Basic Concepts of Atomic Structure

Atoms:

• Atoms are the basis of all matter. 

• They are small and consist of even tinier particles. 

• Neutrons, Protons, and Electrons are the basic particles making up the atom. 

  • They join together with other atoms and create matter. 

  • It takes many atoms to create anything.

• It is the smallest constituent unit of matter that possesses the properties of the chemical element. 

• Atoms do not exist independently, instead, they form ions and molecules which further combine in large numbers to form matter that we see, feel and touch.

Molecules:

• Molecules consist of one or more atoms bound together by chemical bonds. 

• Atoms may be depicted by circle shapes, each of which has a nucleus at the center (containing protons and neutrons), surrounded by one or more concentric circles representing the ‘shells’ or ‘levels’ in which the electrons surrounding the nucleus of the atom are located and markings indicating the electron at each level. 

• A molecule is the smallest thing a substance can be divided into while remaining the same substance. It is made up of two or more atoms that are bound together by chemical bonding.

Nucleus:

• The central part of an atom, containing protons and neutrons.

Subatomic Particles:

Protons:

• Protons are positively charged subatomic particles. 

• The charge of a proton is 1e, which corresponds to approximately 1.602 × 10^-19

• The mass of a proton is approximately 1.672 × 10^-24

• Protons are over 1800 times heavier than electrons.

• The total number of protons in the atoms of an element is always equal to the atomic number of the element.

Neutrons:

• The mass of a neutron is almost the same as that of a proton, i.e., 1.674×10^-24

• Neutrons are electrically neutral particles and carry no charge.

• Different isotopes of an element have the same number of protons but vary in the number of neutrons present in their respective nuclei.

Electrons:

• The charge of an electron is -1e, which approximates to -1.602 × 10^-19

• The mass of an electron is approximately 9.1 × 10-31.

• Due to the relatively negligible mass of electrons, they are ignored when calculating the mass of an atom.

Atomic Number and Mass Number

Atomic Number (Z):

• The number of protons in an atom's nucleus. It defines the element.

Mass Number (A):

• The total number of protons and neutrons in an atom's nucleus.

  •  It is the mass of an atom in a chemical element. 

  • It can be expressed in atomic mass units (denoted by u or amu). 

    • 1 amu is equal to exactly one-twelfth of the mass of 1 atom of C-12 and the relative atomic masses of elements is determined with respect to these C-12 atoms.

Isotopes

• Isotopes are the atoms in which the number of neutrons differs while the number of protons is the same. 

  • From the above definition of mass number and the atomic number, we can conclude that isotopes are those elements having the same atomic number and different mass numbers.

Atomic Structure of Some Elements

Hydrogen:

• The most abundant isotope of hydrogen on the planet Earth is protium. The atomic number and the mass number of this isotope are 1 and 1, respectively.

• Structure of Hydrogen Atom:

  • This implies that it contains one proton, one electron and no neutrons (Total number of neutrons = Mass number – Atomic number)

Carbon:

• Carbon has two stable isotopes – 12C and 13C. Of these isotopes, 12C has an abundance of 98.9%. It contains 6 protons, 6 electrons and 6 neutrons.

• Structure of Carbon Atom:

  • The electrons are distributed into two shells, and the outermost shell (valence shell) has four electrons. 

  • The tetravalency of carbon enables it to form a variety of chemical bonds with various elements.

Oxygen:

• There exist three stable isotopes of oxygen – 18O, 17O and 16O. However, oxygen-16 is the most abundant isotope.

• Structure of Oxygen Atom:         

  • Since the atomic number of this isotope is 8 and the mass number is 16, it consists of 8 protons and 8 neutrons. 6 out of the 8 electrons in an oxygen atom lie in the valence shell.

Electronic Configuration

Electron Shells:

• Electrons are arranged in energy levels or shells around the nucleus.

  • The electrons have to be filled in sublevels called s, p, d, and f by the following rule.

Aufbau principle:

• The filling of electrons should take place by the ascending order of energy of orbitals (regions of space with a high probability of finding an electron)

• Lower energy orbital should be filled first, and higher energy levels.

• The energy of orbital α (p + l) relates to orbitals that have the same (n + l) value, E α n. So, the S subshell has 1 orbital (2 electrons), p has 3 orbitals etc.

• Ascending order of energy 1s, 2s, 2p, 3s, 3p, 4s, 3d, . . .

Pauli’s exclusion principle:

• The maximum number of electrons in an orbital is two. If there are two electrons in an orbital, they must have opposite spin.

Hund’s rule:

• Electrons fill orbits of the same energy so as to give the maximum number of electrons with the same spin, before doubling up with opposite spin.

Stark effect:

• The phenomenon of deflection of electrons in the presence of an electric field.

Zeeman effect:

• The phenomenon of deflection of electrons in the presence of a magnetic field.

Bohr’s Atomic Theory

• Neils Bohr put forth his model of the atom in the year 1915. 

  • This is the most widely used atomic model to describe the atomic structure of an element which is based on Planck’s theory of quantisation.

Postulates:

• The electrons inside atoms are placed in discrete orbits called “stationary orbits”.

• The energy levels of these shells can be represented via quantum numbers.

• Electrons can jump to higher levels by absorbing energy and move to lower energy levels by losing or emitting their energy.

  • As long as an electron stays in its own stationery, there will be no absorption or emission of energy.

• Electrons revolve around the nucleus in these stationary orbits only.

• The energy of the stationary orbits is quantised.

Limitations of Bohr’s Atomic Theory:

• Bohr’s atomic structure works only for single electron species such as H, He+, Li2+, Be3+, ….

• When the emission spectrum of hydrogen was observed under a more accurate spectrometer, each line spectrum was seen to be a combination of a number of smaller discrete lines.

• Both Stark and Zeeman’s effects could not be explained using Bohr’s theory.

Heisenberg’s uncertainty principle

• Heisenberg stated that no two conjugate physical quantities could be measured simultaneously with 100% accuracy. 

• There will always be some error or uncertainty in the measurement.

• Drawback: 

  • Position and momentum are two such conjugate quantities that were measured accurately by Bohr (theoretically).

Quantum Mechanical Model of the Atom

• Orbitals: Regions of space where there is a high probability of finding an electron.

• Types of Orbitals: s (spherical), p (dumbbell-shaped), d, and f.

• Quantum Numbers: Describe the properties of atomic orbitals and the electrons in them.

  • Principal Quantum Number (n): Indicates the energy level.

  • Angular Momentum Quantum Number (l): Indicates the shape of the orbital.

  • Magnetic Quantum Number (m): Indicates the orientation of the orbital.

  • Spin Quantum Number (s): Indicates the spin of the electron.

Periodic Table and Electron Configuration

Periods:

• Horizontal rows in the periodic table. 

  • The period number indicates the highest energy level of electrons in an atom.

Groups:

• Vertical columns in the periodic table.

  • Elements in the same group have similar chemical properties and the same number of valence electrons.

Trends in the Periodic Table

Atomic Radius:

• Decreases across a period and increases down a group.

Ionization Energy:

• The energy required to remove an electron from an atom.

  • Increases across a period and decreases down a group.

Electronegativity:

• A measure of an atom's ability to attract electrons.

  • Increases across a period and decreases down a group.

Key Takeaways

• Understanding atomic structure is fundamental to chemistry and explains the behaviour of elements in the periodic table.

• The arrangement of electrons in atoms determines an element's chemical properties and reactivity.

• Trends in the periodic table provide insights into the properties of elements and their compounds.