CXC Chemistry: Atomic Structure

1. Introduction to Atomic Structure

The concept of the atom has evolved significantly throughout history. The word "atom" comes from the Greek word "atomos," meaning indivisible. However, we now know that atoms consist of smaller subatomic particles.

Historical Development of Atomic Theory

Dalton's Atomic Theory (1808):
- All matter is made up of tiny, indivisible particles called atoms
- Atoms of the same element are identical
- Atoms of different elements have different properties
- Atoms combine in simple whole-number ratios to form compounds
- Atoms cannot be created, destroyed, or divided in chemical reactions
Thomson's Plum Pudding Model (1897):
- After discovering electrons, Thomson proposed that atoms were spheres of positive charge with embedded electrons
- Often described as plums (electrons) in a pudding (positive charge)
Rutherford's Nuclear Model (1911):
- Based on the famous gold foil experiment
- Discovered that atoms have a small, dense, positively charged nucleus
- Most of the atom is empty space
- Electrons orbit the nucleus like planets around the sun
Bohr's Atomic Model (1913):
- Electrons move in specific circular orbits (energy levels) around the nucleus
- Electrons can jump between energy levels by absorbing or emitting energy
- Introduced the concept of quantized energy levels
Modern Quantum Mechanical Model:
- Electrons do not follow specific paths but exist in probability clouds called orbitals
- These orbitals represent regions where electrons are most likely to be found
- Based on quantum mechanics and wave equations

2. Subatomic Particles

Atoms are composed of three fundamental particles: protons, neutrons, and electrons.

Particle	Charge	Relative Mass	Location
Proton	+1	1	Nucleus
Neutron	0	1	Nucleus
Electron	-1	1/1836	Orbitals

Properties of Protons

Positively charged particles
Determine the element's identity (atomic number)
Located in the nucleus
Have a relative mass of 1 atomic mass unit (amu)

Properties of Neutrons

Neutral particles (no charge)
Located in the nucleus alongside protons
Help stabilize the nucleus
Have a relative mass of 1 amu
Different numbers of neutrons create isotopes of an element

Properties of Electrons

Negatively charged particles
Orbit the nucleus in energy levels
Have negligible mass (about 1/1836 of a proton)
Responsible for chemical bonding and reactions
Arranged in energy levels (shells)

3. Atomic Number, Mass Number, and Isotopes

Atomic Number (Z)

Equals the number of protons in an atom
Defines which element an atom is
All atoms of the same element have the same atomic number
In a neutral atom, the number of electrons equals the number of protons

Mass Number (A)

The sum of protons and neutrons in an atom
A = number of protons + number of neutrons
Different isotopes of the same element have different mass numbers

Isotopes

Atoms of the same element with different numbers of neutrons
Have the same atomic number but different mass numbers
Have the same chemical properties but slightly different physical properties
Written as: ^A_ZX where X is the chemical symbol

Example:
Carbon has three naturally occurring isotopes:

Carbon-12: 6 protons, 6 neutrons (98.93% abundance)
Carbon-13: 6 protons, 7 neutrons (1.07% abundance)
Carbon-14: 6 protons, 8 neutrons (trace amounts, radioactive)

Calculating Number of Subatomic Particles

For an atom ^A_ZX:

Number of protons = Z
Number of neutrons = A - Z
Number of electrons (in a neutral atom) = Z
Number of electrons (in an ion with charge n) = Z - n

Example:
For ³⁵₁₇Cl^-:

Protons = 17
Neutrons = 35 - 17 = 18
Electrons = 17 + 1 = 18 (gained one electron to form a negative ion)

4. Electron Configuration and Orbital Arrangement

Energy Levels (Shells)

Electrons are arranged in different energy levels or shells around the nucleus
Each shell can hold a maximum number of electrons
The shells are labeled as K, L, M, N, etc. or by numbers 1, 2, 3, 4, etc.

