AQA A-Level Physical chemistry
This document contains notes for physical chemistry and
important questions along with formulas for each chapter to help
your AQA A-Level exam preparation.
A sample question paper for the AQA A-Level physical
chemistry examination is at the end of the document.
NOTE:- This study material is for REVISING ONLY,
CHAPTER - 1 ATOMIC STRUCTURE
1. The Atom: An atom is the basic unit of matter, composed of a
nucleus (protons and neutrons) surrounded by electrons in energy
levels or shells.
2. Subatomic Particles:
Protons: Positively charged particles found in the nucleus.
Neutrons: Neutral particles also located in the nucleus.
Electrons: Negatively charged particles orbiting the nucleus in
energy levels.
3. Rutherford's Experiment: Ernest Rutherford's gold foil experiment
showed that atoms have a small, dense, positively charged nucleus
and mostly empty space.
4. Bohr's Model: Niels Bohr proposed that electrons orbit the nucleus
in discrete energy levels. Electrons can jump between levels by
absorbing or emitting photons.
5. Quantum Mechanical Model: The modern model of the atom treats
electrons as wave-like particles described by wave functions. It
provides a probability distribution for finding electrons in certain
regions around the nucleus.
6. Quantum Numbers: Four quantum numbers (n, l, m_l, and m_s)
define the energy level, shape, orientation, and spin of an electron.
7. Electron Configuration: The arrangement of electrons in energy
levels and sublevels in an atom. It follows the Aufbau principle,
Hund's rule, and the Pauli exclusion principle.
,8. Aufbau Principle: Electrons fill lower energy levels before moving
to higher energy levels.
9. Hund's Rule: When filling degenerate orbitals (same energy level),
electrons occupy separate orbitals with the same spin before pairing
up.
10. Pauli Exclusion Principle: No two electrons in an atom can have
the same set of quantum numbers. This leads to electron pairing in
orbitals.
11. Periodic Table: Arranged by increasing atomic number. Periods
represent energy levels, and groups have similar electron
configurations.
12. Electron Shielding and Effective Nuclear Charge: Inner
electrons shield outer electrons from the full attractive force of the
nucleus, resulting in effective nuclear charge.
13. Atomic Radius: Half the distance between the nuclei of two
identical atoms. Decreases across periods and increases down groups
due to increasing effective nuclear charge and energy levels.
14. Ionization Energy: Energy required to remove an electron from a
neutral atom. Increases across periods and decreases down groups due
to increasing effective nuclear charge and electron shielding.
15. Electron Affinity: Energy change when an atom gains an electron.
Trends vary across the periodic table.
16. Isotopes: Atoms of the same element with the same number of
protons but different numbers of neutrons.
17. Quantum Tunneling: The phenomenon where particles can pass
through potential energy barriers that classical mechanics would not
permit.
18. Wave-Particle Duality: Electrons and other particles exhibit both
wave-like and particle-like properties.
19. Heisenberg's Uncertainty Principle: It's impossible to know both
the exact position and momentum of a particle simultaneously.
20. Photoelectric Effect: The emission of electrons from a material
when exposed to light of a certain frequency or higher, supporting the
particle-like behavior of light.
, Theories:
1. Bohr's model of the atom: Niels Bohr proposed that electrons in an
atom occupy discrete energy levels or orbits around the nucleus.
Electrons can move between these levels by absorbing or emitting
photons of specific energies.
2. Quantum Mechanics: Quantum mechanics is the theory that describes
the behavior of particles at the atomic and subatomic levels. It
involves the use of wave functions and operators to predict the
probability distribution of particles.
3. Pauli Exclusion Principle: Proposed by Wolfgang Pauli, this principle
states that no two electrons in an atom can have the same set of four
quantum numbers. This leads to the filling of electron orbitals in a
specific manner following the Aufbau principle and Hund's rule.
4. Aufbau Principle: This principle states that electrons fill atomic
orbitals in order of increasing energy, starting with the lowest energy
orbital and moving to higher energy orbitals as the electron
configuration is built up.
5. Hund's Rule: Hund's rule states that when filling degenerate (same
energy level) orbitals, electrons will occupy separate orbitals with the
same spin before pairing up.
6. Schrödinger's equation: The Schrödinger equation is a fundamental
equation in quantum mechanics that describes the behavior of a
particle as a wave.
7. Planck's Quantum Theory: Max Planck proposed that energy is
quantized and can only be emitted or absorbed in discrete units called
quanta. This theory laid the foundation for quantum mechanics.
8. Photoelectric Effect: The photoelectric effect refers to the emission of
electrons from a material when it is exposed to light of a certain
frequency or higher. It supported the idea that light behaves as both a
wave and a particle.
, Formulas:
1. Energy of a photon: E = hf (E = energy of the photon, h = Planck's
constant, f = frequency of the radiation)
2. Energy of an electron in a hydrogen-like atom: E = - (2.18 x 10^-18 J)
* (Z^2 / n^2) (E = energy of the electron, Z = atomic number of the
element, n = principal quantum number)
3. de Broglie wavelength: λ = h / p (λ = wavelength of the particle, h =
Planck's constant, p = momentum of the particle)
4. Heisenberg's uncertainty principle: Δx * Δp ≥ h / (4π) (Δx =
uncertainty in position, Δp = uncertainty in momentum, h = Planck's
constant)
5. Rydberg equation: 1/λ = R * (1/n1^2 - 1/n2^2) (λ = wavelength of
light, R = Rydberg constant, n1 and n2 = integers representing energy
levels)
This document contains notes for physical chemistry and
important questions along with formulas for each chapter to help
your AQA A-Level exam preparation.
