Part 1: Atoms, Elements, Molecules, Compounds, Chemical Reactions, and Modern Atomic Theory
1. Atomic Structure and Basic Chemical Concepts
Definition of an Atom
An atom is the smallest unit of an element that retains the chemical properties of that element. It serves as the basic structural unit of matter and participates directly in chemical reactions.
Atoms are composed of smaller particles known as subatomic particles.
Components of an Atom
Protons
Positively charged particles
Located inside the nucleus
Determine the atomic number
Neutrons
Electrically neutral particles
Found inside the nucleus
Contribute to atomic mass
Electrons
Negatively charged particles
Occupy regions surrounding the nucleus
Responsible for chemical bonding
The nucleus contains almost all the mass of the atom.
Atomic Number
Definition:
The atomic number is the number of protons present in the nucleus of an atom.
Symbol:
Z
Example:
Carbon contains:
Protons = 6
Electrons = 6
Atomic number = 6
Mass Number
Mass number refers to the total number of protons and neutrons in an atom.
Formula:
Mass Number = Protons + Neutrons
Example:
Sodium:
Protons = 11
Neutrons = 12
Mass number:
= 11 + 12
= 23
Important note:
Electrons contribute negligibly to atomic mass.
Elements
An element is a pure substance consisting of only one type of atom and cannot be broken into simpler substances by ordinary chemical methods.
Examples:
Hydrogen (H)
Carbon (C)
Oxygen (O)
Gold (Au)
Silver (Ag)
Molecules
A molecule consists of two or more atoms chemically bonded together.
Molecules may contain:
Similar atoms
Examples:
O2
H2
N2
Different atoms
Examples:
H2O
CO2
NH3
Compounds
A compound is a pure substance composed of two or more different elements chemically combined in fixed proportions.
Examples:
Water (H2O)
Sodium chloride (NaCl)
Carbon dioxide (CO2)
Difference between elements and compounds:
| Element | Compounds |
| One type of atom | Different atoms |
| Cannot be chemically broken down | Can be chemically decomposed |
| Represented by symbols | Represented by formulas |
| Found on the periodic table | Not found on the periodic table |
| Example: Oxygen (O), Iron (Fe) | Example: Water (H2O), Sailt (NaCl) |
Key Takeaways
Atom = smallest unit of an element
Atomic number = number of protons
Mass number = protons + neutrons
Elements contain one type of atom
Molecules contain bonded atoms
Compounds contain different atoms chemically combined
2. Chemical Reactions
A chemical reaction is a process involving the rearrangement of atoms through breaking old chemical bonds and forming new ones.
General form:
Reactants → Products
Example:
2H2 + O2 → 2H2O
Hydrogen reacts with oxygen to produce water.
Types of Chemical Reactions
A. Synthesis Reaction
Two or more substances combine to form one product.
General form:
A + B → AB
Examples:
2H2 + O2 → 2H2O
2Mg + O2 → 2MgO
B. Decomposition Reaction
A compound breaks into simpler substances.
General form:
AB → A + B
Examples:
CaCO3 → CaO + CO2
2H2O → 2H2 + O2
C. Single Replacement Reaction
One element replaces another element in a compound.
General form:
A + BC → AC + B
Examples:
Zn + 2HCl → ZnCl2 + H2
Fe + CuSO4 → FeSO4 + Cu
D. Double Replacement Reaction
Two compounds exchange ions.
General form:
AB + CD → AD + CB
Examples:
AgNO3 + NaCl → AgCl + NaNO3
CaCl2 + 2KOH → Ca(OH)2 + 2KCl
E. Addition Reaction
Multiple molecules combine to form one larger molecule.
Examples:
C2H4 + H2 → C2H6
C2H4 + Cl2 → C2H4Cl2
F. Combustion Reaction
A substance reacts with oxygen and releases energy.
