IAL Chemistry Periodic Trends 2025 - 2026

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Hosni chembio

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Created by

Hosni chembio

teacher

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What is electron shielding?

The blocking of valence shell electron attraction by the nucleus due to the presence of inner-shell electrons.

How does electron shielding change as you move down a group in the Periodic Table?

Electron shielding increases down a group due to the increase in the number of quantum shells.

What is nuclear charge?

The total positive charge of all the protons in the nucleus.

Define effective nuclear charge.

The actual nuclear charge minus the effects of shielding due to inner-shell electrons.

How does atomic radius change as you move down a group?

Atomic radius increases as you go down a group due to the extra number of energy levels.

What happens to atomic radius as you move across a period?

The atomic radius decreases across a period.

What is the ionic radius of cations compared to their original atoms?

Cations have a smaller radius than their original atom.

How does the ionic radius of anions compare to their original atoms?

Anions have a larger radius than their original atom.

What determines the radii of isoelectronic species?

The radii of isoelectronic species are determined by their proton number.

What is the first ionisation energy?

The energy required to remove an electron from a neutral atom in the gaseous state.

How does first ionisation energy change as you go down a group?

The first ionisation energy decreases as you go down a group.

What trend is observed in first ionisation energy as you move across a period?

The first ionisation energy increases as you go across a period.

What are two exceptions to the trend of first ionisation energy across a period?

A slight decrease occurs when moving from group 2 to group 3 and between group 5 and group 6.

What is the second ionisation energy?

The energy required to remove an electron from a 1+ cation in the gaseous state.

How does melting point trend among metallic elements like sodium, magnesium, and aluminium?

There is a gradual increase in melting points among sodium, magnesium, and aluminium.

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Boost your knowledge before diving into past papers. Use these questions to reinforce your understanding of the lessons and supercharge your revision. They'll help you recall key concepts and stay on top of your studies

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1. What is electron shielding?

A The decrease in ionization energy down a group. B The blocking of valence shell electron attraction by the nucleus due to inner-shell electrons. C The total positive charge of all protons in the nucleus. D The increase in atomic radius across a period.

2. How does the effective nuclear charge change across a period?

A It decreases due to increased electron shielding. B It increases due to extra protons found in each element. C It remains constant due to unchanged nuclear charge. D It decreases due to added energy levels.

3. Why does atomic radius decrease across a period?

A Increase in electron shielding across a period. B Addition of more energy levels across a period. C Greater attractive force by increasing positive nuclear charge pulls electrons closer. D Decrease in nuclear charge across a period.

4. What happens to the ionic radius of cations compared to their original atoms?

A Cations have a larger radius due to increased energy levels. B Cations have a smaller radius due to greater attraction to the nucleus. C Cations have the same radius as their original atoms. D Cations have a larger radius due to electron gain.

5. Why does the first ionisation energy decrease down a group?

A Increased shielding and atomic radius outweigh increased nuclear charge. B Increased nuclear charge outweighs shielding effects. C Decreased nuclear charge down a group. D Decreased atomic radius down a group.

6. Why is there a decrease in the first ionisation energy when moving from group 2 to group 3?

A The 2p subshell in boron is further from the nucleus and more shielded than the 2s subshell in beryllium. B Boron has fewer protons than beryllium, reducing nuclear charge. C The 2p subshell in boron is closer to the nucleus than the 2s subshell in beryllium. D The 2s subshell in boron is closer to the nucleus than in beryllium.

7. What causes the exception in first ionisation energy between Group 5 and Group 6?

A Nitrogen's p electrons are paired, increasing ionisation energy. B Nitrogen has a higher nuclear charge than oxygen. C Spin pairing in oxygen causes repulsion, reducing ionisation energy compared to nitrogen. D Oxygen has a full p subshell, increasing ionisation energy.

8. Why do Group 1 elements have the highest second ionisation energy (IE2)?

A They have more protons than other groups, increasing nuclear charge. B They have a full outer shell, making electron removal difficult. C Removing an electron from a lower s subshell requires more energy. D Their electrons are in higher energy levels, requiring more energy to remove.

9. What factor contributes to the increase in melting points of metals like sodium, magnesium, and aluminium?

A Lower cation charge intensity reduces melting points. B Decreased number of valence electrons weakens bonds. C Weaker metallic bonds due to fewer delocalised electrons. D Stronger metallic bonds due to more delocalised electrons and greater cation charge.

10. How does electron configuration affect the second ionisation energy trend across a period?

A Cations with stable electron configurations have higher second ionisation energies. B Cations with half-filled p subshells have lower ionisation energies. C Electron configuration has no effect on second ionisation energy. D Cations with unstable configurations have higher ionisation energies.

11. What is the reason for the low melting points of elements from group 5 onwards?

A Strong covalent bonds between atoms. B High charge density and small ionic radius. C Giant molecular structure with strong bonds. D Weak intermolecular forces between simple molecules.

12. What factor contributes to the high melting point of aluminium?

A Weak metallic lattice and low charge density. B Large ionic radius and low charge density. C High charge density and more delocalised electrons. D Low charge density and fewer delocalised electrons.

13. Why do group 4 elements have the highest melting points among all elements?

A Due to their large atomic radius and low charge density. B Because they have the most delocalised electrons. C Because of their simple molecular structure with weak forces. D Due to their giant molecular structure with strong covalent bonds.

14. How does the reactivity of group 1 metals compare to group 2 metals?

A Both groups have similar reactivity levels. B Group 2 metals are more reactive due to higher ionisation energy. C Group 1 metals are less reactive due to higher ionisation energy. D Group 1 metals are more reactive due to lower ionisation energy.

15. Why is fluorine more reactive than other elements in the p block?

A Fluorine has a weaker nuclear charge than other p block elements. B Fluorine has fewer valence electrons than other p block elements. C Fluorine has a larger radius and more shielding. D Fluorine has a smaller radius and less shielding, leading to a higher effective nuclear charge.

Study Notes

Understanding Atomic Structure and Reactivity

This document explores key concepts related to atomic structure, including valence electrons, atomic radius, and the reactivity of metals and non-metals. It highlights how these factors influence chemical properties and behaviors across the periodic table.

Valence Electrons and Charge

Valence electrons are the outermost electrons that determine an atom's chemical properties. Metals typically have 1 to 3 valence electrons, leading to positive charges when they lose these electrons.
The reactivity of metals is primarily determined by their ability to lose electrons, with lower ionization energy correlating to higher reactivity.

Atomic Radius and Charge Density

The atomic radius decreases across a period (e.g., from sodium to aluminum), which increases charge density as smaller atoms have more concentrated charge.
A smaller atomic radius enhances the attraction between the nucleus and electrons, influencing reactivity and ionization energy.

Metallic Bonding and Melting Point

The strength of metallic bonds, influenced by the number of valence electrons, affects melting points. More valence electrons generally lead to stronger bonds and higher melting points.
Group 1 metals (alkali metals) are highly reactive due to their low ionization energy, while Group 2 metals are less reactive due to higher ionization energy.

Reactivity of Non-Metals

Non-metals, particularly in Group 7 (halogens), are highly reactive due to their tendency to gain electrons. Their reactivity is influenced by effective nuclear charge and atomic radius.
Fluorine, being the smallest halogen, has a high effective nuclear charge, making it the most reactive element in its group.

Key Takeaways

Valence electrons play a crucial role in determining an element's reactivity and charge.
Atomic radius and effective nuclear charge significantly influence ionization energy and electronegativity trends across the periodic table.
Understanding the properties of metals and non-metals, including their bonding characteristics and reactivity patterns, is essential for predicting chemical behavior.