The periodic table is more than a list of elements — it’s a powerful tool that reveals predictable patterns, or periodic trends, in the physical and chemical properties of elements. These trends arise from changes in atomic structure as we move across periods or down groups.
In this guide, we focus on the periodic trends shown by s- and p-block elements, following the Chemistry syllabus, including atomic and ionic sizes, ionization energy, electron gain energy, and electronegativity.
Sizes of Atoms and Ions
Atomic Size (Atomic Radius)
Atomic size refers to the distance between the nucleus and the outermost electron shell.
There are three commonly discussed atomic radii:
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Van der Waals Radius: Half the distance between two non-bonded atoms in neighboring molecules.
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Covalent Radius: Half the distance between two nuclei in a covalent bond.
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Metallic Radius: Half the distance between nuclei of two adjacent atoms in a metallic lattice.
📌 Covalent radius is typically used when discussing nonmetals; metallic radius for metals.
Periodic Trends in Atomic Radii
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Across a Period (→): Atomic radius decreases
→ Due to increasing nuclear charge pulling electrons closer. -
Down a Group (↓): Atomic radius increases
→ Due to addition of new electron shells.
Electron Configurations of Ions
When an atom loses electrons to form a cation, the electrons are removed first from the orbitals with the highest principal quantum number (n).
For instance, in a sodium atom with the configuration 1s² 2s² 2p⁶ 3s¹, the electron is lost from the 3s orbital, as it has the highest value of n.
When atoms become ions:
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Cations (positive ions): Lose electrons → become smaller
→ Sometimes lose an entire outer shell. -
Anions (negative ions): Gain electrons → become larger
→ Increased electron-electron repulsion expands the electron cloud.
Example:
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Na (Z = 11): 1s² 2s² 2p⁶ 3s¹
→ Na⁺ = 1s² 2s² 2p⁶ (smaller size) -
Cl (Z = 17): 1s² 2s² 2p⁶ 3s² 3p⁵
→ Cl⁻ = 1s² 2s² 2p⁶ 3s² 3p⁶ (larger size)
Periodic Trends in Ionic Radii
The radii of cations and anions, compared to their respective parent atoms (in picometres, pm), are shown in the figure below.
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Cationic radii are smaller than atomic radii.
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Anionic radii are larger than atomic radii.
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Across a period (for isoelectronic species): Ionic radii decrease with increasing nuclear charge.
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Down a group: Ionic radii increase due to added shells.
- An isoelectronic series refers to a group of atoms or ions that all have the same number of electrons.
For example, the species O²⁻, F⁻, Ne, Na⁺, and Mg²⁺ each possess 10 electrons. Within such a series, as the atomic number increases, the nuclear charge also increases. Since the electron count stays the same, a stronger attraction is exerted by the nucleus on the electrons, causing the ionic radius to decrease progressively across the series.
Ionization Energy
Ionization energy (IE) is the energy required to remove one mole of electrons from one mole of gaseous atoms.
- The first ionization energy (I₁) is the energy needed to remove the outermost, most weakly held electron from a neutral atom in the gas phase. For instance, in the case of lithium, the first ionization involves the reaction:
Li(g) → Li⁺(g) + e⁻
- The second ionization energy (I₂) is the energy required to remove a second electron — this time from a positively charged ion (monovalent cation) to form a doubly charged ion (divalent cation). For lithium, this process is:
Li⁺(g) → Li²⁺(g) + e⁻
- Each successive ionization energy is higher than the previous one, as it’s increasingly difficult to remove electrons from an already positively charged ion.
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Ionization energy increases with each successive electron removed:
I₁ < I₂ < I₃, and so on. -
Reason for the increase:
Each electron is removed from a more positively charged ion, requiring more energy to overcome the stronger attraction to the nucleus. -
Sharp jump in ionization energy:
A significant increase occurs when an inner-shell electron is removed after all outer-shell electrons have been taken. -
Why inner electrons require more energy:
Inner-shell electrons are closer to the nucleus and experience a stronger electrostatic pull, making them harder to remove.
Periodic Trends in First Ionization Energies
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Across a Period: IE increases
→ Greater nuclear attraction and smaller size make it harder to remove electrons. -
Down a Group: IE decreases
→ Outer electrons are further from the nucleus and shielded by inner shells.
Detail diagrams of first ionization energy variation are given below.
Variation of First Ionization Energies Across Periods 1 and 2
The diagram shows how the first ionization energy changes as we move left to right across the first (H to He) and second (Li to Ne) periods of the periodic table.
Key Trends and Reasons:
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General Increase Across a Period:
Ionization energy generally increases from left to right across both periods.
🔹 This is because the nuclear charge increases (more protons in the nucleus), pulling electrons closer and more tightly, which makes them harder to remove. -
Same Shell, Increasing Attraction:
Even though all elements in a period have electrons in the same principal energy level, the increasing nuclear charge leads to a stronger attraction, so more energy is required to remove an electron. -
Notable Exceptions:
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Between Be and B, and N and O, there is a slight drop in ionization energy.
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Boron vs. Beryllium: The 2p electron in boron is slightly easier to remove than the 2s electron in beryllium due to higher energy and shielding.
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Oxygen vs. Nitrogen: In oxygen, electron pairing in the 2p orbital causes electron–electron repulsion, making it easier to remove one electron compared to nitrogen.
