Inorganic Chemistry Notes

Inorganic Chemistry

• Date: 18/05/2025

• Strength of nuclear charge increases with the number of electrons in the outermost ring of an atom.

• Losing an electron causes an atom to become positively charged (+ve), while gaining an electron causes an atom to become negatively charged (-ve).

• Size of atoms decreases across the periodic table.

  • An atom with a higher atomic number is smaller.

• The number of protons is equal to the number of electrons in a neutral atom.

Course Content

1. Introduction

   • Atomic structure and wave mechanics.

2. Periodic table and periodic properties.

3. Chemical bonding.

4. Chemistry of S & P block.

5. Metal.

6. The principles of qualitative and quantitative analysis.

Atomic Structure and Wave Mechanics

Rutherford Atom Model

• Alpha particles (+ve charge)

• Rutherford designed an experiment using fast-moving alpha particles directed through a thin gold foil.

• Expectation: Alpha particles would be deflected by subatomic particles in the gold atoms.

• Actual Results: The alpha particle scattering experiment gave different results than expected.

Observations

1. Most fast-moving alpha particles passed straight through the gold foil.

2. Some alpha particles were deflected by small angles.

3. One out of every 12,000 particles appeared to rebound (kurudi nyuma).

Conclusions

1. Most of the space inside the atom is empty because most alpha particles passed through without deflection.

2. Very few particles were deflected, indicating the +ve charge of the atom occupies very little space.

3. A very small fraction of alpha particles were deflected, indicating that not all the +ve charge and mass of the gold atom were concentrated in a very small volume within the atom.

Nuclear Model Features

1. There is a +ve charged center in an atom called the nucleus, and nearly all the mass of an atom resides in the nucleus.

2. Electrons revolve around the nucleus in circular paths.

3. The size of the nucleus is very small compared to the size of the atom.

Drawbacks of Rutherford's Model

• The revolution of the electron in a circular orbit is not expected to be stable.

• Any particle in a circular orbit would undergo acceleration.

• During acceleration, charged particles would radiate energy. Therefore, the revolving electron would lose energy and finally spiral into the nucleus.

• If this was the case, the atom would be highly unstable, and hence matter would not exist in the form that we know.

• Atoms are quite stable.

Plank's Quantum Theory

• Plank's Quantum Theory is a natural phenomenon of Quantum mechanics.

• It explains the quantum nature of electromagnetic energy and deals with phenomena such as the photoelectric effect and the nature of radiated emission, which are not explained by the laws of classical mechanics.

• The black body spectrum, which is the radiation emitted by a body that absorbs all the light that falls on it, shows significant divergence at higher frequencies.

• This divergence was at odds with hypothetical results predicted by classical models of physics.

• Planck resolved this by theorizing that a particle in oscillation only absorbs radiation with a minimum amount of energy.

• The value of this minimum energy was discovered by Plank to be proportional to the frequency of the oscillation particle.

• The relation between the energy required to excite the particle to a higher energy level and the frequency of the particle is given by E = h\nu where:

  • E is the minimum energy,

  • h is Plank's Constant,

  • \nu is the frequency.

• This discovery proved significant in the development of Quantum mechanics; therefore, Plank's results show that electromagnetic waves behave as both particles and waves when interacting with matter.

• This dual nature of electromagnetic radiation led to further investigations when someone called De-Broglie, he was successful in obtaining a formula for the wave line of any particle with a mass and momentum.

  • Formula: \lambda = \frac{h}{m \nu}

    • \lambda = wave length

    • h = Plank's constant

    • m = mass

    • \nu = frequency

Electromagnetic Radiation

• This refers to a phenomenon of electromagnetic waves traveling through space.

• They do not need a medium to propagate/move.

• The radiation is sustained by the electric component and magnetic component of the wave, giving rise to each other through space.

• Both components lie in perpendicular lines.

• Electromagnetic radiation consists of:

  • Wave length: This is the distance between 2 consecutive crests or troughs of a wave.

  • Frequency: Number of waves occurring in a given interval of time.