A sample question paper for the AQA A-Level physical
chemistry examination is at the end of the document.
NOTE:- This study material is for REVISING ONLY,
CHAPTER - 1 ATOMIC STRUCTURE
1. The Atom: An atom is the basic unit of matter, composed of a
nucleus (protons and neutrons) surrounded by electrons in energy
levels or shells.
2. Subatomic Particles:
Protons: Positively charged particles found in the nucleus.
Neutrons: Neutral particles also located in the nucleus.
Electrons: Negatively charged particles orbiting the nucleus in
energy levels.
3. Rutherford's Experiment: Ernest Rutherford's gold foil experiment
showed that atoms have a small, dense, positively charged nucleus
and mostly empty space.
4. Bohr's Model: Niels Bohr proposed that electrons orbit the nucleus
in discrete energy levels. Electrons can jump between levels by
absorbing or emitting photons.
5. Quantum Mechanical Model: The modern model of the atom treats
electrons as wave-like particles described by wave functions. It
provides a probability distribution for finding electrons in certain
regions around the nucleus.
6. Quantum Numbers: Four quantum numbers (n, l, m_l, and m_s)
define the energy level, shape, orientation, and spin of an electron.
7. Electron Configuration: The arrangement of electrons in energy
levels and sublevels in an atom. It follows the Aufbau principle,
Hund's rule, and the Pauli exclusion principle.
,8. Aufbau Principle: Electrons fill lower energy levels before moving
to higher energy levels.
9. Hund's Rule: When filling degenerate orbitals (same energy level),
electrons occupy separate orbitals with the same spin before pairing
up.
10. Pauli Exclusion Principle: No two electrons in an atom can have
the same set of quantum numbers. This leads to electron pairing in
orbitals.
11. Periodic Table: Arranged by increasing atomic number. Periods
represent energy levels, and groups have similar electron
configurations.
12. Electron Shielding and Effective Nuclear Charge: Inner
electrons shield outer electrons from the full attractive force of the
nucleus, resulting in effective nuclear charge.
13. Atomic Radius: Half the distance between the nuclei of two
identical atoms. Decreases across periods and increases down groups
due to increasing effective nuclear charge and energy levels.
14. Ionization Energy: Energy required to remove an electron from a
neutral atom. Increases across periods and decreases down groups due
to increasing effective nuclear charge and electron shielding.
15. Electron Affinity: Energy change when an atom gains an electron.
Trends vary across the periodic table.
16. Isotopes: Atoms of the same element with the same number of
protons but different numbers of neutrons.
17. Quantum Tunneling: The phenomenon where particles can pass
through potential energy barriers that classical mechanics would not
permit.
18. Wave-Particle Duality: Electrons and other particles exhibit both
wave-like and particle-like properties.
19. Heisenberg's Uncertainty Principle: It's impossible to know both
the exact position and momentum of a particle simultaneously.
20. Photoelectric Effect: The emission of electrons from a material
when exposed to light of a certain frequency or higher, supporting the
particle-like behavior of light.
, Theories:
1. Bohr's model of the atom: Niels Bohr proposed that electrons in an
atom occupy discrete energy levels or orbits around the nucleus.
Electrons can move between these levels by absorbing or emitting
photons of specific energies.
2. Quantum Mechanics: Quantum mechanics is the theory that describes
the behavior of particles at the atomic and subatomic levels. It
involves the use of wave functions and operators to predict the
probability distribution of particles.
3. Pauli Exclusion Principle: Proposed by Wolfgang Pauli, this principle
states that no two electrons in an atom can have the same set of four
quantum numbers. This leads to the filling of electron orbitals in a
specific manner following the Aufbau principle and Hund's rule.
4. Aufbau Principle: This principle states that electrons fill atomic
orbitals in order of increasing energy, starting with the lowest energy
orbital and moving to higher energy orbitals as the electron
configuration is built up.
5. Hund's Rule: Hund's rule states that when filling degenerate (same
energy level) orbitals, electrons will occupy separate orbitals with the
same spin before pairing up.
6. Schrödinger's equation: The Schrödinger equation is a fundamental
equation in quantum mechanics that describes the behavior of a
particle as a wave.
7. Planck's Quantum Theory: Max Planck proposed that energy is
quantized and can only be emitted or absorbed in discrete units called
quanta. This theory laid the foundation for quantum mechanics.
8. Photoelectric Effect: The photoelectric effect refers to the emission of
electrons from a material when it is exposed to light of a certain
frequency or higher. It supported the idea that light behaves as both a
wave and a particle.
, Formulas:
1. Energy of a photon: E = hf (E = energy of the photon, h = Planck's
constant, f = frequency of the radiation)
2. Energy of an electron in a hydrogen-like atom: E = - (2.18 x 10^-18 J)
* (Z^2 / n^2) (E = energy of the electron, Z = atomic number of the
element, n = principal quantum number)
3. de Broglie wavelength: λ = h / p (λ = wavelength of the particle, h =
Planck's constant, p = momentum of the particle)
4. Heisenberg's uncertainty principle: Δx * Δp ≥ h / (4π) (Δx =
uncertainty in position, Δp = uncertainty in momentum, h = Planck's
constant)
5. Rydberg equation: 1/λ = R * (1/n1^2 - 1/n2^2) (λ = wavelength of
light, R = Rydberg constant, n1 and n2 = integers representing energy
levels)
