The chemical properties of an element are determined by the charge of its
nucleus, i.e. by the number of protons. This number is called the atomic number and
is denoted by the letter Z.
Definition: Atomic number (Z)
The number of protons in an atom.
The mass of an atom depends on how many nucleons its nucleus contains. The
nucleons, i.e. the total number of protons plus neutrons, is called the atomic
and is denoted by the letter A.
Definition: Atomic mass number (A)
The number of protons and neutrons in the nucleus of an atom.
Standard notation shows the chemical symbol, the atomic mass number and the
atomic number of an element as follows:
For example, the iron nucleus which has 26 protons and 30 neutrons, is
where the total nuclear charge is Z = 26 and the mass number A = 56. The
Don’t confuse the notation we have used above, with the way this information
on the Periodic Table. On the Periodic Table, the atomic number usually
appears in the
top lefthand corner of the block or immediately above the element’s symbol.
below the element’s symbol is its relative atomic mass. This is not exactly the
the atomic mass number. This will be explained in section 3.5. The example of
iron is used again below.
You will notice in the example of iron that the atomic mass number is more or
neutrons (i.e. Z = n), will have a mass approximately equal to n u. The reason is that a
C-12 atom has 6 protons, 6 neutrons and 6 electrons, with the protons and neutrons
having about the same mass and the electron mass being negligible in comparison.
A large local charge separation usually results when a shared electron pair is
donated unilaterally. The three Kekulé formulas shown here illustrate this condition.
In the formula for ozone the central oxygen atom has three bonds and a full
charged. The overall charge of the ozone molecule is therefore zero. Similarly,
nitromethane has a positive-charged nitrogen and a negative-charged oxygen, the
total molecular charge again being zero. Finally, azide anion has two negative-
charged nitrogens and one positive-charged nitrogen, the total charge being minus
In general, for covalently bonded atoms having valence shell electron octets, if the
positive charge. If the number of covalent bonds to an atom is less than its normal
valence it will carry a negative charge. The formal charge on an atom may also be
calculated by the following formula:
Polar Covalent Bonds
as a result of shielding by inner
electron shells, the different
atoms of the periodic table
have different affinities for
nearby electrons. The ability of
an element to attract or hold
values was established by
Linus Pauling, and some of
these are given in the table to the right. A larger number on this scale signifies a
greater affinity for electrons. Fluorine has the greatest electronegativity of all the
elements, and the heavier alkali metals such as potassium, rubidium and cesium have
the lowest electronegativities. It should be noted that carbon is about in the middle of
the electronegativity range, and is slightly more electronegative than hydrogen.
When two different atoms are bonded covalently, the shared electrons are attracted to
the more electronegative atom of the bond, resulting in a shift of electron density
toward the more electronegative atom. Such a covalent bond is polar, and will have a
dipole (one end is positive and the other end negative). The degree of polarity and the
magnitude of the bond dipole will be proportional to the difference in
electronegativity of the bonded atoms. Thus a O–H bond is more polar than a C–H
bond, with the hydrogen atom of the former being more positive than the hydrogen
bonded to carbon. Likewise, C–Cl and C–Li bonds are both polar, but the carbon end
is positive in the former and negative in the latter. The dipolar nature of these bonds
is often indicated by a partial charge notation (
δ+/–) or by an arrow pointing to the
negative end of the bond.
Although there is a small electronegativity difference between carbon and
general are considered to be non-polar compounds.
The shift of electron density in a covalent bond toward the more
electronegative atom or group can be observed in several ways. For bonds to
hydrogen, acidity is one criterion. If the bonding electron pair moves away from the
hydrogen nucleus the proton will be more easily transfered to a base (it will be more
acidic). A comparison of the acidities of methane, water and hydrofluoric acid is
instructive. Methane is essentially non-acidic, since the C–H bond is nearly non-
polar. As noted above, the O–H bond of water is polar, and it is at least 25 powers of
ten more acidic than methane. H–F is over 12 powers of ten more acidic than water as
a consequence of the greater electronegativity difference in its atoms.
Electronegativity differences may be transmitted through connecting covalent bonds
by an inductive effect. Replacing one of the hydrogens of water by a more
electronegative atom increases the acidity of the remaining O–H bond. Thus
hydrogen peroxide, HO–O–H, is ten thousand times more acidic than water, and
hypochlorous acid, Cl–O–H is one hundred million times more acidic. This inductive
transfer of polarity tapers off as the number of transmitting bonds increases, and the
presence of more than one highly electronegative atom has a cumulative effect. For
example, trifluoro ethanol, CF
–O–H is about ten thousand times more acidic
than ethanol, CH
A chemical bond represents the net attraction that keeps atoms near each other
oppositely charged particles in atoms--namely electrons and protons.
Because there exists a large number and a diverse arrangement of electrons and
protons in the various atoms of most substances, a precise understanding of all the
complex electrical interactions can be challenging. However, some simplified models
of these interactions allow us to predict many important properties.