• Waves of higher frequency than visible light come into categories of X-rays, gamma rays, and UV-rays, while waves of lower frequency than visible light are infrared, radio waves, and microwaves.

Black Body Radiation (Planck's Radiation)

• A black body is an object that absorbs all radiation that falls on it.

• An ideal black body absorbs and emits radiations of all frequencies.

• Emission of electromagnetic radiation from a black body depends on temperature.

• The higher the temp of the body, the more is the emission of radiation of all wavelengths.

Postulates of Plank's Quantum Theory

1. Energy is not emitted/radiated continuously. It is emitted in small portions in the form of energy packets called photons or quanta.

2. The radiation, when in the form of light, each particle is known as a photon.

3. The energy of a photon & quantum is directly proportional to frequency of the radiation.

4. The total energy of radiation is represented as a whole number multiple of h\nu e.t.c.

   • When energy of a quantum as E = 1h\nu, 2h\nu, 3h\nu

Photoelectric Effect

• If a beam of light of sufficiently high frequency is allowed to strike a clean metal surface, electrons are ejected from the metal surface.

• The photon having sufficient energy, when it strikes a metal surface, the ejection of electron will take place.

• Otherwise, some of the energy absorbed by an electron from the photon is utilized to free the electron inside the metal, while the rest of the energy is converted to kinetic energy.

• Formula: h\nu = w + K.E

  • w -> work function

  • K.E -> kinetic energy

Bohr Model

• According to classical physics, electrons must radiate electromagnetic waves.

• The loss of energy would result in the electrons spiraling into the nucleus. At this time, Bohr was contemplating how Plank's Quantum nature of radiation could be applied to atomic spectrum and in particular the Rydberg formula for hydrogen.

• Bohr pictured the electron in hydrogen orbiting the central atomic nucleus and postulated a number of assumptions.

Assumptions

1. Atoms have stationary states of definite total energy. In this state, the electrons do not radiate.

2. The emission or absorption of electromagnetic radiation occurs only in transitions from one stationary state to another. The frequency of the electromagnetic radiation is proportional to the energy difference between the two states.

3. Classical physics describes the dynamical equilibrium of the atom in a stationary state but does not describe the transition between stationary states.

4. The mean value of the kinetic energy of the electron to nuclear system is quantized; therefore, quantization of energy is equivalent to the angular momentum of the system being an integer multiple of h.

The Hydrogen Spectrum

• Hydrogen atom contains only 1 electron and shows many lines in the spectrum.

• When hydrogen gas is subjected to electric charge, the hydrogen molecules absorb energy and split into atoms.

• The electrons in the atoms absorb energy and get excited, and while coming back to their ground state, they emit radiation of different frequencies, showing the spectrum.

Series and Spectral Regions

| Series | n | Spectral Region |
| :------ | :- | :------------------- |
| Lyman | 1 | UV |
| Balmer | 2 | Visible |
| Paschen | 3 | Infrared (near) |
| Brackett| 4 | Infrared (middle) |
| Pfund | 5 | Infrared (far) |

• n_1 means first energy level.

• n_2 from the second energy level.

• Atoms become excited - which means electrons move around to lower energy levels.

Limitations of Bohr's Model (hydrogen spectrum)

1. It explains the spectrum of single electron species: hydrogen, lithium ion, and helium ion, but not the spectra of multi-electron species.

2. Cannot explain the fine structure in the atomic spectra.

3. Cannot explain the formation of chemical bonds.

4. It failed to explain the duo nature (pairing) of electrons.

Sommerfeld Atomic Model

• Sommerfeld presents a modified and elaborated version of Bohr's model.

• The fine structure of hydrogen emission lines is evidence of a fine structure of the given atom.

• He suggested that it is not necessary that all the electrons of the same orbit have the same energy; the energies of some electrons of the same orbit may be different from the energies of the other electrons. i.e., the main energy levels are subdivided into subenergy levels.

• Some of these electrons can move on elliptical paths or orbits.

• The nucleus of an atom is at one of the two foci of elliptical orbitals, and the value of angular momentum of an electron moving in a circular orbit is not n\frac{h}{2\pi} where n is a whole number (multiple of energy level).

