The electronic configuration of an atom is notation that describe the orbitals in which the electrons occupy and total number of electron in each orbitals.
Niels Bohr proposed that electrons only exist in specific energy states when an electron absorb or emits a specific amount of energy it instantaneously moves from one orbital to another.
Energy Levels of Electrons
As you may remember from chemistry, an atom consists of electrons orbiting
around a nucleus. However, the electrons cannot choose any orbit they wish. They
are restricted to orbits with only certain energies. Electrons can jump from one
energy level to another, but they can never have orbits with energies other than
the allowed energy levels.
The Energy difference between two particular energy level is called the quantum of energy.
- Ground State: When all the electron of an atom are in their lowest possible energy levels.
- Excited State: When on or more of an atom's electrons are in energy levels other than the lowest available level.
Electron Configuration
The letters S,P,D,F are different types of orbital.
Example
For example, let's say you have an atom of lithium. Lithium's
atomic number is 3. So a neutral atom of lithium has 3 protons and 3 electrons.
We would need space for 3 electrons. Write:
There are 2 electrons in the 1s sublevel and 1 electron in the 2s sub-level. 2
+ 1 = 3. That's it!
Now let's try fluorine, which has an atomic number of 9. A neutral
atom of fluorine has 9 protons and 9 electrons. We need enough space for 9
electrons. The 1s orbital can hold 2 electrons and the 2s orbital can hold 2
more electrons. The five remaining electrons must go into the next orbital, the
2p orbital. The 2p orbital can hold up to 6, but we only have 5. So the
following would be the correct electron configuration for a neutral atom of
fluorine. 2 + 2 + 5 = 9.
Writing Electron Configuration for Ion
For a negative ion: Add electrons (equal to the charge) to the last unfilled sub-shell starting where the neutral atom left off.
For a positive ion: 1) Start with the neutral configuration, remove electrons from the outermost shell first. 2) If there 's are electrons on both the S and P orbitals of the outermost shell, the electrons in p orbitals are removed first.
Valance Electrons
The number of valence electrons is just how many electrons an atom has in its
outer shell. It's easy to figure out if you've got a periodic table. All the elements in each column have the same number of
electrons in their outer shells. All the elements in the first column all have a
single valence electron (H, Li, Na, K, etc.). The second column elements all have 2 valence electrons (Be, Mg,
Ca, Sr, etc.). Skipping over the gap, go to the Group 3
elements, which all have 3 valence electrons (B, Al, Ga, etc.). The elements in the next column (C, Si, Ge,
etc.) all have 4 valence electrons. The elements in the next
column (N, P, As, etc.) all have, yes, you guessed it, 5 valence electrons. O, S, Se, and the others in this column have 6 valence
electrons. The halogens in the next-to-last column (F, Cl,
Br, etc.) have 7 valence electrons. The
noble gases in the right-most column (Ne, Ar, Kr, etc.) all have 8 electrons in
their out except for He, which only has 2 electrons. If an atom is an ion, you must include the charge also: For a positive ion, for each charge subtract one electron, *for
instance, Na+ has 1-1 = 0, BUT it
has 8 valence electrons because it has the same
electron configuration as Ne. Just as K+ has the same configuration as Ar. Therefore, the
alkali metal ions with a single positive charge will have 8 valence electrons. For a negative ion, add one electron for
each charge, for instance, O2- has
6+2 = 8 valence
electrons
decrease across a period with increase in number of valence electrons as well as a decrease in atomic radius
increase down the group with increase in number of shells and atomic radius
Atomic Radius
the distance from the atomic nucleus to the outermost stable electron orbital in an atom
decrease as one progresses across a period from left to right because the effective nuclear charge increases, thereby attracting the orbiting electrons and lessening the radius
usually increases while going down a group due to the addition of a new energy level (shell)
Ionization energy
minimum amount of energy required to remove one electron from each atom in a mole of atoms in the gaseous state
increase while one progresses across a period because the greater number of protons (higher nuclear charge) attract the orbiting electrons more strongly
As one progresses down a group on the periodic table, the ionization energy will likely decrease since the valence electrons are farther away from the nucleus and experience a weaker attraction to the nucleus
Electronegativity
a measure of the ability of an atom or molecule to attract pairs of electrons in the context of a chemical bond
as one moves from left to right across a period in the periodic table, the electronegativity increases due to the stronger attraction that the atoms obtain as the nuclear charge increases
Moving down a group, the electronegativity decreases due to the longer distance between the nucleus and the valence electron shell, thereby decreasing the attraction, making the atom have less of an attraction for electrons or protons
Reactivity
how likely or vigorously an atom is to react with other substances
Metals
Period - reactivity decreases as you go
from left to right across a period
Group
- reactivity increases as you go down a group
Non-metals
Period - reactivity increases as you go
from the left to the right across a period
Group - reactivity decreases as you go down the group
Ion Charge
Melting Point and Boiling Point
Metals - the melting point for metals generally decreases as you go down a
group.