Shell	Maximum Number of Electrons
K (1)	2
L (2)	8
M (3)	18
N (4)	32

Subshells and Orbitals

Each energy level contains subshells labeled as s, p, d, and f
Each subshell contains a certain number of orbitals
Each orbital can hold a maximum of 2 electrons

Subshell	Number of Orbitals	Maximum Electrons
s	1	2
p	3	6
d	5	10
f	7	14

Electron Configuration Principles

Electron configuration shows how electrons are distributed among the available orbitals, subshells, and shells. It follows these key principles:

Aufbau Principle:
- Electrons fill orbitals from lowest energy to highest energy
- The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, etc.
Pauli Exclusion Principle:
- No two electrons in an atom can have the same four quantum numbers
- Each orbital can hold a maximum of two electrons with opposite spins
Hund's Rule:
- Electrons occupy orbitals of equal energy singly before pairing up
- These single electrons must have parallel spins

Examples of Electron Configurations

Hydrogen (Z=1): 1s¹
Helium (Z=2): 1s²
Lithium (Z=3): 1s² 2s¹
Carbon (Z=6): 1s² 2s² 2p²
Sodium (Z=11): 1s² 2s² 2p⁶ 3s¹
Chlorine (Z=17): 1s² 2s² 2p⁶ 3s² 3p⁵

Shorthand Electron Configuration

For larger atoms, we can use noble gas notation as shorthand:
Sodium (Z=11): [Ne] 3s¹
Chlorine (Z=17): [Ne] 3s² 3p⁵
Iron (Z=26): [Ar] 4s² 3d⁶

5. Periodic Table and Electron Configuration

Relationship Between Electron Configuration and Periodic Table

The periodic table is arranged by increasing atomic number, and the elements' positions reflect their electron configurations.

Groups (vertical columns): Elements in the same group have the same number of electrons in their outermost shell (valence electrons)
Periods (horizontal rows): Elements in the same period have the same number of electron shells

Blocks in the Periodic Table

The periodic table is divided into four blocks based on the subshell being filled:

s-block: Groups 1 and 2 (the alkali metals and alkaline earth metals)
p-block: Groups 13-18 (includes non-metals, metalloids, and some metals)
d-block: Groups 3-12 (transition metals)
f-block: Lanthanides and actinides (inner transition metals)

Valence Electrons

Valence electrons are electrons in the outermost energy level of an atom
They determine the chemical properties and behavior of an element
Elements in the same group have the same number of valence electrons
The number of valence electrons equals the group number for main group elements (Groups 1, 2, and 13-18)

6. Atomic and Ionic Radii

Atomic Radius

Half the distance between the nuclei of two adjacent atoms of the same element
Trends in the periodic table:
Across a period: Decreases from left to right (due to increasing nuclear charge pulling electrons closer)
Down a group: Increases (due to additional electron shells)

Ionic Radius

The radius of an ion
Depends on whether electrons are gained or lost:
Cations (positive ions) are smaller than their parent atoms (lost an electron shell or electrons from the outermost shell)
Anions (negative ions) are larger than their parent atoms (additional electrons cause increased repulsion)

7. Ionization Energy

Definition of Ionization Energy

The energy required to remove an electron from a gaseous atom or ion
Always an endothermic process (energy is absorbed)
Measured in kJ/mol

First, Second, and Successive Ionization Energies

First ionization energy: Energy required to remove the first electron
Second ionization energy: Energy required to remove the second electron (always higher than the first)
Successive ionization energies: Continue to increase as more electrons are removed

Trends in Ionization Energy

Across a period: Increases from left to right (due to increasing nuclear charge and decreasing atomic radius)
Down a group: Decreases (due to increasing atomic radius and shielding effect)

Factors Affecting Ionization Energy

Nuclear charge: Higher nuclear charge increases ionization energy
Distance from nucleus: Electrons farther from the nucleus have lower ionization energy
Shielding effect: Inner electrons shield outer electrons from the full nuclear charge
Electron configuration: Fully filled or half-filled subshells are more stable and require more energy to remove an electron

8. Electron Affinity

Definition of Electron Affinity

The energy change when a gaseous atom gains an electron to form a negative ion
Can be exothermic (energy released) or endothermic (energy absorbed)
Measured in kJ/mol

Trends in Electron Affinity

Across a period: Generally becomes more exothermic from left to right (except for noble gases)
Down a group: Generally becomes less exothermic

Factors Affecting Electron Affinity

Nuclear charge: Higher nuclear charge makes electron affinity more exothermic
Atomic size: Smaller atoms generally have more exothermic electron affinities
Electronic configuration: Atoms that gain electrons to achieve stable configurations (like halogens) have more exothermic electron affinities

9. Electronegativity

Definition of Electronegativity

The ability of an atom to attract shared electrons in a chemical bond
Measured on the Pauling scale (unitless)

Trends in Electronegativity

Across a period: Increases from left to right (fluorine is the most electronegative element)
Down a group: Decreases