First, we divide bonds up into two major categories: primary bonds and
secondary bonds. Primary bonds are the strong bonds between the tightly clustered
atoms that give any pure substance its characteristic properties. Secondary bonds
(also known as interparticle, intermolecular, or Van der Waals attractions) are the
relatively weaker attractions between nearby atoms or molecules that are important in
most liquids (especially liquid mixtures) and some solids.
Primary Bond Types: There are 3 major types of primary bonds--ionic,
two major types of elements,metals and nonmetals.
Remember that metals are elements that have a relatively weak attraction for
their outermost electrons, while nonmetals are elements with a strong attraction for
these electrons. As you move to the left and down on the Periodic Table, elements get
'more metallic'; as you move to the right and up, elements get 'more nonmetallic'.
Metal + Nonmetal
NaCl (table salt)
Nonmetal + Nonmetal
Metal + Metal
Ionic: Metal + Nonmetal
charged cations and negatively charged anions.
Ionic materials are usually brittle solids at room temperature, and they exist as
highly-ordered 3-dimensional arrangements of vast numbers of ions. The exact
proportions of the different types of ions are given by the compound's chemical
formula, and reflect a balance of total positive and negative charges.
Most geological minerals are ionic compounds.
Covalent: Nonmetal + Nonmetal
relatively strong attraction for those electrons.
Covalently bonded materials come in 2 major types--molecular substances and
Molecular substances: Most covalent substances exist as molecules, which are
small to medium-sized groups of atoms connected by covalent bonds. Because a
molecule is attracted to other molecules in the sample by weaker secondary bonds,
most molecular substances exist as gases or liquids at room temperature. The nature
of the attraction between molecules depends a great deal on the way that the atoms in
the molecule are arranged (i.e., on the molecular shape).
Since the strength of secondary bonds tends to go up with molecular size,
larger molecules can exist as solids, but they usually have low melting points and
poor mechanical strength. Waxes, polymers (plastics), and many biological tissues
are all familiar examples of large molecular substances.
Covalent crystals are an extreme example of a large molecule--they are
highly-ordered 3-dimensional arrangements of trillions of trillions of atoms
connected by covalent bonds. These materials have high melting points and good
mechanical strength. The most familiar examples are diamond and graphite, both
composed of purely carbon atoms (just connected in slightly different geometries).
Other examples include silicon, quartz (silicon dioxide), and silicon carbide;
however, because silicon is actually a metalloidrather than a nonmetal, these latter
materials are entering the "grey area" between the different material and bond types.
The valence electrons are easily dislodged from all of the atoms in the sample,
and behave as a "sea of fluid electrons" surrounding positively charged metal cores.
Most metals are solids at room temperature, and they exist as highly-ordered 3-
dimensional arrangements of vast numbers of atoms (i.e., as metallic crystals). The
strength of metallic bonding varies significantly between different metal elements,
and therefore melting points and mechanical strength also vary. Since the valence
electrons are held loosely in all metallic solids, they are good conductors of
electricity and heat, and they can be bent or shaped without breaking.
The Periodic Law and periodic system
The periodic law was developed independently by Dmitri Mendeleev and
followed by Meyer. They both arranged the elements by their mass and proposed that
certain properties periodically reoccur. Meyer formed his periodic law based on the
atomic volume or molar volume, which is the atomic mass divided by the density in
solid form. Mendeleev's table is noteworthy because it exhibits mostly accurate
values for atomic mass and it also contains blank spaces for unknown elements.
In 1804 physicist John Dalton advanced the atomic theory of matter, helping
scientists determine the mass of the known elements. Around the same time, two
chemists Sir Humphry Davy and Michael Faraday developed electrochemistry which
aided in the discovery of new elements. By 1829, chemist Johann Wolfgang
Doberiner observed that certain elements with similar properties occur in group of
three such as; chlorine, bromine, iodine; calcium, strontium, and barium; sulfur,
selenium, tellurium; iron, cobalt, manganese. However, at the time of this discovery
too few elements had been discovered and there was confusion between molecular
weight and atomic weights; therefore, chemists never really understood the
significance of Doberiner's triad.
In 1859 two physicists Robert Willhem Bunsen and Gustav Robert Kirchoff
discovered spectroscopy which allowed for discovery of many new elements. This
gave scientists the tools to reveal the relationships between elements. Thus in 1864,
chemist John A. R Newland arranged the elements in increasing of atomic weights.
Explaining that a given set of properties reoccurs every eight place, he named it the
law of Octaves.
The Periodic Law
In 1869, Dmitri Mendeleev and Lothar Meyer individually came up with their
own periodic law "when the elements are arranged in order of increasing atomic
mass, certain sets of properties recur periodically." Meyer based his laws on the
atomic volume (the atomic mass of an element divided by the density of its solid
form), this property is called Molar volume.