• n = whole number (main energy level)

• This atomic model successfully explains the finer structure of the micro-structure of the lines obtained in the emission spectrum of hydrogen.

Limitation

• However, this model, just like Bohr's model, fails to explain the emission spectrum of atoms or ions with more than one electron.

De Broglie Relation

• From Bohr's postulates, De Broglie formulated the relation between momentum and wavelength: \lambda = \frac{h}{p}

• If the electron in a hydrogen atom is to be represented by a wave, then the circumference of the orbit must be an integer number of wavelengths: n\lambda = 2\pi r

  • Where r is the radius.

• The equation can be written in terms of electron momentum and radius which is J = rp

  • J = rp where rp = n\hbar

    • quantum momentum number

    • energy levels

Heisenberg Uncertainty Principle (momentum & location)

• It is impossible to measure the position (x) and momentum (p) of a particle with absolute accuracy or precision.

• The more accurately we know one of these values, the less accurately we know the other.

• e.g., when an electron is being viewed, we shine some light on it which imparts some energy to these electrons. This leads to momentum (and location) being altered.

• Alternatively, because electrons move so fast, by the time the incident photons report back to their position, the electrons would have already moved, therefore affecting the calculations about their position.

• This principle is fundamental to understanding the structure of an atom.

• Therefore, it helped overcome the deficiencies or limitations of models of atoms like Rutherford, Bohr's, etc.

• It basically states that the product of the errors in the measurement of momentum and position of electrons is equal to a constant.

  • (\Delta P) (\Delta x) > \frac{h}{4\pi}

    • If the error of measurement of momentum is \Delta P and \Delta x = Position

    • h = Plank's constant = 6.626 X 10^{-34} J/s (Joules per second)

Quantum Numbers

• n, l, ml, and ms are quantum numbers that explain the wave function of a wave equation

  • \Psi-wave function.

• To explain the size of the orbit, energy of electron, the shape of the orbital, the orientation of orbital, and the spin of the electron, quantum numbers are used.

• n is known as the principle quantum number.

• l is known as the azimuthal number.

• m_l is known as the magnetic number.

• m_s is known as the spin number.

Principal Quantum Number (n)

• It is denoted by the number n and given the values of 1, 2, 3, 4…

• It represents the orbits of the main shells around the nucleus and their size and energy, and

• As the value of n increases, the size and energy of orbit increases.

• If the orbit number of orbitals are present, they are given by n^2.

  • N = orbital numbers with that energy levels.

• The number of electrons is given by 2n^2 because each orbit has electrons.

• Main shells - size and energy.

Azimuthal Quantum Number (l)

• Proposed by Sommerfeld denoted & it represents the subshells (sub- levels) present.

• They can have the values of 0 to n-1. and it describes the shape of atomic orbitals.

Magnetic Quantum Number (m_l)

• This describes the number of orbitals present in a given sub-shell and can have values of -l through zero to +l value!

• It describes the orientation of orbitals in space.

The Spin Quantum Number (m_s)

• This describes the spin of electrons in each orbital.

• m_s value of a clockwise electron is positive +\frac{1}{2} and that of an anticlockwise electron is negative -\frac{1}{2}.

The Shapes of s, p, d orbitals

• An electron orbital is a mathematical function used in describing the wave nature mechanism of either an electron or many pairs of electrons in an atom.

• Every energy level is shown by a number and the electron orbital is shown by letters - e.g., 1s.

• the election orbital is shown by a whole number & letter and the is there to shows the level of the energy and the s is there to refer Shape of the orbital.

The s orbital

• s orbitals are found to be spherical symmetric like a hollow bowl with a nucleus at the centre.

• The orbitals grow bigger as the energy levels increase and the electrons get located far away from the nucleus.

• The size of the nucleus ground in the order of 1s < 2s < 3s < 4s.

• Not every electrons are found in s orbitals.

The p orbital

• The only orbital located in the 1st energy level is 1s, however p orbitals are found in the 2nd energy levels - 2s and 2p.

• The difference between a 2s orbital and a 2p orbital is that a s orbital points in no particular direction while a p orbital points in a certain direction.