Examples:
CH4 + 2O2 → CO2 + 2H2O + Heat
C3H8 + 5O2 → 3CO2 + 4H2O + Heat
Energy Changes in Reactions
Exothermic Reaction
Reaction that releases energy.
Examples:
- Combustion
- Respiration
Endothermic Reaction
Reaction that absorbs energy.
Examples:
- Photosynthesis
- Thermal decomposition
Importance of Chemical Reactions
- Essential for biological processes
- Used in industries
- Important in manufacturing drugs and fertilizers
- Involved in everyday activities
Examples:
Cooking
Digestion
Cleaning
Exam Points
- Atomic number depends on proton number.
- Electrons determine chemical behaviour.
- Synthesis forms larger compounds.
- Decomposition breaks compounds apart.
- Combustion reactions release heat and light.
- Exothermic reactions release energy.
- Endothermic reactions absorb energy.
Part 2: Modern Electronic Theory of Atoms
1. Early Atomic Models
Understanding the modern structure of the atom developed gradually through contributions from several scientists. Early models attempted to explain how particles are arranged within atoms.
A. Thomson Atomic Model
Based on cathode ray experiments, J. J. Thomson proposed that atoms consist of negatively charged electrons embedded in a positively charged sphere.
This model became known as the "plum pudding model" because electrons were imagined as particles distributed inside a positively charged substance like fruits inside a pudding.
Main Features
• Atoms contain electrons.
• Positive charge is spread uniformly throughout the atom.
• Total positive and negative charges balance each other.
• Atoms are electrically neutral.
Limitations
• Could not explain the existence of the nucleus.
• Could not explain atomic spectra.
• Could not explain electron arrangement.
Exam Points
• J. J. Thomson discovered electrons.
• Cathode ray experiments led to the Thomson model.
• The model is known as the plum pudding model.
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B. Rutherford Atomic Model
Rutherford proposed that atoms contain a very small, dense, positively charged center called the nucleus, with electrons surrounding it.
This conclusion came from the alpha-particle scattering experiment conducted by Geiger and Marsden.
Observations from Experiment
• Most alpha particles passed through the foil.
• Some particles were slightly deflected.
• Very few were reflected backward.
Conclusions
• Most of the atom consists of empty space.
• Positive charges are concentrated in the nucleus.
• Electrons move around the nucleus.
Features of Rutherford Model
• Dense central nucleus.
• Electrons surround the nucleus.
• Most atomic volume is empty space.
Limitations
• Could not explain why orbiting electrons do not lose energy.
• Could not explain atomic spectra.
Exam High-Yield Points
• Rutherford discovered the nucleus.
• Most of the atom is empty space.
• Positive charges are concentrated in a small region.
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2. Atomic Spectrum and Bohr Theory
Atomic Spectrum
Atomic spectroscopy is the study of radiation emitted or absorbed by atoms.
When atoms absorb energy, electrons become excited and move to higher energy levels. When they return to lower energy levels, energy is released as electromagnetic radiation.
Hydrogen Spectrum
Hydrogen atoms produce characteristic spectral lines when excited.
Important series include:
1. Lyman Series
- Ultraviolet region
2. Balmer Series
- Visible region
3. Paschen Series
- Infrared region
4. Brackett Series
- Infrared region
5. Pfund Series
- Infrared region
6. Humphreys Series
- Infrared region
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Relationship Between Wavelength, Frequency and Energy
Frequency and wavelength are related by:
c=λν
Where:
• c = speed of light
• λ = wavelength
• ν = frequency
Energy of a photon is given by:
E = hν
Where:
• E = energy
• h = Planck constant
• ν = frequency
Example
As wavelength decreases:
• Frequency increases
• Energy increases
Short wavelength therefore corresponds to high energy radiation.
Exam High-Yield Points
• Short wavelength means high energy.
• Long wavelength means low energy.
• Energy and frequency are directly proportional.
