Periodic Table, Periodic Properties and Variations of Properties — Question 16
Back to all questionsBoth Q and R
Reason — Element X with an atomic number of 13 is Aluminium (Al), which is a metal. Element Y with an atomic number of 17 is Chlorine (Cl), which is a non-metal.
X being a metal tends to lose electrons to achieve a stable electron configuration and Y being a non-metal tends to gain electrons to achieve stable electron configuration. Hence, statements Q and R are true whereas statement P is false.
Chapter Overview: The Periodic Table
The Periodic Table is the systematic arrangement of elements in order of increasing atomic number. Mendeleev arranged elements by atomic mass and predicted undiscovered elements, but the Modern Periodic Table (by Moseley) uses atomic number as the basis. The table has 7 periods (horizontal rows) and 18 groups (vertical columns). Elements in the same group share similar chemical properties because they have the same number of valence electrons. Periodicity refers to the recurrence of similar properties at regular intervals. Key periodic properties include atomic size, ionisation energy, electron affinity, electronegativity, and metallic/non-metallic character. Understanding trends across periods (left to right) and down groups is essential for predicting element behaviour. The table is divided into s-block, p-block, d-block (transition metals), and f-block (inner transition metals). ICSE Class X focuses on the first 20 elements (H to Ca), their electron configurations, and the relationship between electronic configuration and position in the table. Students must know how to locate an element given its atomic number and predict its properties from its position.
Key Definitions & Concepts
| Term | Definition |
|---|---|
| Atomic Number (Z) | Number of protons in the nucleus of an atom |
| Period | Horizontal row; period number = number of electron shells |
| Group | Vertical column; elements share the same number of valence electrons |
| Periodicity | Recurrence of similar properties at regular intervals of atomic number |
| Atomic Size | Increases down a group, decreases across a period (left to right) |
| Ionisation Energy | Energy required to remove the outermost electron from an isolated gaseous atom |
| Electron Affinity | Energy released when an electron is added to a neutral gaseous atom |
| Electronegativity | Tendency of an atom to attract shared pair of electrons towards itself |
| Metallic Character | Increases down a group, decreases across a period |
| Modern Periodic Law | Properties of elements are a periodic function of their atomic numbers |
Must-Know Concepts
- Elements in Group 1 (alkali metals) have 1 valence electron; Group 17 (halogens) have 7 valence electrons
- Noble gases (Group 18) have complete octets and are chemically inert
- Period number = number of occupied electron shells
- Group number for main group elements: Group = valence electrons (for groups 1-2); Group = 10 + valence electrons (for groups 13-18)
- Across a period: atomic size decreases, ionisation energy increases, electronegativity increases, metallic character decreases
- Down a group: atomic size increases, ionisation energy decreases, electronegativity decreases, metallic character increases
- Mendeleev's table was arranged by atomic mass; Modern table by atomic number
Mendeleev's vs Modern Periodic Table
| Feature | Mendeleev's Table | Modern Table |
|---|---|---|
| Basis | Atomic mass | Atomic number |
| Groups | 8 groups (I to VIII) | 18 groups (1 to 18) |
| Periods | 7 periods (some incomplete) | 7 periods (complete) |
| Isotopes | No place for isotopes | Same place (same Z) |
| Noble Gases | Not included initially | Group 18 |
Important Diagrams to Practice
- First 20 elements with electron configurations in the periodic table layout
- Trend arrows showing variation of atomic size, IE, EA, and electronegativity across a period and down a group
- Electron dot structures for elements of Period 2 and Period 3
Common Mistakes
- Confusing atomic number with mass number when placing elements
- Saying ionisation energy increases down a group (it decreases)
- Writing wrong electron configuration for elements like K (2,8,8,1 not 2,8,9)
- Forgetting that noble gases have the highest ionisation energy in each period
- Mixing up electronegativity trend (increases across period, decreases down group)
Scoring Tips
- Always write electron configuration first when asked to identify period and group
- For trend questions, state the trend AND give the reason (nuclear charge, shielding effect)
- Remember: Period = number of shells, Group = valence electrons
- Practice placing elements 1-20 rapidly; this appears in almost every exam
Frequently Asked Questions
Why does atomic size decrease across a period?
As we move across a period, the nuclear charge (number of protons) increases while electrons are added to the same shell. The increased nuclear charge pulls the electron cloud closer, reducing atomic size.
Why are noble gases chemically inert?
Noble gases have completely filled outermost shells (duplet for He, octet for others), making them extremely stable with no tendency to gain, lose, or share electrons.
What is the difference between electron affinity and electronegativity?
Electron affinity is the energy change when an isolated gaseous atom gains an electron. Electronegativity is the tendency of a bonded atom to attract shared electrons. EA is a measurable quantity; electronegativity is a relative scale.
How did Mendeleev predict undiscovered elements?
Mendeleev left gaps in his table for elements not yet discovered and predicted their properties (atomic mass, density, oxide formulae) based on the properties of surrounding elements. His predictions for eka-aluminium (Gallium) and eka-silicon (Germanium) proved remarkably accurate.