Inorganic Chemistry

Nuclear Charge: Increases with the number of electrons in the outermost ring of an atom.
Ions: Atoms losing electrons become positively charged (+ve), while those gaining electrons become negatively charged (-ve).
Atomic Size: Decreases across the periodic table; atoms with higher atomic numbers are generally smaller.
Neutral Atoms: Number of protons equals the number of electrons.

Experiment: Rutherford used fast-moving alpha particles (+ve charge) through a thin gold foil.
Observations:
- Most alpha particles passed straight through.
- Some were deflected at small angles.
- About 1 in 12,000 appeared to rebound.
Conclusions:
1. Most of the space in an atom is empty.
2. The positive charge occupies very little space.
3. Most of the +ve charge and mass are concentrated in a small volume.
Nuclear Model Features:
1. A +ve charged center called the nucleus contains nearly all the mass.
2. Electrons revolve around the nucleus in circular paths.
3. The nucleus is very small compared to the atom.
Drawbacks of Rutherford's Model:
- Revolution in a circular orbit is unstable.
- Charged particles radiate energy during acceleration.
- Electrons would lose energy and fall into the nucleus, making the atom unstable.

Explains the quantum nature of electromagnetic energy.
*Deals with phenomena like the photoelectric effect.
Resolved issues with classical physics predictions of black body spectrum.
Key Idea: Particles in oscillation absorb radiation with a minimum amount of energy.
$E = h\nu$ where:
- $E$ = minimum energy
- $h$ = Plank's Constant
- $\nu$ = frequency
Electromagnetic waves behave as both particles and waves when interacting with matter.

Electromagnetic waves traveling through space without needing a medium.
Sustained by electric and magnetic components perpendicular to each other.
Wave Characteristics:
- Wavelength: Distance between two consecutive crests or troughs.
- Frequency: Number of waves in a given time interval.
Types: Includes X-rays, gamma rays, UV rays (higher frequency), infrared, radio waves, and microwaves (lower frequency).

A black body absorbs all radiation falling on it and emits radiation of all frequencies.
Emission depends on temperature; higher temperature means more radiation emission.

Energy is emitted in small packets called photons or quanta.
Light form is known as photons.
Energy of a photon is directly proportional to frequency: $E = h\nu$ , $2h\nu$ , etc.

Electrons are ejected from a metal surface when struck by light of sufficiently high frequency.
Photon energy is used to free electrons and convert the rest to kinetic energy.
Formula: $h\nu = w + KE$ , where $w$ is the work function and $KE$ is kinetic energy.

Postulates:
1. Atoms have stationary states with definite energy; electrons don't radiate in these states.
2. Radiation occurs during transitions between stationary states with frequency proportional to energy difference.
3. Classical physics describes equilibrium in stationary states but not transitions.
4. Kinetic energy of the electron-nucleus system is quantized.

Hydrogen atoms absorb energy and split; electrons get excited and emit radiation when returning to ground state, creating a spectrum.
Spectral Series:
- Lyman (UV)
- Balmer (Visible)
- Paschen, Brackett, Pfund (Infrared)

Only explains spectra of single-electron species (hydrogen, lithium ion, helium ion).
Cannot explain fine structure in atomic spectra.
Cannot explain chemical bond formation.
Failed to explain electron pairing.

Modified Bohr's model to explain the fine structure of hydrogen emission lines.
Suggested sub-energy levels within main energy levels. Some electrons move on elliptical paths.

It is impossible to measure both position ( $x$ ) and momentum ( $p$ ) of a particle with absolute accuracy.
The more accurately one value is known, the less accurately the other is known.
$(\Delta p)(\Delta x) \geq \frac{h}{4\pi}$
- $\Delta p$ = momentum error
- $\Delta x$ = position error

$n, l, m<em>l, m</em>s$ explain the wave function of a wave equation.
Used to describe orbit size, electron energy, orbital shape and orientation, and electron spin.
Types:
- $n$ (Principal Quantum Number): Represents main shells and their size and energy.
- $l$ (Azimuthal Quantum Number): Represents subshells (0 to $n-1$ ) and orbital shape.
- $m_l$ (Magnetic Quantum Number): Describes the number of orbitals and their orientation (- $l$ to + $l$ ).
- $m_s$ (Spin Quantum Number): Describes electron spin (+1/2 or -1/2).

Electron orbital: Mathematical function that describes warelike Mechanism. *S orbitals: spherical syntetric like a holley bowl with a nucleus at the centre. *The orbitals grow bigger as the energy levels increase. *P orbitals: points in a certain direction at 2nd energy level.
- d Orbitals: at the 3rd energy level.

No two electrons can have the identical set of quantum numbers in simultaneously . Every electron should have the single quantums state.
Helium atoms has 2 electrons and they occupy the 1s shell

Atomic orbital with the lowest energy are field first by elections before occupying the upper atomic orbituals.

According to this rule election pairing in p/d/f orbitals Cannot
occur. Until each orbitual of sub-shell contains electrons is silagly.

Modern Periodic Law: Properties of elements are periodic functions of their atomic numbers. *Features of a Periodic Table:
- Group 1 elements are called. Alicalian metals except. hydrogen
- Group 2 bave called alkaline earth metals
- Group 3 to 15 are called Transition metals
- Group 16 are called Chalcogens
- Group 17 are called hallogen stops
- Group 18 are called noble gases
- Elements with atomic numbers 58 to 71 are called Lanthanoids
- Elements with atomic numbers 90-103 are called Actinoids

Atomic Radius: Distance from the nucleus to the outermost shell. Increases down the group (more energy levels) and decreases across the period (increased nuclear charge).