8.1 Modern Periodic Table
The modern periodic table arranges elements in order of increasing atomic number. When elements are arranged horizontally by atomic number, their properties repeat at regular intervals (periodicity). Elements with similar properties are placed in vertical columns (groups).
KEY NOTES: Modern Periodic Table
BASIS OF MODERN PERIODIC TABLE
• Atomic number (proton number) determines position
• Earlier: Mendeleev used atomic mass
• Henry Moseley established atomic number as basis
• Periodic law: Properties repeat periodically with atomic number
PERIODS (HORIZONTAL ROWS)
• 7 periods total
• Show gradual change in properties
• Each period = completion of an electron shell
• Period number = Principal quantum number (n)
GROUPS (VERTICAL COLUMNS)
• 18 groups total
• Elements have same valence electrons
• Show similar chemical properties
• Also called families
Memory Trick:
“Periods are ROWS (like rows in theater), Groups are COLUMNS (like columns in temple)”
Remember: Modern periodic table = Atomic number basis (not atomic mass).
8.2 Salient Features of Modern Periodic Table
Table 8.1: Periods in Periodic Table
| Period No. | Name | Elements | Shell Being Filled | Notes |
|---|---|---|---|---|
| 1 | Very Short | 2 | 1st (K) | H and He only |
| 2 | Short | 8 | 2nd (L) | Li to Ne |
| 3 | Short | 8 | 3rd (M) | Na to Ar |
| 4 | Long | 18 | 4th (N) | K to Kr, includes transition metals |
| 5 | Long | 18 | 5th (O) | Rb to Xe |
| 6 | Very Long | 32 | 6th (P) | Cs to Rn, includes lanthanides |
| 7 | Very Long | 32 | 7th (Q) | Fr to Og, includes actinides |
Interesting Information!
Mendeleev vs Moseley: Mendeleev arranged elements by atomic mass, Moseley by atomic number. Moseley’s work established the modern basis of periodic law.
Periodic Table Variations: E.G. Mazurs collected 700 different published versions! Some were spirals, circles, triangles, but rectangular structure remains most common.
s-BLOCK
Groups 1-2
Alkali & Alkaline Earth Metals
p-BLOCK
Groups 13-18
Metalloids & Non-metals
d-BLOCK
Groups 3-12
Transition Metals
f-BLOCK
Lanthanides & Actinides
Bottom two rows
KEY NOTES: Table 8.2 – Groups & Electronic Configuration
GROUP 1: ALKALI METALS
• Configuration: ns¹
• 1 valence electron
• Form 1+ ions
• Very reactive, soft metals
GROUP 2: ALKALINE EARTH METALS
• Configuration: ns²
• 2 valence electrons
• Form 2+ ions
• Less reactive than group 1
GROUPS 3-12: TRANSITION METALS
• Configuration: ns² (n-1)d¹⁻¹⁰
• Inner d-subshell filling
• Variable oxidation states
• Form colored compounds
GROUPS 13-18: p-BLOCK
• Configuration: ns² np¹⁻⁶
• Group 13: ns² np¹ (Boron family)
• Group 14: ns² np² (Carbon family)
• Group 18: ns² np⁶ (Noble gases)
Solved Exercise from Textbook
1. Element with electronic configuration s²p¹:
• Period: 2 (n=2 for s and p orbitals)
• Group: 13 (2+1=3 valence electrons)
• Block: p-block (p orbital filling)
• Example: Boron (B)
2. Element with eight electrons in outermost shell:
• Group: 18 (Noble gases: ns² np⁶)
• Physical state: Gas (all noble gases are gases)
• Example: Neon (Ne), Argon (Ar)
3. Group 16 element that is a gas:
• Oxygen (O) or Sulfur (S) but Oxygen is gas
• p-block element
8.3 Similarities in Chemical Properties of Elements in Same Group
Elements in same group have same valence electrons → similar chemical properties. This is the fundamental basis for periodic classification.
