Periodic Table & Properties

The periodic table is a map of the elements based on electron configuration. Elements in the same group share valence electrons and similar chemistry; across a period, atomic radius, ionization energy, and electronegativity shift with effective nuclear charge. When Z_eff rises, electrons are held tighter, so radius falls while ionization energy and electronegativity rise. Here you pick a property and compare its periodic trend on a mini table.

Why Does the Periodic Table Look Like This?

💡 Key: Electron Configuration Determines All Properties

①Same group (column): same valence electrons → similar chemistry

②Same period (row): same number of shells → similar atomic size

③Mendeleev left blanks and predicted properties of 'undiscovered' elements

④Those predictions came true — the genius of the periodic table

⑤A 'map of elements' — knowing the position predicts properties

Periodic Properties of Elements

Property (0=radius, 1=IE, 2=electronegativity)0

🔍 Compare the Three Trends!

①Atomic radius: decreases across a period (Z_eff↑)

②Ionization energy: increases across a period

③Electronegativity: F(4.0) is max — noble gases not defined

④All three are explained by effective nuclear charge (Z_eff)

Effective Nuclear Charge — the Underlying Cause

Atomic Radius

Period: → decreases | Group: ↓ increases

Z_eff ↑ pulls electrons more strongly → radius ↓

Ionization Energy

Period: → increases | Group: ↓ decreases

Minimum energy to remove an electron

Electronegativity

Period: → increases | Group: ↓ decreases

Tendency to attract bonding electrons (F=4.0 max)

💡 What is Effective Nuclear Charge?

①Number of protons = nuclear charge Z

②Inner electrons screen the nuclear charge — shielding effect

③Z_eff = Z − screening

④Across a period: Z↑, shielding ~constant → Z_eff↑

⑤Z_eff↑ → tighter pull → radius↓, IE↑, EN↑

Exceptions in Ionization Energy

📐 Group 2 → 13 and Group 15 → 16 Exceptions

①Be(900) > B(801): jumping from 2s² to 2p¹ costs less energy

②N(1402) > O(1314): half-filled 2p³ is extra stable, 2p⁴ has more repulsion

③These exceptions reveal the stability of half/full subshells

④In exams: ionization energy is not strictly increasing across a period

⑤Be ready to explain why group 2/15 IE exceeds the next group

Work It Out

Example 1

Within the same period, how does atomic radius change as atomic number increases? Why?

Across a period the proton count rises, so the nucleus pulls electrons more strongly.

Same period: protons ↑ → effective nuclear charge ↑

The number of electron shells stays the same, so the radius shrinks.

Same shell count → radius decreases

▸ Atomic radius decreases

A larger effective nuclear charge pulls the same-shell electrons in harder, shrinking the radius.

Example 2

Between ₁₁Na and ₁₇Cl, which has the larger first ionization energy? (both period 3)

Moving right across a period, effective nuclear charge grows, so ionization energy rises.

Same period: ionization energy ↑ to the right

Cl(17) sits to the right of Na(11), so its ionization energy is larger.

Cl(17) > Na(11)

▸ Cl is larger

Cl holds its electrons more tightly, so removing one needs more energy.

Summary

2023 KICE mock Chemistry I type, adapted

Among these period-3 elements, which has the largest atomic radius?

①Na

②Mg

③Al

④Si

⑤Cl

▸ ① Na

Within a period, an element farther left has a smaller effective nuclear charge, so a larger radius.

Na is the leftmost period-3 element, so its radius is the largest.

Atomic number 11 (Na) is leftmost in period 3 → maximum radius

🎯 Exam Points

①Radius: decreases across period, increases down group

②IE: opposite trend to radius (watch exceptions!)

③Electronegativity: F > O > N > Cl (memorize!)

④Group 18 (noble) — electronegativity not defined

⑤Z_eff is the underlying cause: Z_eff↑ → radius↓, IE↑, EN↑

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Electron Configuration

Isotopes & Atomic Mass

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