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3. Bohr Theory of the Hydrogen Atom
Bohr Postulates
Bohr proposed several assumptions regarding electron movement around the nucleus.
Postulate 1
Electrons move in specific allowed circular paths called stationary states.
Postulate 2
Electrons do not emit energy while moving in allowed orbits.
Postulate 3
Energy is emitted or absorbed only when electrons move between energy levels.
Energy change is expressed as:
E = hν
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Ground State and Excited State
Ground State
The lowest energy level of an atom.
For hydrogen:
n = 1
Excited State
Any state above the ground state.
Examples:
• n = 2
• n = 3
• n = 4
When energy is absorbed:
Electron moves upward.
When energy is released:
Electron returns downward.
Example
Hydrogen absorbs energy:
n = 1 → n = 2
Hydrogen releases energy:
n = 2 → n = 1
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Successes of Bohr Model
• Explained hydrogen spectrum.
• Introduced discrete energy levels.
• Explained ground and excited states.
Limitations of Bohr Model
• Applicable mainly to hydrogen and one-electron species.
• Failed for multi-electron atoms.
• Did not explain electron wave behavior.
Exam High-Yield Points
• Hydrogen has discrete energy levels.
• Ground state has minimum energy.
• Excited state has higher energy.
• Bohr model works mainly for one-electron systems.
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4. Wave-Particle Duality
Wave-particle duality states that matter and radiation possess both wave and particle properties.
Louis de Broglie proposed that electrons can behave like waves.
This idea explained why electrons occupy only certain energy levels.
Evidence
Electron diffraction experiments confirmed the wave nature of electrons.
Examples
Particle behavior:
• Photoelectric effect
Wave behavior:
• Diffraction
• Interference
Exam High-Yield Points
• de Broglie proposed wave nature of electrons.
• Electrons possess dual characteristics.
• Electrons behave as particles and waves.
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5. Heisenberg Uncertainty Principle
Werner Heisenberg stated that the exact position and momentum of an electron cannot be known simultaneously.
The principle is represented by:
Δx × mΔv ≥ h/4π
Where:
• Δx = uncertainty in position
• Δv = uncertainty in velocity
• m = mass
• h = Planck constant
Implications
If position is determined accurately:
• Velocity becomes uncertain.
If velocity is determined accurately:
• Position becomes uncertain.
Example
A moving car can be tracked accurately because of its large mass.
Electrons are extremely small; therefore exact measurements become impossible.
Exam High-Yield Points
• Exact position and momentum cannot be measured simultaneously.
• Greater accuracy in position gives lower accuracy in velocity.
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6. Quantum Mechanical Model
The quantum mechanical model explains electron behavior using probability rather than exact paths.
Electrons are not considered to travel in fixed circular orbits.
Instead, they occupy regions called orbitals.
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Atomic Orbital
An orbital is a region around the nucleus where the probability of finding an electron is highest.
Important points:
• Orbitals are probability regions.
• Electrons do not move in fixed paths.
• Orbital shape depends on energy level.
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Electron Probability Density
The probability of locating an electron is expressed using:
ψ²
Higher ψ² values indicate a greater chance of finding an electron.
Interpretation
• High dot concentration = high electron probability
• Low dot concentration = low electron probability
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Key Takeaways
• Thomson discovered electrons.
• Rutherford discovered the nucleus.
• Hydrogen spectrum consists of discrete lines.
• Bohr introduced stationary energy levels.
• Ground state represents minimum energy.
• Electrons possess wave-particle duality.
• Exact position and momentum cannot be simultaneously determined.
• Quantum mechanics describes electrons using probability.
• Orbitals are regions of highest electron probability.
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Exam Revision Summary
1. Thomson → Plum pudding model
2. Rutherford → Nuclear model
3. Bohr → Fixed energy levels
4. de Broglie → Wave nature of electrons
5. Heisenberg → Uncertainty principle
6. Quantum mechanics → Probability model of electrons