Non-metals - the melting point for non-metals generally increases as
you go down a group.
The known elements (>60) were arranged in order of atomic weights and
observed similarities between the first and ninth elements, the second and tenth
elements etc.
He proposed the 'Law of Octaves'.
Also he assigned Hydrogen an arbitary mass of 1.
Dmitri Mendeleev (Russian) in 1869
First produced a table based on atomic weights but arranged 'periodically' with
elements with similar properties under each other
Gaps were left for elements
that were unknown at that time and their properties predicted (the elements were
gallium, scandium and germanium)
The order of elements was re-arranged if their
properties dictated it, eg, tellerium is heavier than iodine but comes before it
in the Periodic Table
Mendeleev's Periodic Table was important because it enabled the
properties of elements to be predicted by means of the 'periodic law':
properties of the elements vary periodically with their atomic weights.
Mendeleev's 1871 periodic table
Modern Periodic Table
Our modern day periodic table is expanded beyond Mendeleev's initial 63 elements. Most of the current periodic tables include 108 or 109 elements.
Groups
The modern periodic table of the elements contains 18 groups, or vertical columns.
Elements in a group have similar chemical and physical properties because they have the same number of outer electrons.
Elements in a group are like members of a family--each is different, but all are related by common characteristics.
To avoid confusion, the Roman numerals and letters designating groups will eventually be replaced by the numerals from one to eighteen.
Periods
Each of the table's horizontal rows is called a period.
Along a period, a gradual change in chemical properties occurs
from one element to another.
Changes in the properties occur because the number of protons and electrons increases from left to right across a period or row.
The increase in number of electrons is important because the outer electrons determine the element's chemical properties.
Metals
They are usually shiny, very dense, and only melt at high temperatures.
Their shape can be easily changed into thin wires or sheets without breaking.
Metals will corrode, gradually wearing away, like rusting iron.
Heat and electricity travel easily through metals, which is why it is not wise to stand next to a flagpole during a thunderstorm!
Nonmetals
On the right side of the periodic table, are very different from metals.
Their surface is dull and they don’t conduct heat and electricity.
As compared to metals, they have low density and will melt at low temperatures.
The shape of nonmetals cannot be changed easily because they are brittle and will break.
Metalloids
Elements that have properties of both metals and nonmetals.
They can be shiny or dull and their shape is easily changed.
Electricity and heat can travel through metalloids but not as easily as they travel through metals.
Atom is very small, about 1000millionths of a millimeter in diameter. Although the atom is very small, but it is located in the center of the atom's nucleus and tiny electronic center of movement of these electrons around the atomic nucleus like the planets of the solar system around the sun.
From the British chemist and physicist J.John, Dalton after the creation of atomic theory, for a long time people thought that atoms like a despicably small glass medicine ball, which is no longerthere is no pattern. In 1869, German scientists Hittorf found that cathoderay, a large number of scientists studied the cathode ray tube, which lasted more that twenty years. Ultimatelym by Joseph, John Thomson discovered the existence of electrons. Under normal circunstances, the atom is uncharged, since ran out of the small 1700 times more than its quality from the atomic negatively charged electrons, which shows te atomic internal structure of the original house also shows that there is something positively charged, they and electronics carried by the negatively charged atoms neutral.
