Particles in the atom and atomic radius - Chemistry (9701) - Cambridge A-Level

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Chemistry 9701 Study Notes: Particles in the Atom and Atomic Radius (AS Level 1.1)

Welcome to the building blocks of Chemistry! This chapter, "Particles in the Atom," is absolutely fundamental. Understanding what atoms are made of and how big they are helps us explain why elements behave the way they do—from how they react to how they form chemical bonds.

Don't worry if the concepts seem abstract; we will use analogies and clear steps to break everything down!

1. The Structure of the Atom: The Tiny Universe

1.1 The Nuclear Model

Imagine an atom is like a massive sports stadium. The nucleus would be a tiny fly sitting right in the centre spot. That gives you an idea of the scale!

The syllabus requires you to understand that the atom is mostly empty space surrounding a very small, dense centre called the nucleus (LO 1).

The Nucleus contains two types of particles: protons and neutrons.
Electrons are found orbiting the nucleus in specific paths called shells (or energy levels) in the vast empty space outside the nucleus.

1.2 The Subatomic Particles: Mass and Charge

Atoms are made of three essential subatomic particles. It is vital to know their relative masses and charges (LO 2).

Quick Review Table: Subatomic Particles

Particle	Location	Relative Mass	Relative Charge
Proton	Nucleus	1	+1
Neutron	Nucleus	1	0 (Neutral)
Electron	Shells/Orbitals	\( \frac{1}{1836} \) (effectively 0)	-1

Memory Aid: P A N D A

A simple way to remember the relationship between protons, electrons, and charge in a neutral atom:

Protons = Atomic Number = Neutrons (not always true, but often close to atomic number for small elements) = Different from electrons in ions = Atoms are neutral.

1.3 Distribution of Mass and Charge (LO 4)

Because protons and neutrons both have a relative mass of 1, and electrons have virtually zero mass:

The vast majority of the atom's mass is concentrated in the nucleus.
The charge of the atom is distributed between the positive nucleus (protons) and the negative electron shells (electrons). In a neutral atom, these charges exactly cancel out.

Key Takeaway for Section 1: Atoms have a tiny, dense, positively charged nucleus (p + n) surrounded by negatively charged electrons occupying most of the atom’s volume.

2. Defining Elements: Atomic and Mass Numbers

2.1 Key Definitions (LO 3)

These terms are fundamental for identifying elements and calculating particle numbers:

Atomic Number (\(Z\)) (or Proton Number): The number of protons in the nucleus of an atom.
This number defines the element. If you change the number of protons, you change the element!
Mass Number (\(A\)) (or Nucleon Number): The total number of protons and neutrons (nucleons) in the nucleus.

How to Calculate the Number of Neutrons:

\[ \text{Number of Neutrons} = \text{Mass Number} (A) - \text{Proton Number} (Z) \]

2.2 Calculating Particles in Atoms and Ions (LO 6)

To determine the number of protons, neutrons, and electrons, follow these steps:

Step 1: Find the Protons (\(p\))

\( p = Z \) (Atomic Number). This never changes for a given element.

Step 2: Find the Neutrons (\(n\))

\( n = A - Z \) (Mass Number minus Atomic Number).

Step 3: Find the Electrons (\(e^-\))

For a neutral atom: \( e^- = p \) (The charges must balance).
For an ion: You must adjust the electron count based on the charge:
- If the ion is positive (a cation, e.g., Na\(^+\)), it has lost electrons. \( e^- = p - \text{Charge} \).
- If the ion is negative (an anion, e.g., Cl\(^-\)), it has gained electrons. \( e^- = p + \text{Charge} \).

Example: The Oxide Ion, O\(^{2-}\)

Oxygen has \(Z=8\) and a typical \(A=16\).

Protons: \( p = 8 \)
Neutrons: \( n = 16 - 8 = 8 \)
Electrons: The 2- charge means it gained 2 electrons. \( e^- = 8 + 2 = 10 \).

Key Takeaway for Section 2: Atomic number identifies the element. Mass number dictates the total particles in the nucleus. Ions have a different number of electrons than their parent atom.

3. Behaviour in an Electric Field (LO 5)

We can prove the properties of subatomic particles by observing how beams of them behave when passed through an electric field (like two charged plates).