• p-Orbitals: 2px , 2py , 2pz and 3px , 3py , 3pz.

d Orbitals

• In the 3rd energy level s,d orbitals are present.

• 3s, 3p and d are also present.

• The total number of orbitals in the 3rd energy level are 9 (3s, 3p, 5d orbitals).

• d orbitals are named as

  • 3d_{xy}

  • 3d_{xz}

  • 3d_{yz}

  • 3d_{(x^2 - y^2)}

  • 3d_{(z^2)}

• Every orbital contains 4 lobes and each lobe points in the middle of 2 axis and not along the axis.

The Pauli Exclusion Principle

• This principle describes the arrangement of electrons in an atom & along with Aufbau's and Hund's rule.

• It states that no 2 electrons can have the identical set of quantum numbers simultaneously.

• Every electron should have a unique quantum state; therefore, no 2 electrons with up spin can be arranged together and vice versa in a single quantum state.

• The helium atom has 2 electrons, and they occupy the outermost shell with opposite spins.The 2 electrons are in 1s sub-shell, where n=1, l=0, m_l = 0.

• However, both electrons will have differences in m_l : +\frac{1}{2}, -\frac{1}{2}

• Hydrogen atom has 1 electron in 1s sub-shell with 1 up spin electron.

Aufbau Principle

• Describes the arrangement of elections in an atom in order of their increasing energy level.

• This means that the atomic orbital with the lowest energy are filled first by electrons before occupying the upper atomic orbitals.

• The order in which the atomic orbitals are filled by elections is like 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p etc.

• An arrow is used to depict the electrons in Aufbau principle. Each shell carries a max electrons as follows:

  • n=1, s, max of 2

  • n=2, p, max of 6

  • n=3, d, max of 10

  • n=4, f, max of 14

Hund's Rule

• It states that for a given electron configuration, the term with Maximum Multiplicity falls lowest in energy.

• According to this rule, electron pairing in p,d & f orbitals cannot occur until each orbital of a sub-shell contains is singly occupied.

• Nitrogen Atom:

  • 1s^2 2s^2 2p^3

• The electron enters an empty orbital before pairing.

• Electrons don't share orbitals until all the orbitals are singly filled.

Periodic Properties

Modern Periodic Law

• It states that the Physical & Chemical properties of elements are Periodic functions of their atomic numbers.

• If elements are arranged in their order of increasing number, the elements with Similar Properties are repeated after certain regular Intervals.

Long Form of Periodic Table

• This is a perfect matching of the electronic configuration of elements on one hand and Physical & Chemical Properties on the other hand.

Principles

1. An atom losses electrons from or gains electrons on the outer most shell of an atom during a chemical reaction

2. Sharing of electrons by an atom with another atom is largely through the outer most shell. Therefore electrons in the outer most shell of an atom largely determine the chemical Properties of an elements.

Features of a Periodic Table

1. There are 18 vertical columns (groups) in the periodic table and 7 horizontal rows called periods, giving a total number of 114 elements.

2. Of all the known elements, 90 are naturally occurring, and others are made through nuclear transformation or synthesized artificially.

Periodic Table Groups

• Group 1 elements are called Alkali Metals (except hydrogen)

• Group 2 are called alkaline earth metals

• Group 3 to 12 are called Transition metals

• Group 16 are called Chalcogens

• Group 17 are called Halogens

• Group 18 are called Noble gases

• Elements with atomic numbers 58 to 71 are called Lanthanoids (inner transition series)

• Elements with atomic numbers from 90-103 are called Actinoids (Inner transition series)

• All elements except transitions and Inner transition elements are collectively known as main group elements

Periodic Properties

1. Atomic Radius

2. Ionization Energy

3. Electronic affinity

4. Electromagnet negativity

• The repetition of properties of elements at regular Intervals in the periodic table is called periodicity

Atomic Radius

• It is the distance between the center of a nucleus to the outermost shell of an atom.

• Atomic radius increases down the group because of the increase in the number of energy levels.

• Atomic radius decreases across the Period from left to right because an Increase in the number of protons (nuclear charge).