STRUCTURES AND PROPERTIES

STRUCTURE AND PROPERTIES

The Structure of an Atom

Ø  All matter is composed of atoms, existing individually or in combination with each other.

Ø  An atom is an extremely small electrically-neutral particle. It is the smallest unit involved in the chemical change of matter.

Ø  The views on the atomic structure which are accepted today have developed from the classical Rutherford-Bohr theory.

Ø  According to this theory, the atom is made of a central positively charged. Nucleus containing positively charged particles called Protons and neutral particles called Neutrons, both having unit mass. The nucleus is surrounded by negatively charged particles called Electrons which carry one unit negative charge and negligible weight.

Ø  The electrons are said to revolve around the nucleus in fixed orbits or energy levels. While the electron moves in such a level, it possesses a definite quantity of energy and it neither emits nor absorbs energy. The electrons are arranged in the orbits so that the maximum number of electrons in the various orbits starting from the one nearest the nucleus is 2, 8, 18, 32, 18,8.

Ø   The outermost orbit of electrons in different atoms (except those of inert gases), is incomplete and the electrons in it are known as the Valence Electrons.

Ø  Together the neutrons and protons give the nucleus its mass, but the protons alone give the nucleus its positive charge.

Atomic Number

            The number of protons or electrons present in an atom is called atomic number. The symbol Z is often used for atomic number (or number of protons).

No. of protons in an atom = No. of electrons

For example - Hydrogen has an atomic number of 1 and carbon has an atomic number of 6.

Atomic Mass Number or Atomic Weight

Ø  The sum of the total number of protons, Z, and the total number of neutrons, N, is called the atomic mass number. The symbol is A.

Ø  Not all atoms of the same element have the same atomic mass number, because, although the Z is the same, the N and thus the A are different.

Ø  The masses of atomic particles are given in atomic mass units (amu).

Ø  A proton has a mass of 1.0 amu and a positive charge (+1). The neutron also has a mass of 1.0 amu but is neutral in charge. The electron has a mass of 0.00055 or 1/1835 amu and a negative charge (-1).

Ø  Atoms of the same element with different atomic mass numbers are called isotopes.

Ø  Isotopes are atoms of the same element having different masses, due to varying numbers of neutrons.

Molecules

            Molecules are groups or clusters of atoms held together by means of chemical bonding. There are two types of molecule.

1.Molecules of an Element

            In certain cases, two single atoms of an element can be attracted to one another by a bond to form a molecule.

Examples -- hydrogen, oxygen, and bromine. The molecular formulas for these are H, O, and Br. Most gaseous elements exist as 2 2 2 molecules of two atoms.

2.Molecules of a Compound

            Two atoms of different elements held together by a bond form a compound. The molecule is the primary particle of a chemical compound.

Examples --hydrogen chloride (HCl), water (H₂O), methane (CH₄) and  ammonia (NH₃).

Molecular Weight

            The weight of a molecule, the molecular weight, is the total mass of the individual atoms. Therefore, it is fairly simple to calculate the mass of any molecule if its formula is known.

For Example – Molecular weight of water (H₂O) = At.wt. of H₂ + At.wt. of O = 2+16 = 18.

ATOMIC ORBITALS

The electron shells are around the nucleus, and the shells are referred to by number. The first, or No. 1, shell is the one nearest the nucleus; the second, or No. 2, shell is next; then the third, or No. 3, shell; and so on in numerical order. The principal cells are also called as the capital letters K, L, M, N, O respectively.

             In general, electrons closer to the nucleus have a lower energy state. Atomic electrons always seek the lowest energy state available. The electron shells represent major energy states of electrons. Each shell contains one or more sub shells called orbitals, each with a slightly different energy. In order of increasing energy, the orbitals are designated by the small letters s, p, d, f, g, h.

            No two shells consist of the same number of orbitals. In general, each higher shell contains a new type of orbital:

The first shell contains an s orbital,

The second shell contains s and p orbitals,

The third shell contains s, p, and d orbitals,

The fourth shell contains s, p, d, and f orbitals, and so on.

Shape of s Orbitals.

            An s orbital has the shape of a sphere (Fig. 5.8.). In this orbital there is an equal probability of finding the electron in any direction away from the nucleus. The difference between an electron in 1s and in 2s orbital is that; the electron in the 2s orbital is further away from the nucleus and has greater energy. In the s sub shell for which there is only one orbital.

Shape of p Orbitals

             A p orbital is dumb-bell shaped (Fig. 5.9). There are two lobes associated with this orbital. They are located on opposite sides of the nucleus and directed along a particular axis. In the p sub shell there are three orbitals of equal energy for electrons to occupy. These three p orbitals are directed along the three coordinate axes

ELECTRON CONFIGURATION

            The distribution of electrons in shells, sub shells and orbitals is governed by the following rules:

The total number of electrons that can be accommodated in a Shell is equal to 2n², where n refers to the principal quantum number of the shell. Thus the first shell can accommodate 2 electrons, the second 8, the third 18, and the fourth 32.

The total number of electrons that can be accommodated in a Sub shell is equal to twice the number of orbitals it contains. Thus an s sub shell can have 2 electrons because it has only one orbital, p sub shell can have 6 electrons due to three orbitals, d can have 10 electrons due to five orbitals, and f can have 14 electrons due to seven orbitals. 

1.      Pauli's Exclusion Principle - The total number of electrons that can be accommodated in an orbital is 2 and these two electrons must have opposite spins. Electrons with opposite spins are given the symbols ­ and ¯.

2.      Electrons occupy the orbital of the lowest energy first and then the filling of the orbitals of higher energy starts. Thus the electrons fill the 1s orbital before occupying 2s orbital. The energy of orbitals of the various sub shells follows the sequence given below:

                        1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s

 In a sub shell, all the available degenerate orbitals (i.e., the orbitals of same energy) are occupied singly first and then pairing of electrons in each orbital occurs. Thus, of the three degenerate orbitals of p sub shell (Px, Py and pz), no one will have two electrons as long as any other is vacant: This is called Hund's Rule, or the Principle of Maximum Multiplicity.

            The distribution of electrons in an atom is always written in terms of electrons which each of the sub shells occupies. In common practice, the number of electrons in the sub shell are indicated as superscripts (Fig. 5.7).

            The complete electron configuration of any element is written by listing the sub shells in the order of increasing energy and the number of electrons occupying each sub shells are indicated by superscripts.

n   Ne à 1s² 2s² 2p⁶                          (10 electrons)

n  F à 1s² 2s² 2p⁵                              (9 electrons)

n  Mg à 1s² 2s² 2p⁶ 3s²                  (12 electrons)

n  Mg²⁺ à 1s² 2s² 2p⁶                      (10 electrons)

            The outermost shell is known as the Valence Shell and the electrons that occupy this shell are called the Valence Electrons.