Atomic structure and the periodic table

Atomic structure and the periodic table

An atom contains a small positively charged nucleus surrounded by electrons. The nucleus contains protons and neutrons. A proton has relative charge +1 and an electron has relative charge −1; a neutron has no electrical charge. Protons and neutrons have similar relative masses, while an electron has a much smaller mass. Most atomic mass is therefore concentrated in the nucleus.

The atomic number Z is the number of protons and identifies an element. The mass number A is the total number of protons and neutrons in one atom. Neutron number is A − Z. A neutral atom has equal numbers of protons and electrons. An ion forms when electrons are gained or lost; ordinary ion formation does not change the number of protons or turn the element into another element.

Isotopes are atoms of the same element with the same proton number but different neutron numbers. Carbon-12 and carbon-14 both contain six protons, but they contain six and eight neutrons respectively. They have closely related chemical behaviour because their neutral atoms have the same electron arrangement, although their masses and some physical properties differ. Some isotopes are radioactive, but isotope does not itself mean radioactive.

For the first twenty elements, a simple shell model is useful for introductory electron arrangements. Sodium has eleven electrons arranged 2,8,1 in a neutral atom, while chlorine has seventeen arranged 2,8,7. The electrons in the outermost occupied shell strongly influence common bonding and ion formation. This model is a learning representation; more detailed atomic theory describes orbitals and energy levels rather than little planetary tracks.

The periodic table arranges elements by increasing atomic number. Elements in a main group often have related outer-electron patterns and chemical properties. Periods relate to occupied electron shells in the simple model. Metals commonly lose electrons when forming positive ions, while many nonmetals gain electrons or share them in covalent bonds. Do not assume every atom follows one simple reaction pattern in every compound.

When reading nuclide notation or a table of particles, distinguish mass number from relative atomic mass. Relative atomic mass is a weighted mean for the element’s isotopic mixture compared with the carbon-12 reference. It need not be a whole number. In calculations involving ions, first count protons and neutrons, then adjust the electron count using the ion charge; a 2+ ion has lost two electrons relative to the neutral atom.

Using atomic and ionic notation

In nuclide notation, write the mass number at the upper left and atomic number at the lower left of the element symbol; any ionic charge belongs at the upper right. For an aluminium-27 ion Al³⁺ with Z = 13, there are 13 protons, 14 neutrons and 10 electrons. Subtracting electrons creates positive charge because protons then outnumber electrons. Adding electrons creates negative charge. The neutron count is independent of ordinary ionic charge, so do not subtract three neutrons when interpreting Al³⁺. Atomic number cannot change while the atom remains the same element. Distinguish an atom, an ion and a molecule before interpreting any particle-count table.

What the historical models explained

Dalton's early atomic theory helped organise chemical-combination patterns but did not describe subatomic structure or isotopes. Discovery of electrons showed that atoms were divisible into smaller components. Rutherford-type scattering evidence suggested a small concentrated nucleus and mostly empty space. Bohr's model introduced allowed energies to explain important features of hydrogen spectra, but it was not a complete account of all atoms. Modern orbital descriptions improve on fixed planetary paths. These models illustrate scientific revision: a model is useful within its evidence and limitations, not a drawing that must remain literally true forever. Explain which observation a model addresses instead of memorising names without their associated ideas.

Shells, subshells and first-element configurations

For introductory configurations, an s subshell holds at most two electrons and a p subshell at most six. Carbon can be written 1s² 2s² 2p², and sodium 1s² 2s² 2p⁶ 3s¹. For the first twenty neutral atoms, the common filling sequence includes 1s, 2s, 2p, 3s, 3p and then 4s. Calcium therefore has a simple shell arrangement 2,8,8,2. A shell's theoretical maximum capacity is not a rule saying every lower shell must reach that maximum before the next shell begins to fill. Use a consistent configuration scheme and do not treat a drawn shell as a rigid physical track.

Periodic position and outer electrons

In the simple main-group approach, the period indicates the highest occupied principal shell. Related valence configurations help explain similar reactions within a group. Group 1 elements commonly form +1 ions, Group 2 +2 ions and Group 17 −1 ions in familiar ionic compounds. Noble-gas outer configurations are associated with relative chemical unreactivity under ordinary conditions, although not all are completely incapable of forming compounds. Group labels vary between older school conventions and modern numbering, so identify the element or family when ambiguity is possible. Transition elements can show variable oxidation states and coloured compounds; a simple main-group charge shortcut should not be applied to every metal.

Explaining periodic trends

Across a period, increasing nuclear charge generally pulls outer electrons closer when shielding does not increase enough to offset it, so atomic radius tends to decrease. Down a main group, additional occupied shells and shielding generally increase atomic radius. First ionisation energy is the energy required to remove one electron from each gaseous atom in one mole of gaseous atoms. It generally rises across a period and falls down a group, with subshell and pairing exceptions. Electronegativity describes attraction for bonding electrons rather than the energy of removing an isolated electron. A trend explanation needs nuclear attraction, distance and shielding, not just the statement that the periodic table says so.

Relative atomic mass calculations

Relative atomic mass is a weighted isotopic mean on the carbon-12 reference scale, so it has no unit. If an element contains 75% isotope of relative mass 35 and 25% isotope of relative mass 37, its mean is (75 × 35 + 25 × 37)/100 = 35.5. The mean does not imply that an individual atom has a fractional neutron. If the average mass and two isotope masses are given, let one fraction be x and the other 1 − x, then solve the weighted-average equation. The isotopic abundances belong to the stated sample; do not silently replace supplied numbers with a memorised natural mixture.

Applications and exam mistakes

Electron arrangements help predict simple ionic formulas and compare reactivity, while isotope data help identify tracers and mass-spectrometric peaks. Radioactive tracers require controlled handling; do not assume every isotope is suitable for a medical or industrial use merely because its mass differs. A mass spectrum plots measured ion abundance against mass-to-charge ratio, so charge state matters in detailed interpretation. In ordinary school problems with singly charged ions, peaks often correspond to isotope masses. Keep atomic mass, mass number, nucleon number and molar mass distinct. They are related concepts with different definitions and units, even when some numerical values resemble one another.

Self-check with explanations

An oxide ion O²⁻ with mass number 16 and atomic number 8 contains eight protons, eight neutrons and ten electrons. Its two extra electrons do not make it neon because element identity depends on protons, not total electron count. Magnesium and aluminium are in the same period, but their common ions Mg²⁺ and Al³⁺ can share the same electron count. An element whose neutral configuration ends in 3s² is magnesium in the first-twenty-element sequence and belongs to Period 3, Group 2. Finally, an average isotopic mass must lie between the contributing isotope masses; an answer outside that range signals an abundance or arithmetic error.