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This pattern highlights how atomic structure, particularly nuclear charge and electron configuration, influences the energy needed to remove electrons, helping us understand chemical reactivity and periodic trends.
Note: Small dips occur due to electron pairing and subshell stability (e.g., Be vs. B, N vs. O).
Electron Gain Energy (Electron Affinity)
Electron gain energy refers to the energy change that occurs when an electron is added to a neutral atom in the gaseous state. For most elements, this process releases energy, meaning the value is negative.
For instance, when a chlorine atom gains an electron, it releases -349 kJ mol⁻¹, indicating that the process is energetically favorable:
Cl(g) + e⁻ → Cl⁻(g) ΔE EG = –349 kJ mol⁻¹
However, some atoms have positive electron gain energy values, such as beryllium (Be) and nitrogen (N).
This happens because:
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Be has a stable 2s² configuration.
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N has a half-filled 2p³ subshell.
In both cases, adding another electron would cause electron–electron repulsion, making the process energetically unfavorable.
Examples:
N(g) + e⁻ → N⁻(g) ΔE EG = +134 kJ mol⁻¹
Be(g) + e⁻ → Be⁻(g) ΔE EG = +231 kJ mol⁻¹
Periodic Trends:
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Across a period: Electron gain energy becomes less positive (or more negative), meaning atoms are more likely to gain electrons.
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Down a group: It becomes more positive, indicating decreased tendency to accept electrons due to increased atomic size and shielding.
Relationship Between Electron Gain Energy and Electron Affinity
By international convention:
Electron gain energy (ΔE EG) = – Electron affinity (EA)
So, electron affinity represents the energy released when an electron is added to a gaseous atom, and it is numerically equal but opposite in sign to ΔE EG.
For example:
A⁻(g) → A(g) + e⁻ ΔE = EA,
which is the reverse of the electron gain process.
In summary, electron affinity is a key property that reflects how strongly an atom attracts an added electron and it closely mirrors the trends and behavior observed in electron gain energy.
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If energy is released, the value is negative → Atom strongly attracts electrons.
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If energy is absorbed, the value is positive → Atom resists gaining electrons.
Periodic Trends
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Across a Period: Electron gain energy becomes more negative (more exothermic).
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Down a Group: Becomes less negative → Larger atoms attract extra electrons less strongly.
📌 Halogens have the most negative electron affinities.
📌 Noble gases have positive electron gain energy — they don’t readily gain electrons.
Electronegativity
Electronegativity is the ability of an atom in a chemical bond to attract bonding electrons.
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Measured using Pauling scale (F = 3.98 is the highest).
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Not an actual energy value, but a relative scale.
The Role of Electronegativity in Chemical Bonding
Electronegativity is a key concept in chemistry—it helps us understand how atoms interact when they bond. By looking at the difference in electronegativity between two atoms, we can predict the type of bond they will form:
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Nonpolar covalent bond: Formed when the electronegativity difference is less than 0.4. Electrons are shared almost equally.
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Polar covalent bond: Occurs when the difference falls between 0.4 and 1.7. Electrons are shared unequally, creating partial charges on atoms.
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Ionic bond: Forms when the difference is greater than 1.7. One atom effectively takes an electron from the other.
Understanding this helps us predict a compound’s properties like polarity, solubility, and reactivity.
Electronegativity of Fluorine, Oxygen, and Chlorine
Among all elements, fluorine tops the electronegativity chart—it has the strongest tendency to attract electrons in a bond. Next comes oxygen, followed by chlorine.
What does this mean in practice?
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Fluorine is almost always assigned a negative oxidation state, no matter what it’s bonded to. Its extreme electronegativity makes it highly reactive—it aggressively pulls electrons toward itself.
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Oxygen also tends to be negative in compounds, except in rare cases like oxygen-fluorine compounds, where fluorine’s stronger pull takes precedence.
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Chlorine, while slightly less electronegative, still shows strong electron-attracting behavior, contributing to its reactivity in halogen reactions.
Why Noble Gases Don’t Have Electronegativity Values
You won’t find electronegativity values for noble gases (like helium, neon, or argon) on most charts. That’s because these elements don’t usually form chemical bonds—they already have full outer electron shells and are generally inert. Since electronegativity is about how strongly an atom attracts electrons in a bond, it doesn’t apply well to noble gases under normal conditions.
How to Determine the Electronegativity of a Molecule
To estimate the electronegativity behavior in a molecule, examine each individual bond:
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Identify the two atoms involved in the bond.
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Look up their electronegativity values (typically found on a periodic table).
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Calculate the difference between the two.
The magnitude of this difference tells you how polar that bond is:
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Smaller difference → more equal sharing → nonpolar
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Larger difference → unequal sharing → polar or even ionic
This approach is especially useful in predicting molecular geometry, polarity, and how the molecule will interact with others (e.g., in solvents or reactions).
Periodic Trends
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Across a Period: Electronegativity increases
→ Due to smaller size and higher nuclear charge. -
Down a Group: Electronegativity decreases
→ Atoms are larger and exert weaker pull on bonding electrons.
Summary Table of Periodic Trends (s- and p-Block Elements)
Final Tips for Students
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Always relate trends to atomic structure: shell number, nuclear charge, and shielding.
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Practice comparing elements in same period or group.
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Be familiar with exceptions and reasons behind anomalies (like subshell stability).
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Link these trends to chemical properties such as reactivity, bonding, and compound formation.