3.1 Electric Field and Deflection

If beams of protons, neutrons, and electrons are accelerated to the same velocity and then passed through an electric field, they will behave differently based on their charge and mass.

Neutrons: Since they are neutral (charge = 0), they are not deflected at all.

Protons: Since they are positive (charge = +1), they are deflected towards the negative plate.

Electrons: Since they are negative (charge = -1), they are deflected towards the positive plate.

3.2 The Key Factor: Charge-to-Mass Ratio

The amount of deflection depends on the charge-to-mass ratio (\( Q/M \)):

Greater charge = Greater deflection
Smaller mass = Greater deflection

The electron has a charge of -1 but a mass that is nearly zero. Therefore, the electron has an extremely large magnitude for its charge-to-mass ratio compared to the proton (which has mass 1).
Result: Electrons are deflected the most significantly because they are so light.

Did you know? This principle is crucial in mass spectrometry, a technique used to analyze isotopes and molecular structures!

Key Takeaway for Section 3: Charged particles are deflected by an electric field; neutral particles are not. Electrons deflect the most due to their very low mass.

4. Atomic Radius and Periodic Trends (LO 7)

The atomic radius is a measure of the size of an atom. Since we cannot precisely define the edge of an electron cloud, it is usually defined as half the distance between the nuclei of two identical, bonded atoms.

4.1 Trends Down a Group (e.g., Li to Na to K)

As you move down a group in the Periodic Table, the atomic radius increases.

Explanation:

More Electron Shells: Each new element down the group adds a completely new principal quantum shell (a new energy level) for the electrons.
Increased Shielding: The inner shells of electrons "shield" the outer valence electrons from the full attractive force of the nucleus.
The outer electrons are further from the nucleus and feel less attraction, leading to a larger atomic size.

4.2 Trends Across a Period (e.g., Na to Ar)

As you move across a period (from left to right) in the Periodic Table, the atomic radius decreases.

Explanation:

Increasing Nuclear Charge: The number of protons (nuclear charge) increases steadily (e.g., Na has 11, Mg has 12, P has 15, etc.).
Same Principal Shell: All valence electrons are being added to the same main energy shell.
Stronger Attraction: The increased positive nuclear charge pulls the electron cloud closer to the nucleus, causing the atom to shrink. (Shielding remains roughly constant as only valence electrons are being added).

4.3 Ionic Radius vs. Atomic Radius

Ionic radius follows similar periodic trends, but you must first compare the ion to its parent atom:

A. Cations (Positive Ions) are SMALLER than their Parent Atoms

Example: Na atom vs. Na\(^+\) ion.
Why? When forming a cation, the atom loses its outermost electron shell entirely (e.g., Na loses its 3s electron). The remaining electrons are drawn closer by the same number of protons.

B. Anions (Negative Ions) are LARGER than their Parent Atoms

Example: Cl atom vs. Cl\(^-\) ion.
Why? When forming an anion, the atom gains electrons into the existing outer shell. This increases the inter-electron repulsion among the valence electrons, causing the electron cloud to swell and increase the radius.

Crucial Point for Ionic Trends:

For a series of isoelectronic ions (ions with the same number of electrons, e.g., O\(^{2-}\), F\(^-\), Na\(^+\), Mg\(^{2+}\)), the size is determined purely by the nuclear charge (number of protons).

The one with the most protons (Mg\(^{2+}\) with 12 protons) will pull the electron cloud the tightest, making it the smallest.
The one with the fewest protons (O\(^{2-}\) with 8 protons) will have the weakest pull, making it the largest.

Key Takeaway for Section 4: Atomic size increases down a group (more shells/shielding) but decreases across a period (greater nuclear charge pulling the same shell tighter). Cations are smaller; anions are larger.

Quick Review Box: Particles and Radius (1.1 Summary)

Atom Structure: Dense nucleus (p + n) surrounded by electrons in shells.
Identifying Particles: Protons (\(Z\)), Neutrons (\(A-Z\)), Electrons (adjust for charge).
Electric Field: Electrons deflect most; neutrons don't deflect at all.
Atomic Radius Down Group: Increases (due to addition of shells).
Atomic Radius Across Period: Decreases (due to increased nuclear charge).
Ionic Size: Cations < Atom, Anions > Atom.

You've successfully covered the core concepts of atomic particles and size! This foundational knowledge will be vital for understanding bonding and periodicity later on.