A unit of matter, the smallest unit of an element, consisting of a dense, central, positively charged nucleus surrounded by a system of electrons, equal in number to the number of nuclear protons, centimeter and characteristically remaining undivided in chemical reactions except for limited removal, transfer, or exchange of certain electrons.
John Dalton
From his experiments and observations, he suggested that atoms were like tiny, hard balls. An element is a substance made from only one type of atom. An element cannot be broken down into any simpler substances. Element had its own atoms that differed from others in mass. Dalton believed that atoms were the fundamental building blocks of nature and could not be split. In chemical reactions, the atoms would rearrange themselves and combine with other atoms in new ways.
J.J. Thomson
At the end of the nineteenth century, a scientist called J.J. Thomson discovered the electron. This is a tiny negatively charged particle that is much, much smaller than any atom. Thomson proposed a different model for the atom. He said that the tiny negatively charged electrons must be embedded in a cloud of positive charge (after all, atoms themselves carry no overall charge, so the charges must balance out). Thomson imagined the electrons as the bits of plum in a plum pudding.
Ernest Rutherford
In 1911, Ernest Rutherford interpreted these results and suggested a new model for the atom. The positive charge must be concentrated in a tiny volume at the centre of the atom; otherwise the heavy alpha particles fired at the foil could never be repelled back towards their source. On this model, the electrons orbited around the dense nucleus.
Niels Bohr
Bohr suggested that the electrons must be orbiting the nucleus in the centre of an atom, containing protons and neutrons. The energy must be given out when 'excited' electrons fall from a high energy level to a low one.
Bohr and Arnold Sommerfeld
Bohr and a German physicist, Arnold Sommerfeld expanded the original Bohr model to explain these variations. According to the Bohr-Sommerfeld model, not only do electrons travel in certain orbits but the orbits have different shapes and the orbits could tilt in the presence of a magnetic field. Orbits can appear circular or elliptical, and they can even swing back and forth through the nucleus in a straight line.
Wolfgang Pauli
Pauli gave a rule governing the behavior of electrons within the atom that agreed with experiment. If an electron has a certain set of quantum numbers, then no other electron in that atom can have the same set of quantum numbers. Physicists call this "Pauli's exclusion principle." It provides an important principle to this day and has even outlived the Bohr-Sommerfeld model that Pauli designed it for.
Louis de Broglie
In 1924 a Frenchman named Louis de Broglie thought about particles of matter. He thought that if light can exist as both particles and waves, why couldn't atom particles also behave like waves? In a few equations derived from Einstein's famous equation, (E=mc2) he showed what matter waves would behave like if they existed at all. (Experiments later proved him correct.)
Erwin Schrödinger
In 1926 the Austrian physicist, Erwin Schrödinger had an interesting idea: His theory worked kind of like harmonic theory for a violin string except that the vibrations traveled in circles. The world of the atom, indeed, began to appear very strange. It proved difficult to form an accurate picture of an atom because nothing in our world really compares with it.
Heisenberg
In 1927 Heisenberg formulated an idea, which agreed with tests, that no experiment can measure the position and momentum of a quantum particle simultaneously. Scientists call this the "Heisenberg uncertainty principle." This implies that as one measures the certainty of the position of a particle, the uncertainty in the momentum gets correspondingly larger. Or, with an accurate momentum measurement, the knowledge about the particle's position gets correspondingly less.
James Chadwick
Not until 1932 did the English physicist James Chadwick finally discover the neutron. He found it to measure slightly heavier than the proton with a mass of 1840 electrons and with no charge (neutral). The proton-neutron together, received the name, "nucleon."
Paul Dirac
In 1928, Paul Dirac produced equations which predicted an unthinkable thing at the time- a positive charged electron. He did not accept his own theory at the time. In 1932 in experiments with cosmic rays, Carl Anderson discovered the anti-electron, which proved Dirac's equations. Physicists call it the positron.
