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Topic 1: Properties and structure of atoms
Chemistry | GRADE 11 | QCAA Board
“Atoms can be modelled as a nucleus surrounded by electrons in distinct energy levels.”
The notes are structured for Grade 11 / QCAA-style conceptual understanding.
All matter is made of atoms, which are extremely small particles that retain the chemical properties of an element.
Although atoms are very small \( ≈ (10^{-10} \) m), scientists have developed models to explain their internal structure.
Modern atomic theory describes an atom as:
"A dense nucleus surrounded by electrons arranged in discrete energy levels"
This model explains many chemical and physical properties of elements.
2. Subatomic Particles
Atoms consist of three main subatomic particles.
| Particle | Symbol | Charge | Relative Mass | Location |
|---|---|---|---|---|
| Proton | \( (p^+) \) | +1 | 1 | Nucleus |
| Neutron | \( (n^0) \) | 0 | 1 | Nucleus |
| Electron | \( (e^-) \) | −1 | \(\frac{1}{1836} \) | Electron cloud |
Key Points
• Protons determine the identity of an element
• Neutrons contribute to atomic mass
• Electrons are responsible for chemical reactions
3. The Atomic Nucleus
The nucleus is the central core of the atom.
Characteristics of the Nucleus
• Extremely small but very dense
• Contains protons and neutrons
• Contains almost all the mass of the atom
Typical size comparison:
| Structure | Approximate Size |
|---|---|
| Atom radius | \( (10^{-10} \) ) m |
| Nucleus radius | \( (10^{-15}) \) m |
This means the nucleus is about 100,000 times smaller than the atom.
4. Atomic Number and Mass Number
Atomic Number (Z)
The atomic number is the number of protons in the nucleus.
\( Z = \text{Number of protons} \)
Example:
Hydrogen Z = 1, Carbon Z = 6, Oxygen Z = 8
Since atoms are electrically neutral:
\( \text{Number of electrons} = \text{Number of protons} \)
Mass Number (A)
The mass number is the total number of protons and neutrons.
\( A = p + n \)
Example: Carbon-12
\( A = 12 \)
\( Z = 6 \)
Neutrons:
\( n = A - Z \)
\( n = 12 - 6 = 6 \)
5. Isotopes
Isotopes are atoms of the same element that have:
• the same number of protons
• different numbers of neutrons
Example:
| Isotope | Protons | Neutrons | Mass Number |
|---|---|---|---|
| Carbon-12 | 6 | 6 | 12 |
| Carbon-13 | 6 | 7 | 13 |
| Carbon-14 | 6 | 8 | 14 |
Important Properties
• Chemical properties are almost identical
• Physical properties may differ (mass related)
Example uses:
• Carbon-14 → Radiocarbon dating
6. Historical Development of Atomic Models
Scientists developed several models of the atom.
Dalton’s Atomic Model (1803)
John Dalton proposed that:
• Matter is made of indivisible atoms
• Atoms of the same element are identical
• Atoms combine in simple ratios
Limitations:
• Did not explain subatomic particles
Thomson’s Plum Pudding Model (1897)
J.J. Thomson discovered the electron.
Model:
• Atom is a positively charged sphere
• Electrons embedded inside like plums in pudding
Limitations:
• Could not explain scattering experiments
Rutherford Nuclear Model (1911)
Rutherford performed the gold foil experiment.
Observation:
• Most alpha particles passed through
• Few were strongly deflected
Conclusion:
• Atom is mostly empty space
• Positive charge concentrated in a small nucleus
Bohr Model (1913)
Niels Bohr proposed that:
• Electrons move in fixed circular orbits
• Each orbit has a specific energy
These orbits are called energy levels.
7. Electron Energy Levels
Electrons occupy discrete energy levels (shells) around the nucleus.
These levels are often labelled:
n = 1, 2, 3, 4 ...
where n is the principal quantum number.
Energy Level Characteristics
| Energy Level | Shell Name |
|---|---|
| n = 1 | K |
| n = 2 | L |
| n = 3 | M |
| n = 4 | N |
Key Ideas
• Electrons cannot exist between energy levels
• They jump between levels by absorbing or releasing energy
This explains atomic emission spectra.
8. Maximum Electrons in Energy Levels
The maximum number of electrons in a shell is given by:
\( \text{Maximum electrons} = 2n^2 \)
Where (n) is the energy level.
| Energy Level | Maximum Electrons |
|---|---|
| 1 | 2 |
| 2 | 8 |
| 3 | 18 |
| 4 | 32 |
Example:
Oxygen (8 electrons)
Electron arrangement: 2, 6
9. Electron Configuration
Electrons fill the lowest energy levels first.
Example:
Hydrogen (1 electron)
\( 1s^1 \)
Carbon (6 electrons)
\( 1s^2 , 2s^2 , 2p^2 \)
Oxygen (8 electrons)
\( 1s^2 , 2s^2 , 2p^4 \)
10. Electron Excitation and Emission
When atoms absorb energy:
• electrons move to higher energy levels
This is called excitation.
\( e^{-}_{low} + energy \rightarrow e^{-}_{high} \)
When electrons fall back to lower levels:
• energy is released as light (photons)
\( e^{-}_{high} \rightarrow e^{-}_{low} + photon \)
This produces atomic emission spectra.
11. Atomic Emission Spectra
Each element produces unique spectral lines.
Example:
• Hydrogen emission spectrum
• Sodium yellow flame test
This occurs because electrons transition between specific energy levels.
12. Modern Quantum Mechanical Model
The modern model improves Bohr's model.
Electrons do not move in fixed circular orbits.
Instead, they exist in orbitals, which are probability regions where electrons are likely to be found.
Orbitals include:
• s orbitals
• p orbitals
• d orbitals
• f orbitals
Example shapes:
| Orbital | Shape |
|---|---|
| s | spherical |
| p | dumbbell |
13. Importance of Atomic Structure
Understanding atomic structure helps explain:
Chemical Bonding
How atoms combine to form molecules.
Periodic Table Trends
Patterns such as:
• atomic radius
• ionisation energy
• electronegativity
Spectroscopy
Identification of elements through spectral lines.
Chemical Reactivity
Valence electrons determine how atoms react.
14. Summary
Key concepts:
• Atoms consist of protons, neutrons, and electrons
• The nucleus contains most of the mass
• Electrons occupy distinct energy levels
• Energy levels explain atomic spectra
• The modern model describes electrons in orbitals
Atomic structure forms the foundation of chemistry, explaining the behaviour of elements and their interactions.
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