KEY NOTES: Group-wise Chemical Properties
GROUP 1: ALKALI METALS
• Configuration: ns¹ → lose 1 electron → M⁺
• Highly electropositive
• React with water: 2M + 2H₂O → 2MOH + H₂↑
• React with halogens: 2M + X₂ → 2MX
• Reactivity increases down group
GROUP 2: ALKALINE EARTH METALS
• Configuration: ns² → lose 2 electrons → M²⁺
• Less reactive than Group 1
• Basic oxides: MO + H₂O → M(OH)₂
• Reactivity increases down group
GROUP 17: HALOGENS
• Configuration: ns² np⁵ → gain 1 electron → X⁻
• Highly electronegative
• React with metals: 2M + X₂ → 2MX
• Reactivity decreases down group
• “Halogen” = salt former
GROUP 16: OXYGEN FAMILY
• Configuration: ns² np⁴ → gain 2 electrons → X²⁻
• Oxygen: strong electronegative element
• O₂ + 2e⁻ → O²⁻
• Form oxides with metals
Relationship: Group Number ↔ Ion Charge
For Main Group Elements:
• Left to right: Charge on cation = Group number
• Group 1 → 1+ (Na⁺, K⁺)
• Group 2 → 2+ (Mg²⁺, Ca²⁺)
• Group 3 → 3+ (Al³⁺)
Right to left: Charge on anion = 18 – Group number
• Group 17 (18-17=1) → 1- (Cl⁻, F⁻)
• Group 16 (18-16=2) → 2- (O²⁻, S²⁻)
• Group 15 (18-15=3) → 3- (N³⁻, P³⁻)
Memory Trick:
“Group 1 = lose 1 e⁻, Group 2 = lose 2 e⁻, Group 17 = gain 1 e⁻, Group 16 = gain 2 e⁻”
For cation charge: Group number = Charge
For anion charge: 18 – Group number = Charge
8.4 Variation of Periodic Properties in Periods and Groups
Periodic properties show predictable trends across periods and down groups due to recurring electronic configurations. Key properties include atomic radius, ionization energy, electron affinity, and electronegativity.
Atomic Radius
Period: Decreases →
Group: Increases ↓
Ionization Energy
Period: Increases →
Group: Decreases ↓
Electron Affinity
Period: Increases →
Group: Decreases ↓
Electronegativity
Period: Increases →
Group: Decreases ↓
KEY NOTES: Periodic Properties
ATOMIC RADIUS
• Definition: Half distance between nuclei of bonded identical atoms
• Unit: pm (picometers, 1pm = 10⁻¹²m)
• Period trend: Decreases (increased nuclear charge pulls electrons)
• Group trend: Increases (new shells added, shielding effect)
• Example: Li (152pm) to Ne (69pm) in period 2
IONIZATION ENERGY
• Definition: Energy to remove most loosely bound electron
• Unit: kJ/mol
• First IE: M(g) → M⁺(g) + e⁻
• Period trend: Increases (harder to remove electrons)
• Group trend: Decreases (easier to remove electrons)
• Example: Li (520) to Ne (2081) kJ/mol in period 2
ELECTRON AFFINITY
• Definition: Energy change when electron added to gaseous atom
• Usually negative (energy released)
• Period trend: Increases (more negative)
• Group trend: Decreases (less negative)
• Example: F (-328) to I (-295) kJ/mol in group 17
ELECTRONEGATIVITY
• Definition: Ability to attract shared electrons
• Pauling scale (F = 4.0 highest)
• Period trend: Increases
• Group trend: Decreases
• Most electronegative: F, O, N, Cl (top-right corner)
Table 8.4: Atomic Radii – Period 2 Elements
| Element | Li | Be | B | C | N | O | F | Ne |
|---|---|---|---|---|---|---|---|---|
| Atomic Radius (pm) | 152 | 113 | 88 | 77 | 75 | 73 | 71 | 69 |
Observation: Atomic radius decreases across period due to increasing nuclear charge pulling electrons closer.
Table 8.5: Atomic Radii – Group 1 Elements
| Element | Li | Na | K | Rb | Cs |
|---|---|---|---|---|---|
| Atomic Radius (pm) | 152 | 186 | 227 | 248 | 265 |
| Inner Shell e⁻ | 2 | 10 | 18 | 36 | 54 |
Observation: Atomic radius increases down group due to addition of new shells and shielding effect.
Interesting Information!
Atomic size doesn’t always increase with atomic number! Size depends on electron shell diameter, not just number of protons. For example, Na (Z=11) is larger than Mg (Z=12) because Mg has stronger nuclear pull on same shell.
Solved Exercises with Explanations
(i) In which period and group will you place the element which is an important part of the solar cell?
Show ExplanationExplanation: Silicon (Si) is used in solar cells. Electronic configuration: 1s²2s²2p⁶3s²3p² or [Ne]3s²3p².
• