Grade 10-11 (age 15-17)

high school Periodic Table & Properties

The periodic table lists the elements in order of atomic number, folded so that elements with similar electron configurations line up in the same column. Elements in a group share the same number of valence electrons and behave alike, while across a period the atomic radius, ionization energy and electronegativity change in a regular way. All three trends come down to how strongly the nucleus holds the outermost electrons, that is, the effective nuclear charge and the number of electron shells. The small periodic table in the diagram colors the real values for elements 1 to 20, so switching the property brings out the trends across and down at once.

Why Is the Periodic Table Shaped Like This?

💡 Position Is Electron Configuration

①Elements in a group (column) have the same number of valence electrons; every group-1 element ends in ns¹ and is a reactive metal that readily loses one electron

②Elements in a period (row) have the same number of occupied shells; every period-3 element has three shells

③The table's shape follows the order in which orbitals fill: groups 1 and 2 fill s orbitals, groups 13 to 18 fill p orbitals

④Mendeleev left gaps and predicted the properties of elements not yet found, and gallium and germanium, discovered later, matched his predictions well

Periodic Properties of Elements 1 to 20

PropertyAtomic Radius

🔍 Comparing the Three Trends

①Atomic radius shrinks from left to right across a period: in period 3, Na at 186 pm is nearly twice Cl at 99 pm

②Down a group it grows, because each step adds an electron shell

③Ionization energy runs the other way, rising to the right and falling down a group; the dashed cells are exceptions to that flow

④Electronegativity is largest for F (4.0); group 18 hardly forms bonds, so it has no assigned value

The Root of All Three: Effective Nuclear Charge

Atomic Radius

across a period → smaller | down a group ↓ larger

Across a period the effective nuclear charge grows and pulls the electrons closer; down a group added shells make atoms larger

Ionization Energy

across a period → larger | down a group ↓ smaller

Energy needed to remove 1 mol of electrons from 1 mol of gaseous atoms

Electronegativity

across a period → larger | down a group ↓ smaller

How strongly an atom pulls on a shared electron pair in a covalent bond; F is highest at 4.0

💡 What Is Effective Nuclear Charge?

①The number of protons in the nucleus is the nuclear charge Z

②Inner electrons stand between the nucleus and the outer electrons and weaken its pull; this is called shielding

③The nuclear charge an outermost electron actually feels is the effective nuclear charge, which is smaller than Z

④Moving right across a period adds one proton at a time, but electrons in the same shell shield each other poorly, so the effective nuclear charge grows

⑤A larger effective nuclear charge holds electrons more tightly: the radius shrinks while ionization energy and electronegativity rise

Exceptions in Ionization Energy

📐 Group 2 > Group 13, Group 15 > Group 16

①Be (900) > B (801): the outermost electron of B is in 2p, which is higher in energy than 2s and easier to remove

②N (1402) > O (1314): oxygen's fourth 2p electron must pair up in an orbital that already holds an electron, and the repulsion makes it easier to remove

③The same happens in period 3: Mg (738) > Al (578) and P (1012) > S (1000). Choosing ionization energy in the diagram marks these eight cells with dashed outlines

④So ionization energy does not always rise across a period; it dips once after group 2 and once after group 15

Work It Out

Example 1

Across a period, how does atomic radius change as the atomic number increases, and why?

Moving right across a period adds protons, and electrons added to the same shell shield each other poorly.

same period: more protons → larger effective nuclear charge

The number of shells stays the same while the pull of the nucleus grows, so the electrons are drawn inward and the radius shrinks.

same number of shells + larger effective nuclear charge → smaller radius

▸ The atomic radius decreases.

In the diagram, the period-3 row turning bluer from Na at 186 pm to Cl at 99 pm is this trend.

Example 2

Which has the larger first ionization energy, ₁₁Na or ₁₇Cl? (Both are in period 3.)

Moving right across a period, the effective nuclear charge grows and electrons become harder to remove.

same period: ionization energy ↑ to the right

Cl lies to the right of Na, so its ionization energy is larger. The actual values are Na 496 and Cl 1251 kJ/mol.

Cl (1251) > Na (496)

▸ Cl

Losing one electron gives Na the same stable configuration as Ne, so it easily becomes Na⁺, while Cl tends to take an electron rather than give one up.

Summary

2023 KICE mock Chemistry I type, adapted

Which of these period-3 elements has the largest atomic radius?

①Na

②Mg

③Al

④Si

⑤Cl

▸ ① Na

Within a period, elements further left have a smaller effective nuclear charge and a larger radius.

Na is the leftmost (group 1) element of period 3, so its radius is the largest.

Na (186 pm) > Mg (160) > Al (143) > Si (117) > Cl (99)

🎯 Exam Points

①Atomic radius shrinks to the right across a period and grows down a group

②Ionization energy and electronegativity follow the opposite trend; watch for the group-2 and group-15 exceptions in ionization energy

③F has the highest electronegativity, followed by O

④Group 18 has no assigned electronegativity

⑤All three trends are explained by effective nuclear charge and the number of electron shells

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Electron Configuration

Isotopes & Atomic Mass

Last updated 2026-10-17

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