chem

MATTER


Solids State
Arrangement: Particles are packed in a fixed, regular pattern and are in direct contact with their neighbors.
Movement: Particles lack the freedom to move from place to place; they can only vibrate in their fixed positions.
Properties:
- Incompressible: They cannot be squeezed into a smaller volume because there is very little space between particles.
- Fixed Shape: They do not flow and maintain a constant shape and volume.
Liquid State
Arrangement: The particles still touch their neighbour, but they slide over each other.
Movement: The particles can slide over each other.
Properties: liquids cannot be compressed but can flow and they have no fixed shape.

Gas state
Arrangement: The particles are spread out.

Movement: The particles move freely in all directions.

Properties: Gases can be compressed and can flow.


Mass - The amount of matter in an object. Mass of an object refers to a fied quantity that is independent of the object's location

Weight - It is the pull of gravity of an object and depends on the object's location

Volume - It is a scalar quantity, the amount of space that an object occupies.

Density - It is a substance's mass in a certain volume. It tells you how heavy something is for its size.

• DENSITY = MASS/VOLUME
• MASS = DENSITY X VOLUME
• VOLUME = MASS/DENSITY





ATOMS

​Ancient Philosophical Concepts

• ​Greek philosopher Democritus proposed that matter is made of indivisible, solid particles called "atomos".

• ​Indian philosopher Acharya Kanada independently introduced a similar concept called "Anu".

• ​These early ideas were philosophical speculations with no experimental evidence.

​Dalton’s Atomic Theory & Chemical Laws

• ​John Dalton introduced the first scientific atomic model, viewing atoms as solid, indivisible spheres.

• ​His model provided a physical basis for fundamental chemical laws, including the Law of Conservation of Mass and the Law of Definite Proportions.

​Discovery of the Electron (Plum Pudding Model)

• ​J.J. Thomson used cathode ray tube experiments to discover the first subatomic particle: the negatively charged electron.

• ​He proposed the Plum Pudding Model, depicting electrons embedded inside a positively charged fluid to form a neutral atom.

• ​Rutherford’s Gold Foil Experiment & Nuclear Model

• ​Ernest Rutherford bombarded thin gold foil with positively charged alpha particles.

• ​Because most particles passed through while a few bounced back, he concluded that the atom is mostly empty space with a dense, positively charged center called the nucleus.

• ​His planetary model had electrons orbiting the nucleus, but it couldn't explain atomic stability under Maxwell's laws of electromagnetism (which predicted orbiting charges should continuously lose energy and collapse).

​Bohr’s Quantized Model

• ​Niels Bohr resolved the stability problem by proposing that electrons move in fixed, quantized energy levels (shells) where they do not lose energy.

• ​He formulated mathematical equations for orbital radii, electron velocities, and energy levels.

• ​Bohr’s model successfully explained the spectral emission lines of hydrogen, as electrons emit or absorb light photons when jumping between discrete energy levels.


• ​Rutherford identified the proton, and later James Chadwick discovered the neutrally charged neutron.
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​Wave-Particle Duality & Heisenberg's Uncertainty Principle

• ​Louis de Broglie demonstrated that matter, including subatomic particles, possesses wave-particle duality.

• ​Warner Heisenberg introduced the Uncertainty Principle, proving it is impossible to simultaneously determine both the exact position and velocity of an electron.

​The Quantum Mechanical Model

• ​Erwin Schrödinger and Heisenberg developed quantum mechanics, replacing fixed circular orbits with probabilistic electron clouds or orbitals.

• ​Instead of defining an exact path, the quantum model defines a region (orbital) where an electron has roughly a 90% probability of being located upon measurement.




ATOMIC THEORY & QUANTUM MECHANICS

Bohr's Atomic Model

What Bohr Got Right

• Electrons occupy fixed energy levels
• Electrons absorb/emit energy
• Explains hydrogen spectrum

But...

• Does not explain atoms with many electrons and electron behavior accurately


Max Planck (1900)

• He is the father of Quantum theory

• He discovered that energy is not released continuously, it comes in tiny packets called Quanta.

• Equation:    E = hf

• E = energy of one quantum
• h = Planck's constant (6.626 x 10^-34 J-s)
• f = Frequency of the radiation (hz)


Albert Einstein (1905)

• Explained that light behaves like particles called photons.

• Photons can knock electrons out of metals, it became known as the Photoelectric Effect.