Mendeleev's Periodic Table
physical and chemical properties in the increasing order of the atomic masses, in the
form of a table. Mendeleev had observed that properties of the elements orderly recur
in a cyclic fashion. He found that the elements with similar properties recur at regular
intervals when the elements are arranged in the order of their increasing atomic
masses. He concluded that 'the physical and chemical properties of the elements are
periodic functions of their atomic masses'. This came to be known as the law of
chemical periodicity and stated:
Based on this law all the known elements were arranged in the form of a table
called the 'Periodic Table'. Elements with similar properties recurred at regular
intervals and fell in certain groups or families. The elements in each group were
similar to each other in many properties. The elements with dissimilar properties
from one another were separated. Mendeleev's periodic table contains eight vertical
columns of elements called 'groups' and seven horizontal rows called 'periods', Each
group has two sub-groups A and B. The properties of elements of a sub-group
resemble each other more markedly than the properties of those between the elements
of the two sub-groups.
Mendeleev's periodic table is an arrangement of the elements that group similar
elements together. He left blank spaces for the undiscovered elements (atomic
masses, element: 44, scandium; 68, gallium; 72, germanium; & 100, technetium) so
that certain elements can be grouped together. However, Mendeleev had not
predicted the noble gases, so no spots were left for them.
columns, called groups. Molar volume increases from top to bottom of a group
The alkali metals (Mendeleev's group I) have high molar volumes and they
also have low melting points which decrease in the order:
C) > Na (97.8
C) > K (63.7
C) > Rb (38.9
C) > Cs (28.5
Atomic Number as the Basis for the Periodic Law
Assuming there were errors in atomic masses, Mendeleev placed certain
elements not in order of increasing atomic mass so that they could fit into the proper
groups (similar elements have similar properties) of his periodic table. An example of
this was with argon (atomic mass 39.9), which was put in front of potassium (atomic
mass 39.1). Elements were placed into groups that expressed similar chemical
In 1913 Henry G.J. Moseley did researched the X-Ray spectra of the elements
and suggested that the energies of electron orbitals depend on the nuclear charge and
the nuclear charges of atoms in the target, which is also known as anode, dictate the
frequencies of emitted X-Rays. Moseley was able to tie the X-Ray frequencies to
numbers equal to the nuclear charges, therefore showing the placement of the
elements in Mendeleev's periodic table. The equation he used:
νν: X-Ray frequency
ZZ: Atomic Number
AA and bb: constants
Atomic numbers, not weights, determine the factor of chemical properties. As
mentioned before, argon weights more than potassium (39.9 vs. 39.1, respectively),
yet argon is in front of potassium. Thus, we can see that elements are arranged based
on their atomic number. The periodic law is found to help determine many patterns of
many different properties of elements; melting and boiling points, densities, electrical
conductivity, reactivity, acidic, basic, valance, polarity, and solubility.
The table below shows that elements increase from left to right accordingly to
their atomic number. The vertical columns have similar properties within their group
for example Lithium is similar to sodium, beryllium is similar to magnesium, and so
So, elements in Group 1 (periodic table) have similar chemical properties, they
are called the alkaline earth metals.
Short form periodic table
The short form periodic table is a table where elements are arranged in 7 rows,
periods, with increasing atomic numbers from left to right. There are 18 vertical
columns known as groups. This table is based on Mendeleev's periodic table and the
Long form Periodic Table
In the long form, each period correlates to the building up of electronic shell;
the first two groups (1-2) (s-block) and the last 6 groups (13-18) (p-block) make up
the main-group elements and the groups (3-12) in between the s and p blocks are
called the transition metals. Group 18 elements are called noble gases, and group 17
are called halogens. The f-block elements, called inner transition metals, which are at
the bottom of the periodic table (periods 8 and 9); the 15 elements after barium
(atomic number 56) are called lanthanides and the 14 elements after radium (atomic
number 88) are called actinides.
The Law of Conservation of Mass is that, in a closed system, matter cannot be
The Law of Conservation of Mass is a relation stating that in a chemical
reaction, the mass of the products equals the mass of the reactants. Antoine Lavoisier
stated, "atoms of an object cannot be created or destroyed, but can be moved around
and be changed into different particles".
The principle of conservation of mass was first outlined by Mikhail
Lomonosov (1711–1765) in 1748. He proved it by experiments—though this is
Antoine Lavoisier (1743–1794) had expressed these ideas in
(1728–1799), Henry Cavendish(1731–1810), and Jean Rey (1583–1645).
The conservation of mass was obscure for millennia because of the buoyancy
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
even though they were carried out with other intentions. His
more than from 2 to 4 parts in 100,000.
The difference in the accuracy aimed at and
When elements and compounds react to form new products, mass cannot be
lost or gained.
"The Law of Conservation of Mass" definition states that "mass cannot be
created or destroyed, but changed into different forms".
So, in a chemical change, the total mass of reactants must equal the total
mass of products.
The law of conservation of mass can also be stated "no atoms can be lost or
made in a chemical reaction", which is why the total mass of products must equal the
total mass of reactants you started with.
calculate the quantities of reactants and products involved in a reaction and the
simplest formula of a compound
chemical symbol equation, the total of relative formula masses of the reactants is
equal to the total relative formula masses of the products.
2.3 reactivity series of metals.
and structurally analytical progression of a series of metals, arranged by their
"reactivity" from highest to lowest.
It is used to summarize information about
the reactions of metals with acids and water, double displacement reactions and the
extraction of metals from their ores.