potential energy into some other kind of active energy, such as kinetic energy
or photons, which easily escape a bound system. The difference in system masses,
called a mass defect, is a measure of the binding energy in bound systems – in other
words, the energy needed to break the system apart. The greater the mass defect, the
larger the binding energy. The binding energy (which itself has mass) must be
released (as light or heat) when the parts combine to form the bound system, and this
is the reason the mass of the bound system decreases when the energy leaves the
The total invariant mass is actually conserved, when the mass of the binding
energy that has escaped, is taken into account.
Exceptions or caveats to mass/matter conservation
Matter is not perfectly conserved
The principle of matter conservation may be considered as an approximate
physical law that is true only in the classical sense, without consideration of special
relativity andquantum mechanics. It is approximately true except in certain high
A particular difficulty with the idea of conservation of "matter" is that "matter"
is not a well-defined word scientifically, and when particles that are considered to be
"matter" (such as electrons and positrons) are annihilated to make photons (which are
often not considered matter) then conservation of matter does not take place over
time, even within isolated systems. However, matter is conserved to such an extent
that matter conservation may be safely assumed in chemical reactions and all
situations in which radioactivityand nuclear reactions are not involved.
Even when matter is not conserved, the collection of mass and energy within
the system are conserved.
Open systems and thermodynamically closed systems
Mass is also not generally conserved in open systems. Such is the case when
various forms of energy are allowed into, or out of, the system (see for example,
binding energy). However, again unless radioactivity or nuclear reactions are
involved, the amount of energy escaping such systems as heat, work, or
electromagnetic radiation is usually too small to be measured as a decrease in system
The law of mass conservation for isolated systems (totally closed to all mass
in modern physics. The reason for this is that relativistic equations show that even
"massless" particles such as photons still add mass and energy to isolated systems,
allowing mass (though not matter) to be strictly conserved in all processes where
energy does not escape the system. In relativity, different observers may disagree as
to the particular value of the conserved mass of a given system, but each observer
will agree that this value does not change over time as long as the system is isolated
(totally closed to everything).
In general relativity, the total invariant mass of photons in an expanding
volume of space will decrease, due to the red shift of such an expansion (see Mass in
general relativity). The conservation of both mass and energy therefore depends on
various corrections made to energy in the theory, due to the changing gravitational
potential energy of such systems.