Periodic Trends
Section: 8. The Periodic Table | Syllabus: Cambridge AS Level Physics 9702
Introduction to Periodic Trends Periodic trends are patterns in element properties that repeat across the Periodic Table. These trends exist because of the systematic way elements are arranged by atomic number and electron configuration.
Periodic Trend A predictable pattern in a property of elements that changes systematically as you move across periods or down groups in the Periodic Table. Why Do Trends Exist? Periodic trends occur due to three main factors: Nuclear charge - The number of protons in the nucleus Electron shells - The number of energy levels containing electrons Shielding effect - Inner electrons shield outer electrons from the full nuclear charge Electron configuration - The arrangement of electrons in shells Understanding Shielding Inner electron shells shield outer electrons from the full positive charge of the nucleus.
More inner shells = greater shielding = outer electrons feel less nuclear attraction. Atomic Radius Atomic radius is the distance from the nucleus to the outermost electron shell. It follows predictable patterns across the Periodic Table.
Atomic Radius The distance from the center of the nucleus to the outer edge of the electron cloud, typically measured in picometers (pm). Trend Across a Period (Left to Right) Atomic radius decreases as you move from left to right across a period.
Nuclear charge increases (more protons) Number of electron shells stays the same Shielding effect remains roughly constant Stronger nuclear pull on electrons draws them closer Atoms become smaller despite having more electrons Example: In Period 3, sodium (Na) is larger than chlorine (Cl) because Cl has a stronger nuclear charge pulling electrons closer.
Trend Down a Group (Top to Bottom) Atomic radius increases as you move down a group. Number of electron shells increases Outer electrons are further from the nucleus Increased shielding from inner electrons Nuclear charge increases but shielding effect is greater Atoms become larger moving down the group Example: In Group 1, cesium (Cs) is much larger than lithium (Li) because Cs has 6 electron shells compared to Li's 2 shells.
Quick Memory Trick Think "Down = Bigger, Across = Smaller". Elements get bigger going down and smaller going right. Diagram Periodic Table with arrows showing atomic radius increasing down groups and decreasing across periods, with size comparison circles for representative elements Ionization Energy Ionization energy is the energy required to remove an electron from a gaseous atom.
It measures how strongly an atom holds onto its electrons. First Ionization Energy The minimum energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.
Trend Across a Period (Left to Right) Ionization energy generally increases across a period. Atomic radius decreases Nuclear charge increases Electrons are held more tightly to the nucleus More energy needed to remove an electron Noble gases have the highest ionization energies in their periods Example: It takes much more energy to remove an electron from neon (Ne) than from sodium (Na).
Trend Down a Group (Top to Bottom) Ionization energy decreases down a group. Atomic radius increases Outer electrons are further from nucleus Greater shielding by inner electrons Electrons are held less tightly Less energy needed to remove an electron Example: Cesium (Cs) has a much lower ionization energy than lithium (Li) because its outer electron is much further from the nucleus.
Direction Ionization Energy Reason Across period → Increases Stronger nuclear pull, smaller atoms Down group ↓ Decreases Weaker pull due to distance and shielding Exceptions to the Trend There are small dips in ionization energy between Groups 2 and 13, and between Groups 15 and 16.
These occur due to electron pairing and subshell configurations, but the overall trend still holds. Diagram Graph showing first ionization energy vs atomic number for periods 2 and 3, displaying the general upward trend with small exceptions Electronegativity Electronegativity measures an atom's ability to attract electrons in a chemical bond.
It helps predict bond types and chemical behavior. Electronegativity A measure of the tendency of an atom to attract a bonding pair of electrons toward itself in a covalent bond. Trend Across a Period (Left to Right) Electronegativity increases across a period.
Atomic radius decreases Nuclear charge increases Atoms attract bonding electrons more strongly Fluorine is the most electronegative element Noble gases are not assigned electronegativity values (they don't bond) Trend Down a Group (Top to Bottom) Electronegativity decreases down a group.
Atomic radius increases Bonding electrons are further from nucleus Shielding effect reduces nuclear attraction Atoms attract bonding electrons less strongly Element Electronegativity (Pauling Scale) Group Fluorine (F) 4.0 Highest of all elements Oxygen (O) 3.5 Group 16 Chlorine (Cl) 3.0 Group 17 Sodium (Na) 0.9 G…
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