Why Periodic Trends Matter<\/h3>\n\n\n\n
Understanding periodic trends helps you predict and explain the chemical behavior<\/strong> of elements in the periodic table. These trends reflect changes in atomic structure<\/strong>, including nuclear charge<\/strong>, electron shielding<\/strong>, and the number of energy levels<\/strong>.<\/p>\n\n\n\n 1. Atomic Radius Across a Period (Left to Right)<\/strong> 2. Atomic Radius Down a Group (Top to Bottom)<\/strong> Note:<\/strong> There are exceptions due to subshell stability (e.g., Be vs. B, N vs. O).<\/p>\n\n\n\n Note:<\/strong> Fluorine is the most electronegative element; francium is among the least.<\/p>\n\n\n\nFlash Cards \u2013 Periodic Trend: Atomic Radius<\/h3>\n\n\n\n
Atomic radius decreases<\/strong> across a period.
As you move from left to right, each atom gains more protons<\/strong>, increasing the nuclear charge<\/strong>. This stronger pull draws electrons closer<\/strong> to the nucleus. Because electrons enter the same energy level<\/strong>, shielding remains mostly constant. As a result, atoms become smaller<\/strong> across the period.<\/p>\n\n\n\n
Atomic radius increases<\/strong> down a group.
Each element adds a new energy shell<\/strong>, which expands the electron cloud<\/strong>. Even though the nuclear charge also increases, the added inner electrons<\/strong> provide more shielding<\/strong>, which reduces the effective pull<\/strong> on the outer electrons. This causes the atoms to grow larger<\/strong> as you move down the group.<\/p>\n\n\n\n
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<\/a><\/figure>\n\n\n\n2. Periodic Trends-Ionization Energy (IE)<\/h4>\n\n\n\n
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Higher effective nuclear charge (Z_eff)<\/strong> means electrons are held more tightly, requiring more energy to remove<\/strong> the outermost electron.<\/li>\n\n\n\n
Outer electrons are farther from the nucleus and more shielded<\/strong>, so they are held less tightly and are easier to remove<\/strong>.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n3. Periodic Trends-Electron Affinity (EA)<\/h4>\n\n\n\n
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Atoms more readily accept electrons due to increasing Z_eff, releasing more energy when gaining an electron.<\/li>\n\n\n\n
Larger atoms have more distance between the nucleus and incoming electron, and increased shielding<\/strong> reduces the attraction, so less energy is released<\/strong>.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n4. Periodic Trends-Electronegativity (EN)<\/h4>\n\n\n\n
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Atoms more strongly attract shared electrons<\/strong> in covalent bonds, especially nonmetals like fluorine.<\/li>\n\n\n\n
As atomic radius increases, the nucleus is farther from bonding electrons<\/strong>, so the attraction weakens.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n5. Flash cards -Periodic Trends-Effective Nuclear Charge (Z_eff)<\/h4>\n\n\n\n
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More protons in the nucleus, but the same number of inner (shielding) electrons, so outer electrons feel a stronger pull<\/strong>.<\/li>\n\n\n\n
While nuclear charge increases, so does the number of inner shells, resulting in similar net attraction<\/strong> felt by valence electrons.<\/li>\n<\/ul>\n\n\n\n
\n\n\n\n\ud83d\udd38 Flash Cards- Periodic Trends-Application of Trends:<\/h4>\n\n\n\n
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