Every calculation and bonding model in chemistry begins with a disciplined picture of the atom. This chapter separates what is fixed in the nucleus from what can change when ions form, then uses isotopes and abundance data to explain why periodic-table masses are usually not whole numbers.

Write a nuclide as ${}_Z^{A}X$, where $Z$ is the proton number and $A$ is the mass number. Protons determine the element: any atom with $Z=17$ is chlorine. The neutron number is $A-Z$. In a neutral atom, electrons equal protons because the total charge is zero.

The mass number is not the periodic-table value. It applies to one isotope and is always a whole number, whereas $A_r$ is an abundance-weighted mean for a natural sample. Keep the labels visible in calculations: this prevents confusing 35Cl with chlorine's approximate $A_r$ of 35.5.

Atomic number $Z$ counts protons; mass number $A$ counts protons plus neutrons. Isotopes have the same $Z$ but different neutron numbers, so neutral atoms of an element have similar electron arrangements but different masses. A positive ion has lost electrons, not protons. Relative atomic mass is an abundance-weighted mean of isotope masses, not usually a whole number. Mass-spectrometric peak heights must be interpreted as abundances under the given measurement model.

Ion formation changes electrons, never protons or neutrons. A positive ion has lost electrons, so $electrons=protons$ minus the positive charge. A negative ion has gained electrons, so $electrons=protons$ plus the magnitude of the negative charge.

For example, ${}_{13}^{27}{\mathrm{Al}}^{3+}$ contains 13 protons, 14 neutrons and 10 electrons. It is still aluminium after losing three electrons because its nucleus still contains 13 protons. This distinction is essential when linking electron loss to oxidation and bonding later in the course.

When proton, neutron and electron beams enter the same uniform electric field at the same speed, separate direction from amount of deflection. A charged particle feels an electric force; an uncharged neutron does not. Positive and negative particles curve towards opposite plates.

Isotopes have nearly identical chemical properties because chemical behaviour depends mainly on the electron arrangement, which is the same for neutral isotopes of one element. Their different masses affect physical processes, including diffusion, and produce separate peaks in a mass spectrum.

A mass spectrometer turns particles into ions, separates them by mass-to-charge ratio, and detects their relative abundance. A peak at a particular $m/z$ value is evidence for ions of a particular isotope or molecular fragment; peak height or area represents relative abundance, not necessarily a count of atoms in one molecule.

A sample with more of a heavy isotope has a larger $A_r$. Multiply each isotopic mass by its fractional abundance, add the contributions, and round only at the end. When abundances are percentages, either divide each by 100 first or divide the final sum by 100.

The result should lie between the lightest and heaviest isotopic masses and closer to the more abundant isotope. That simple estimate is a powerful error check before committing to a final answer.

Across a period, atomic radius generally decreases. Proton number rises while added electrons enter the same principal shell, so shielding changes little and the stronger nuclear attraction pulls the outer electrons closer. Down a group, radius increases because an additional occupied shell places the outer electrons farther from the nucleus and increases shielding.

Do not describe radius as depending only on proton number. A large nuclear charge can be outweighed by extra shells and shielding. Compare species fairly: a set of isoelectronic ions has the same number of electrons, so the species with more protons is smaller because the same electron cloud is pulled in more strongly.

A mass spectrum is more than a list of isotope masses. The position of a peak gives an ion's mass-to-charge ratio, while its relative abundance gives evidence about how common that ion is in the sample. For a singly charged atomic ion, the m/z value is numerically the isotope mass. A weighted mean of the isotope masses gives the relative atomic mass printed in a Periodic Table; it is not the mass of any one atom.

The same logic extends to molecules. A diatomic molecule made from two isotopes produces more than one molecular-ion peak because different isotope pairings are possible. For chlorine, the 35Cl–35Cl, 35Cl–37Cl, and 37Cl–37Cl combinations give an approximately 9:6:1 cluster. The middle peak is twice as likely as either mixed ordering because the two atoms can occupy either position.

A spectrum is experimental evidence for isotopic composition. Treat the supplied peak data as measurements: identify the peaks, use their relative heights or areas, and state the conclusion with suitable precision.

A short, labelled inventory prevents the most common ion and isotope errors.