Chemistry | The nuclear atom

Study Notes: Structure 1.2 – The Nuclear Atom

Hello future Chemist! Welcome to the exciting world of atomic structure. This chapter, "The nuclear atom," is fundamental—it’s where we define the tiny particles that make up all matter. Understanding the structure of the atom is the core foundation for everything else you will study in chemistry, especially within this section on Models of the particulate nature of matter. Don't worry if the concepts seem abstract; we will use clear analogies to build up your understanding step by step!

1. The Core Components: Subatomic Particles

Every atom is made up of three primary subatomic particles. It’s essential to know their location within the atom, their relative mass, and their relative charge.

Analogy: The Atom as a Stadium

Imagine a giant sports stadium. The nucleus is like a tiny marble or golf ball placed right on the center field spot. The rest of the stadium—the empty space, the stands, the air—is where the electrons are flying around incredibly quickly. An atom is mostly empty space!

The particles are:

• Protons (p+)
• Neutrons (n0)
• Electrons (e-)

Summary of Subatomic Particles

This table compares the particles based on their fundamental properties:

Particle	Location	Relative Mass (amu)	Relative Charge
Proton	Nucleus	1	+1
Neutron	Nucleus	1	0 (Neutral)
Electron	Outside the nucleus (Shells/Cloud)	\(1/1836\) (Effectively 0)	-1

Key Takeaway: The nucleus contains almost all the mass (Protons and Neutrons), while the electrons determine the size and chemical behaviour of the atom.

2. Defining the Atom: Atomic and Mass Numbers (SL/HL)

To identify an atom, we use two key numbers. These numbers are crucial for determining the structure of any element.

Atomic Number (Z)

The Atomic Number (Z) is the most important number because it defines the element.

• It equals the number of protons in the nucleus.
• If the number of protons changes, the element itself changes!

Memory Aid: Z marks the Zone (identity).

Mass Number (A)

The Mass Number (A) represents the total number of particles in the nucleus (since electrons have negligible mass).

• It is the sum of protons and neutrons.
• The number of neutrons can be found by: \( \text{Neutrons} = A - Z \).

The Neutral Atom and Ions

In a neutral atom (which has no overall charge):

\( \text{Number of Protons} = \text{Number of Electrons} \)

If an atom gains or loses electrons, it becomes an ion and carries an electrical charge:

• If electrons are lost (fewer e- than p+), the atom becomes a positive ion (cation).
• If electrons are gained (more e- than p+), the atom becomes a negative ion (anion).

Remember: Only the number of electrons changes when an atom forms an ion. The number of protons (Z) never changes in a chemical reaction.

3. Variety in Atoms: Isotopes (SL/HL)

While all atoms of an element must have the same number of protons (same Z), they don't necessarily have the same number of neutrons. This leads us to the concept of isotopes.

What are Isotopes?

Isotopes are atoms of the same element that have the same number of protons (same Z) but a different number of neutrons (different A).

Example: Hydrogen Isotopes
Hydrogen (Z=1) has three common isotopes, all with 1 proton:

• Hydrogen-1 (Protium): 1 proton, 0 neutrons (A=1).
• Hydrogen-2 (Deuterium): 1 proton, 1 neutron (A=2). This is used in heavy water.
• Hydrogen-3 (Tritium): 1 proton, 2 neutrons (A=3). This is radioactive.

Properties of Isotopes

• Chemical Properties: Isotopes have virtually identical chemical properties because these properties are determined by the electrons, and isotopes have the same number of electrons.
• Physical Properties: Isotopes have different physical properties (like density and melting/boiling points) because of the difference in mass.

Did you know? Many isotopes are radioactive, meaning their nuclei are unstable and decay over time, emitting radiation. This instability often occurs when the ratio of neutrons to protons is too high or too low.

Key Takeaway: Isotopes are "versions" of an element. They have the same chemical identity (Z) but different masses (A).

4. Calculating Average Mass: Relative Atomic Mass (\(A_r\)) (SL/HL)

When you look at the Periodic Table, the atomic mass listed (e.g., Carbon is 12.01) is rarely a whole number. This is because it is the average mass of all the naturally occurring isotopes of that element.

Definition: Relative Atomic Mass (\(A_r\))

The Relative Atomic Mass (\(A_r\)) is the weighted average of the masses of all naturally occurring isotopes of an element, relative to \(1/12\) the mass of a Carbon-12 atom.

Since the different isotopes don't exist in equal amounts, we must calculate a weighted average based on their percentage abundance.

Step-by-Step Calculation of \(A_r\)

To calculate the weighted average, you need the mass of each isotope and its natural abundance (usually given as a percentage).

Formula:

$$ A_r = \sum (\text{Isotope Mass} \times \text{Fractional Abundance}) $$

Example Steps (Using Element X):

Assume Element X has two isotopes:

Isotope X-20: Mass = 20.00 amu, Abundance = 75.0%

Isotope X-22: Mass = 22.00 amu, Abundance = 25.0%

1. Convert Abundance to Fractional Abundance (Decimal):
   • \( 75.0\% \rightarrow 0.750 \)
   • \( 25.0\% \rightarrow 0.250 \)
2. Calculate the Contribution of Each Isotope:
   • Contribution 1: \( 20.00 \times 0.750 = 15.00 \)
   • Contribution 2: \( 22.00 \times 0.250 = 5.50 \)
3. Sum the Contributions:
   • \( A_r = 15.00 + 5.50 = 20.50 \text{ amu} \)

The Relative Atomic Mass (\(A_r\)) of Element X is 20.50.

Crucial Tip for Struggling Students: The calculated \(A_r\) must always fall between the lightest and heaviest isotopic mass. In the example above, 20.50 is between 20.00 and 22.00. Since 20.50 is closer to 20.00, it tells you that the lighter isotope (X-20) is the most abundant, which matches the data (75%).

5. Measuring Mass: The Mass Spectrometer (SL/HL Concept)

How do scientists know the exact mass and natural abundance of isotopes? They use a powerful tool called the Mass Spectrometer.

While you do not need to know the complex engineering, you must understand the basic principle: it separates ions based on their mass-to-charge ratio (\(m/z\)).

The Purpose and Process

The mass spectrometer determines the exact masses and the relative abundances of the isotopes of an element.

The process generally involves four main stages (I.A.D.D.):

1. Ionization: The sample atoms are turned into positive ions (cations), usually by knocking off an electron.
2. Acceleration: These positive ions are accelerated by an electric field into a beam.
3. Deflection: The ion beam passes through a strong magnetic field. The magnetic field deflects the lighter ions (or those with higher charge) more than the heavier ions.
4. Detection: The detector records where the ions hit, giving two pieces of data:
   • The position of impact reveals the mass (how much it was deflected).
   • The intensity of the signal reveals the relative abundance (how many ions of that mass hit the detector).

The final result is a mass spectrum (a graph) that shows peaks corresponding to the masses of the isotopes and the height of the peaks corresponding to their relative abundance.

Analogy: Imagine throwing metal balls of different weights (the isotopes) through a strong fan (the magnetic field). The fan pushes the lighter balls further off course (greater deflection) than the heavy balls.

Key Takeaway: The mass spectrometer provides the necessary data (masses and abundances) to accurately calculate the relative atomic mass (\(A_r\)) listed on the periodic table.