Atomic Structure and the Periodic Table

Introduction

Asalam-o-Alaikum, students. I am Dr. Fatima Malik. Welcome to this essential topic which serves as the bedrock for your entire A Level Chemistry journey. Understanding atomic structure is like learning the alphabet before you can read a book. Every chemical reaction, every physical property, and every trend you will study – from the reactivity of metals used in Pakistani steel mills to the function of fertilisers developed by ENGRO – is governed by the principles we will cover here.

The atom is the fundamental unit of an element, and its structure dictates its identity and chemical behaviour. We will explore the subatomic particles that constitute the atom, how they are arranged, and how this arrangement gives rise to the ordered beauty of the Periodic Table. Mastering this topic will not only secure you marks in Paper 2 but will provide the conceptual framework necessary to tackle more complex topics in A2 Chemistry, such as transition metal chemistry and reaction mechanisms. Pay close attention, as a strong foundation here is non-negotiable for success.

Core Theory

Fundamental Particles

All atoms are composed of three fundamental particles. You must know their relative mass and charge by heart.

| Particle | Relative Mass | Relative Charge | Location |

|----------|---------------|-----------------|----------|

| Proton | 1 | +1 | Nucleus |

| Neutron | 1 | 0 | Nucleus |

| Electron | 1/1836 (≈0) | -1 | Orbitals |

The Proton Number (Z), also called the atomic number, defines the element. The Nucleon Number (A), or mass number, is the total number of protons and neutrons in the nucleus.

Isotopes and Relative Atomic Mass (Ar)

Isotopes are atoms of the same element (same Z) but with a different number of neutrons (different A). For example, chlorine exists mainly as Chlorine-35 (17p, 18n) and Chlorine-37 (17p, 20n). Because of isotopes, the mass of an element is an average. The Relative Atomic Mass (Ar) is the weighted mean mass of an atom of an element compared to 1/12th the mass of a carbon-12 atom.

We calculate it using data from a mass spectrometer:

Ar = Σ (isotopic mass × % abundance) / 100

The Mass Spectrometer

This instrument provides evidence for isotopes and helps determine Ar. Its operation involves four key stages:

1. Ionisation: The sample is vaporised and a high-energy electron beam knocks an electron off the atoms/molecules, forming positive ions (e.g., Mg(g) → Mg⁺(g) + e⁻).
2. Acceleration: The positive ions are accelerated by an electric field so they all have the same kinetic energy.
3. Deflection: The ions are deflected by a magnetic field. The amount of deflection depends on the mass-to-charge ratio (m/z). Lighter ions are deflected more; heavier ions are deflected less.
4. Detection: The deflected ions hit a detector, which measures the abundance of each ion. A mass spectrum is produced, plotting relative abundance against m/z.

ASCII Diagram of Deflection:

Ions from Accelerator

│

▼

┌───────────────────┐

│ Magnetic Field │

│ (into page) │

│ .------> Detector (Heavy ions, e.g., ³⁷Cl⁺)

│ / │

│ / │

│ '-------------> Detector (Light ions, e.g., ³⁵Cl⁺)

└───────────────────┘

Electronic Configuration

Electrons exist in principal energy levels or shells (n=1, 2, 3...). These shells contain subshells (s, p, d, f), which in turn contain orbitals. An orbital is a region of space where there is a 95% probability of finding an electron. Each orbital can hold a maximum of 2 electrons.

• s-subshell: 1 spherical orbital (holds 2e⁻)
• p-subshell: 3 dumbbell-shaped orbitals (px, py, pz), (holds 6e⁻)
• d-subshell: 5 complex-shaped orbitals (holds 10e⁻)

Electrons fill these orbitals according to three rules:

1. Aufbau Principle: Orbitals are filled in order of increasing energy (1s, 2s, 2p, 3s, 3p, 4s, 3d...).
2. Pauli Exclusion Principle: An orbital can hold a maximum of two electrons, and they must have opposite spins (represented as ↑ and ↓).
3. Hund's Rule: When filling degenerate orbitals (orbitals of the same energy, like the three p-orbitals), electrons fill each orbital singly before any orbital is doubly occupied.

Example: Electronic configuration of Silicon (Z=14): 1s² 2s² 2p⁶ 3s² 3p².

The 3p electrons would be arranged as: [↑ ] [↑ ] [ ] in the px, py, and pz orbitals.

Periodic Trends: First Ionisation Energy (IE)

First Ionisation Energy is the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

Equation: X(g) → X⁺(g) + e⁻

Across Period 3 (Na to Ar):

• General Trend: First IE increases.
• Reason: Across the period, the number of protons (nuclear charge) increases. Electrons are being added to the same principal shell, so shielding remains relatively constant. This leads to a stronger electrostatic force of attraction between the nucleus and the outer electron, requiring more energy to remove it.
• Anomalies (Dips):
• Mg to Al: There is a slight decrease. The outer electron in Al (3p¹) is in a higher energy subshell than in Mg (3s²). The 3p subshell is also shielded by the 3s electrons. Therefore, less energy is needed to remove this electron.
• P to S: There is a slight decrease. In Phosphorus (3p³), each p-orbital is singly occupied. In Sulfur (3p⁴), one p-orbital contains a pair of electrons. The repulsion between these two spin-paired electrons makes it easier to remove one of them, lowering the IE.