• He won the nobel prize for explaining this phenomenon, not for relativity.

Louis De Broglie (1924)

• He concluded that electrons exhibit wave-particle duality.

• Electrons are particles that also behave as waves.


Erwin Schrodinger (1926)

• Instead of predicting the next path, scientists should calculate the probability of finding an electron.

• Electrons do not travel in fixed orbits, they occupy regions where they are most likely to be found.


Werner Heisenberg (1927)

• Developed the Uncertainty Principle

• It is impossible to know both the exact position and the exact momentum of an electron at the same time.

• The more precisely we measure momentum, the less precisely we know the position.



ORBIT VS ORBITAL

Orbit (BOHR MODEL)

• Fixed circular path
• Exact location
• Defined and Predictable.

Orbital (QUANTUM MECHANICAL MODEL)

• Region of high probability
• Likely location
• Cannot know exact position and path

Electron Cloud Model

• A region around the nucleus where an electron is most likely to be found.

• It represents probability.



QUANTUM NUMBERS

Principal Quantum Number

• Energy Level. It tells you the main electron shell or distance from the nucleus.

• Think of it like the specific floor of an House / Floor

• n = Number of the shell

• As n increases, the electron is farther from the nucleus and energy increases.


Angular Quantum Number

• It determines the geometric shape and type of the orbital is spherical, is dumbbell-shaped, is clover-shaped, etc.

• As L increases, the shape becomes larger and more complex.

• There is only 1 s orbital, but multiple orientations exist for p, d, and f.

• s = 1 p = 3  d = 5  f = 7

• Multiply each value by 2 to find the amount of electrons that an electron can accomodate.
Magnetic Quantum Number

• Orientation. It determines how the orbital is tilted or pointing in 3D space (along the X, Y, or Z axes).

• It is a map that tells us the direction.

• s (L=0) = 1 orientation (0) p (L=1) = 3 orientations ( -1, 0, +1)
• d (L=2) = 5 orientations (-2, -1, 0, +1, +2)
• f (L=3) = 7 orientations ( -3, -2, -1, 0, +1, +2, +3)


Spin Quantum Number

• Spin. It represents the rotation direction of the electron itself, which can only be spin-up or spin-down

• Only one orbital can hold only two electrons and they must have opposite spins.

• spin up = (+1/2) spin down = (-1/2)











ELECTRON CONFIGURATION

• It shows the distribution of electrons in the orbitals in atom.

• It can be represented using either orbital diagram or spdf notation.


Ground state electron configuration

• Arrangement of electrons in an atom's orbitals that corresponds to the lowest possible energy state.

Excited state electron configuration

• When one or more electrons have absorbed energy, the electron moves to a higher energy orbitals.

Orbital diagram

• Uses a box or horizontal line to represent an orbital, while an arrow is used to indicate an electron.

• Upward arrow = +1/2 OR Clockwise electron spin

• Downward arrow = -1/2 OR Counterclockwise electron spin

• Electron spins are also referred to as spin-up or spin-down.

• Orbitals are filled in order of increasing energy, wherein a maximum of two electrons is allowed per orbital.

Degenerate Orbitals

• Orbitals that have the same energy in the absence of a magnetic/electric field. It applies to p, d, and f orbitals.

SPDF Notation

• The electron configuration of an atom written by specifying occupied subshells along with the number of electrons.

• It is done by combining the principal quantum number (n) and the subshell letter (s, p, d, f), followed by a superscript that shows the number of electrons in that subshell (#).

















RULES IN WRITING ELECTRON CONFIGURATION


Aufbau/"Building Up" Principle

• Electrons occupy orbitals in order of increasing energy

• It implies that electrons first occupy the lowest energy levels before moving to higher ones



Pauli Exclusion Principle

• Two electrons cannot have the same set of quantum numbers.

Hund's Rule

• Electrons fill a subshell singly before forming any pairs and each electron in a single occupied orbital has the same spin.

Octet Rule

• Atoms gain, lose, and share electrons to fill their valence shell with 8 electrons.


PARAMAGNETIC VS DIAMAGNETIC

Paramagnetic

• UNPAIRED electrons
• Attracted to a magnetic field

Diamagnetic

• PAIRED electrons
• Repelled by a magnetic field.


• ALL materials are diamagnetic, but in atoms with unpaired electrons, paramagnetism overcomes diamagnetism.