Significance in Chemical Bonding

Large electronegativity difference (>1.7): Ionic bond formation is likely
Moderate electronegativity difference (0.5-1.7): Polar covalent bond formation
Small electronegativity difference (<0.5): Nonpolar covalent bond formation

10. Radioactivity and Nuclear Chemistry

Types of Radioactive Decay

Alpha (α) Decay:
- Emission of an alpha particle (²₄He nucleus)
- Decreases atomic number by 2 and mass number by 4
- Example: ²³₈U → ²³₄Th + ⁴₂He
Beta (β) Decay:
- Emission of a beta particle (electron)
- Increases atomic number by 1 but mass number remains the same
- Occurs when a neutron converts to a proton
- Example: ¹⁴C → ¹⁴N + ⁰₋₁e
Gamma (γ) Radiation:
- Emission of high-energy electromagnetic radiation
- No change in atomic number or mass number
- Often accompanies alpha or beta decay
- Example: ⁶⁰Co → ⁶⁰Co* → ⁶⁰Co + γ

Half-life

Time taken for half of the radioactive nuclei in a sample to decay
Independent of the initial amount of material
Unique to each radioisotope

Half-life Calculation:
If the initial number of atoms is N₀ and the number after time t is N:
N = N₀ × (1/2)^(t/t₁/₂)
Where t₁/₂ is the half-life.

Applications of Radioisotopes

Carbon-14 dating: Determining the age of archaeological artifacts
Medical diagnostics: Using tracers like technetium-99m for imaging
Cancer treatment: Using radiation to kill cancer cells
Food preservation: Killing bacteria in food
Smoke detectors: Using americium-241

11. Practice Questions

Concept Questions

Describe the difference between Rutherford's and Bohr's models of the atom.
Explain why atomic radius decreases across a period in the periodic table.
Arrange the following in order of increasing first ionization energy: Na, Mg, Al, Si, P.
Compare and contrast ionic and covalent bonding, relating this to electronegativity differences.

Calculation Problems

An atom has 35 protons, 45 neutrons, and 36 electrons. Determine:
- a) The atomic number
- b) The mass number
- c) The charge on the ion
Write the electron configuration for:
- a) Calcium (Z=20)
- b) Iron (Z=26)
- c) Bromine (Z=35)
Complete the following nuclear equations:
- a) ²³₈U → ²³₄Th + ?
- b) ¹⁴C → ? + ⁰₋₁e
- c) ? → ²³⁴Pa + ⁰₋₁e
A sample of radioactive material has a half-life of 5 days. If the initial mass is 80g, how much remains after 15 days?

12. Exam Tips for CXC Chemistry

Essential Knowledge

Know your definitions: Make sure you can define key terms like atomic number, mass number, isotope, etc.
Practice calculations: Be comfortable calculating numbers of subatomic particles, half-life problems, etc.
Understand trends: Know the periodic trends for atomic radius, ionization energy, electronegativity, etc., and be able to explain WHY these trends occur.

Exam Technique

Draw diagrams: Practice drawing atomic models, electron configuration diagrams, and nuclear decay equations.
Link to chemical properties: Be able to connect atomic structure to chemical behavior, especially for questions about bonding and reactivity.
Use proper terminology: Use the correct scientific language when describing atomic structure.
Know the exceptions: Understand exceptions to general rules, such as why certain electron configurations differ from what's expected.
Show your work: In calculations, clearly show each step to maximize your marks.

13. Additional Resources

Study Techniques

Create flashcards for key concepts and definitions
Draw mind maps connecting related concepts
Use colors to highlight patterns in the periodic table
Practice writing electron configurations regularly
Create mnemonics to remember the order of orbital filling

14. Glossary of Key Terms

Fundamental Concepts

Atom: The smallest particle of an element that retains its chemical properties
Atomic Number: The number of protons in an atom's nucleus
Mass Number: The sum of protons and neutrons in an atom
Isotope: Atoms of the same element with different numbers of neutrons

Electron Structure

Electron Configuration: The arrangement of electrons in an atom's orbitals
Orbital: A region of space where an electron is likely to be found
Valence Electrons: Electrons in the outermost energy level

Atomic Properties

Ionization Energy: Energy required to remove an electron from a gaseous atom
Electron Affinity: Energy change when a gaseous atom gains an electron
Electronegativity: Ability of an atom to attract shared electrons in a bond
Half-life: Time taken for half of a radioactive sample to decay