Down Group 2 (Be to Ba):

• Trend: First IE decreases.
• Reason: Moving down the group, the number of electron shells increases. Although nuclear charge increases, the effect is outweighed by the increased **shielding** from inner shells and the larger **atomic radius**. The outer electron is further from the nucleus and feels a weaker attraction, making it easier to remove.

Key Definitions

• Proton Number (Z): The number of protons in the nucleus of an atom.
• Nucleon Number (A): The total number of protons and neutrons in the nucleus of an atom.
• Isotopes: Atoms of the same element with the same number of protons but a different number of neutrons.
• Relative Atomic Mass (Ar): The weighted mean mass of an atom of an element relative to 1/12th of the mass of a carbon-12 atom.
• First Ionisation Energy: The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
• Atomic Orbital: A region of space around the nucleus that can hold up to two electrons with opposite spins.
• Electronegativity: The ability of an atom to attract the bonding pair of electrons in a covalent bond.

Worked Examples (Pakistani Context)

Example 1: Calculating Ar from Mass Spectrometry Data

A mass spectrometer analysis of a sample of lead (Pb) found in the Saindak Copper-Gold Project in Balochistan shows four main isotopes. Calculate the relative atomic mass of this lead sample to two decimal places.

| Isotope | m/z | Relative Abundance (%) |

|---------|-----|------------------------|

| ²⁰⁴Pb | 204 | 1.4 |

| ²⁰⁶Pb | 206 | 24.1 |

| ²⁰⁷Pb | 207 | 22.1 |

| ²⁰⁸Pb | 208 | 52.4 |

Solution:

We use the formula: Ar = Σ (isotopic mass × % abundance) / 100

Step 1: Multiply the mass of each isotope by its percentage abundance.

(204 × 1.4) = 285.6

(206 × 24.1) = 4964.6

(207 × 22.1) = 4574.7

(208 × 52.4) = 10899.2

Step 2: Sum these values.

1. 6 + 4964.6 + 4574.7 + 10899.2 = 20724.1

Step 3: Divide by 100.

Ar = 20724.1 / 100 = 207.241

Answer: The relative atomic mass of lead is **207.24** (to 2 d.p.).

Example 2: Atomic Structure and Fertiliser Production

ENGRO Fertilizers, with its major plant in Daharki, Sindh, is one of Pakistan's largest producers of urea, a nitrogen-based fertiliser. The production starts with the Haber process, reacting nitrogen and hydrogen to make ammonia (NH₃). The properties of nitrogen are directly related to its atomic structure.

Question: Explain, with reference to its electronic configuration, why nitrogen gas (N₂) is very unreactive, requiring the high temperatures and pressures of the Haber process.

Solution:

1. Identify Nitrogen's position and configuration: Nitrogen is in Group 15, Period 2. Its atomic number is 7. Its electronic configuration is 1s² 2s² 2p³. It has 5 valence electrons.

1. Describe the bonding: To achieve a stable octet, each nitrogen atom needs to gain 3 electrons. It does this by forming a triple covalent bond with another nitrogen atom, sharing three pairs of electrons.

N ≡ N

1. Relate bonding to energy: This N≡N triple bond is extremely strong and has a very high bond enthalpy (945 kJ mol⁻¹). A large amount of energy is required to break this bond before the nitrogen atoms can react with hydrogen.

1. Conclusion: The unreactivity (or inertness) of nitrogen gas is due to the strength of the N≡N triple bond, a direct consequence of nitrogen's electronic structure and its need to share three electrons to complete its outer shell. This is why industrial processes like the Haber process at the ENGRO plant require extreme conditions (high temperature and pressure, plus an iron catalyst) to provide sufficient energy to break this bond and initiate the reaction.

Exam Technique

• Paper 2 (Structured Questions): For questions asking you to "explain the trend" in ionisation energy or atomic radius, use a structured approach. I call it the **S-E-C** method:

1. State the trend (e.g., "First ionisation energy increases across Period 3").
2. Explain the reasons, always mentioning the three key factors: nuclear charge, electron shielding, and atomic radius/distance of the outer electron from the nucleus.
3. Compare the factors. For example, "Although shielding increases, the increase in nuclear charge has a greater effect, leading to a stronger attraction."

• For IE anomaly questions: Be precise. For the Mg-Al dip, state that the electron is removed from a higher energy 3p subshell, which is further shielded by the 3s electrons. For the P-S dip, you MUST mention **electron-pair repulsion** in the 3p orbital of sulfur.
• Common Mistakes:
• Confusing "nuclear charge" (number of protons) with "effective nuclear charge" or "nuclear attraction". Be precise.
• Stating that shielding "cancels out" the nuclear charge. It only lessens the attraction.
• Forgetting to mention that the atoms must be in a gaseous state in the definition of ionisation energy. This is a classic mark to lose.
• Mark Scheme Tips: Examiners look for key phrases. Ensure "nuclear charge", "shielding", and "atomic radius" are in your answers for periodic trends. When describing electronic configurations, use the standard notation (e.g., 1s² 2s²...) unless asked for 'electrons-in-boxes'.