State that isotopes of the same element have the same chemical properties because they have the same number of electrons and therefore the same electronic configuration

Isotopes – Same Chemical Behaviour, Different Nuclei

1. Definition of isotopes (Core 2.3)

• An isotope is one of two or more atoms of the same element that have the same atomic number (Z) – i.e. the same number of protons – but different mass numbers (A) because they contain different numbers of neutrons.
• Mass number: \(A = Z + N\) where \(N\) = number of neutrons.

2. Reading an isotope symbol (Core)

The general form is ^{mass number}Element‑symbol_{atomic number} (charge is optional). Examples:

• ¹²C – carbon‑12 (mass number 12, atomic number 6)
• ³⁵Cl – chlorine‑35 (mass number 35, atomic number 17)
• ¹⁴C⁺ → carbon‑14 ion (mass number 14, atomic number 6, +1 charge)

3. Nuclear composition of isotopes (Core)

• Protons (p⁺): determine the element (the atomic number Z).
• Neutrons (n⁰): vary between isotopes, giving different mass numbers and influencing nuclear stability.
• Electrons (e⁻): a neutral atom has the same number of electrons as protons ( e⁻ = Z ). Ions have a different electron count, but the element (and its chemistry) is still defined by Z.

4. Electronic configuration and chemical behaviour (Core)

All chemical properties – bonding, ionisation, reactivity, etc. – are governed by the arrangement of electrons around the nucleus. Because every isotope of an element has the same number of electrons, they all have the same electronic configuration and therefore exhibit identical chemical behaviour.

5. Physical consequences of different neutron numbers (Supplement 1.2)

• Isotopes differ in mass, so physical properties such as density, boiling point, and rate of diffusion are affected.
• According to Graham’s law, the rate of diffusion is inversely proportional to the square‑root of the molar mass; therefore a heavier isotope diffuses more slowly than a lighter one.

6. Using isotopic data to calculate relative atomic mass (Supplement 2.3)

For any element the relative atomic mass \(A_r\) is the weighted average of the mass numbers of its naturally occurring isotopes:

\[

Ar = \sum{i} \bigl(fi \times Ai\bigr)

\]

where \(fi\) = fractional (decimal) natural abundance of isotope i and \(Ai\) = its mass number.

Worked example 1 – chlorine

1. Identify the isotopes and their fractional abundances:

   • \(^{35}\)Cl, \(f = 0.758\)
   • \(^{37}\)Cl, \(f = 0.242\)

2. Apply the formula:

   \[

   A_r(\text{Cl}) = (0.758 \times 35) + (0.242 \times 37)

   = 26.53 + 8.95 = 35.48 \approx 35.5

   \]

Worked example 2 – carbon

1. Isotopes present naturally:

   • \(^{12}\)C, \(f = 0.989\)
   • \(^{13}\)C, \(f = 0.011\)

2. Calculate:

   \[

   A_r(\text{C}) = (0.989 \times 12) + (0.011 \times 13)

   = 11.868 + 0.143 = 12.011 \approx 12.0

   \]

7. Typical isotopes (Core)

8. Link to later topics (Stoichiometry)

When performing stoichiometric calculations you use the relative atomic mass (\(A_r\)) given on the periodic table. This value already incorporates the contribution of all naturally occurring isotopes, so you do not need to consider isotopic composition unless a problem explicitly mentions a particular isotope (e.g., radiocarbon dating with \(^ {14}\)C).

9. Key points to remember (Core – exam focus)

• Isotopes have the same atomic number (Z) → same element.
• All isotopes of an element have the same number of electrons in a neutral atom, giving an identical electronic configuration and therefore identical chemical properties.
• Differences in neutron number affect only the mass of the nucleus and related physical properties (density, diffusion rate), not the chemistry.