Introduction to periodic table development and behaviour
Atomic structure and the periodic table • The periodic table
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Historical development of the periodic table
Early attempts order elements by atomic mass and note repeating properties. Recognition of periodic patterns leads to arrangements that group similar elements. A major advance occurs when gaps in the table predict the existence and properties of undiscovered elements, demonstrating the table's explanatory power. Further refinement follows the discovery that atomic number, not atomic mass, provides the correct ordering principle. Measurement of nuclear charge and the identification of the proton confirm ordering by atomic number and resolve previous anomalies in element placement.
Atomic number and periodic position
Atomic number equals the number of protons in the nucleus and increases by one for each step along the table. Position in the table follows increasing atomic number, so neighbouring elements differ by one proton and one electron in a neutral atom. Ordering by atomic number produces repeating patterns of properties (periodicity). Shell completion points (noble gases) and gradual changes across a period arise because atomic number controls nuclear charge and therefore the energy levels of electrons.
Electron arrangement, periods and groups
Periods correspond to the highest occupied principal electron shell. The period number equals the number of the outer electron shell that is filling. Groups contain elements with the same number of electrons in their outer shell (valence electrons), which cause similar chemical behaviour within a group. Valence-electron arrangements determine the types of ions an element forms: metals typically lose valence electrons to form positive ions; non-metals typically gain or share electrons to form negative ions or covalent bonds.
Relationship between electron arrangement and chemical reactions
Chemical reactions involve the gain, loss or sharing of valence electrons. Elements with one valence electron react to lose that electron readily and form +1 ions; elements with seven valence electrons react to gain one electron and form −1 ions. The ease of losing or gaining electrons depends on ionisation energy and electron affinity, which change with position on the table. Across a period, increasing nuclear charge (with similar shielding) pulls electrons closer, raising ionisation energy and reducing metallic character. Down a group, increased shell number increases shielding and atomic radius, lowering ionisation energy and often increasing metallic reactivity for metals while decreasing reactivity for non-metals.
Predicting reactions and relative reactivity from position
Group identity and period number enable prediction of likely reactions and products. Alkali metals (group 1) react with water to form hydroxides and hydrogen gas, with reactivity increasing down the group due to easier electron loss. Halogens (group 7) react with metals to form salts and show decreasing reactivity down the group because electron acceptance becomes harder. Transition metals show variable electron configurations and variable oxidation states, limiting simple predictions. Noble gases show very low reactivity due to full valence shells, with reactivity increasing slightly under extreme conditions or with heavier noble gases forming weak compounds in specific circumstances.
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