Chemistry

Shapes of molecules, bond angles

VSEPR Theory stands for Valence Shell Electron Pair Repulsion Theory.
It is used to predict the shape of molecules based on electron pair repulsion.
The theory states that electron pairs around a central atom arrange themselves to minimize repulsion.
Both bonding pairs and lone pairs of electrons influence molecular geometry.
Lone pairs occupy more space than bonding pairs due to greater repulsion.
The molecu

Valence Bond Theory (VBT) explains the formation of chemical bonds through the overlap of atomic orbitals.
In VBT, atoms combine by overlapping their valence orbitals to form bonds.
The overlapping orbitals contain unpaired electrons, which pair up to form a covalent bond.
Bond strength depends on the extent of overlap; greater overlap results in stronger bonds.
VBT distinguishes between sigma (σ) and pi (π) bonds based

Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.
Measured on the Pauling scale, where fluorine has the highest value (4.0).
Electronegativity increases across a period from left to right due to increasing nuclear charge and a smaller atomic radius.
Electronegativity decreases down a group as atomic size increases, reducing the pull on shared electrons.
Non-metals like oxygen, nitrogen, and fluor

The atomic radius is the distance from the nucleus to the outermost shell of an atom.
Atomic radius decreases across a period due to increased nuclear charge, which pulls electrons closer to the nucleus.
Atomic radius increases down a group because of the addition of new electron shells.
Cations (positively charged ions) have smaller radii than their parent atoms due to the loss of electrons and reduced electron-electron re

Proposed by Dmitri Mendeleev in 1869.
Based on the periodic recurrence of properties of elements when arranged in order of increasing atomic mass.
Contained 63 known elements at the time.
Mendeleev organized elements into rows and columns to form the periodic table.
He left gaps in the table for undiscovered elements and predicted their properties accurately, such as gallium (eka-aluminum) and g

Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons.
As a result, isotopes have the same atomic number but different mass numbers.
Examples of isotopes include:
- Hydrogen: Protium (¹H), Deuterium (²H), Tritium (³H).
- Carbon: Carbon-12 (¹²C), Carbon-13 (¹³C), Carbon-14 (¹⁴C).
Isotopes exhibit identical chemical properties but different

Quantum numbers describe the state and position of an electron in an atom.
There are four quantum numbers:
- Principal Quantum Number (n): Indicates the main energy level or shell of the electron.
- Azimuthal Quantum Number (l): Defines the subshell and shape of the orbital (0 for s, 1 for p, 2 for d, 3 for f).
- Magnetic Quantum Number (m_l): Specifies the orientation of the orbital in space.
- S

Proposed by Niels Bohr in 1913.
Based on Rutherford’s model and Planck's quantum theory.
Electrons revolve around the nucleus in fixed circular paths called orbits or energy levels.
These energy levels are represented as n = 1, 2, 3, ... (principal quantum numbers).
Each orbit has a fixed amount of energy, and electrons do not radiate energy while in a stable orbit.
Electrons can transition betwee

Proposed by John Dalton in 1803.
It was the first scientific theory to describe the nature of matter in terms of atoms.
Matter is made up of small, indivisible particles called atoms.
Atoms of a given element are identical in size, mass, and other properties.
Atoms of different elements differ in size, mass, and other properties.
Atoms cannot be created, divided, or destroyed during chemical reactions.
Atoms combine in simple whole-number ratios to form