The nucleus and nucleons

• An atom contains a dense, positively charged nucleus, with negatively charged electrons occupying the space outside it.

• The nucleus contains protons and neutrons, collectively called nucleons.

• Protons carry positive charge, while neutrons have no charge.

• Almost all atomic mass is concentrated in the nucleus because the electron mass is negligible.

• The number of protons is the atomic number $Z$.

The diagram represents the nucleus as an assembly of protons and neutrons. Together, these particles are called nucleons and account for almost all the atom's mass. Source

Using the nuclear symbol ${}^{A}_{Z}X$

• The nuclear symbol is written as ${}^{A}_{Z}X$, where $X$ is the element symbol.

• $Z$ is the atomic number and equals the number of protons.

• $A$ is the mass number and equals the total number of nucleons: $A=p+n$.

• Number of protons $=Z$.

• Number of neutrons $=A-Z$.

• For a neutral atom, number of electrons $=Z$.

Isotopes

• Isotopes are atoms of the same element containing different numbers of neutrons.

• Isotopes therefore have the same atomic number $Z$ but different mass numbers $A$.

• The proton number remains constant because changing the proton number would produce a different element.

• Different neutron numbers give isotopes different masses.

• Their physical properties can differ where those properties depend on particle mass.

• Specific isotope examples do not need to be memorized for this syllabus statement.

Finding isotopic abundance

• For two isotopes, if one abundance is $x$, the other abundance is $(100-x)$.

• Insert these expressions into the weighted-mean equation for $A_r$.

• For isotope masses $m_1$ and $m_2$: $A_r=\frac{m_1x+m_2(100-x)}{100}$.

• Rearrange the equation to calculate the unknown value of $x$.

• Check that every abundance lies between $0$ and $100$.

• All percentage abundances must total $100$.

Subatomic particles

Deducing electrons in ions

• Forming an ion changes the number of electrons, not the numbers of protons or neutrons.

• For a positive ion of charge magnitude $q$, electrons $=Z-q$.

• For a negative ion of charge magnitude $q$, electrons $=Z+q$.

• For ${}^{24}_{12}\mathrm{Mg}^{2+}$: protons $=12$, neutrons $=12$, electrons $=10$.

• For ${}^{19}_{9}\mathrm{F}^{-}$: protons $=9$, neutrons $=10$, electrons $=10$.

• Exam check: never change the proton number to account for ionic charge.

Calculating relative atomic mass

• Relative atomic mass $A_r$ is calculated as an abundance-weighted mean of the isotopic masses.

• For percentage abundances: $A_r=\frac{\sum(m_i a_i)}{100}$.

• More generally: $A_r=\frac{\sum(m_i a_i)}{\sum a_i}$.

• A non-integer $A_r$ results when an element occurs naturally as isotopes with different masses and abundances.

• Example: isotopic masses $10$ and $11$ at $20$ and $80$ give $A_r=\frac{10(20)+11(80)}{100}=10.8$.

• Always multiply each isotopic mass by its corresponding relative abundance.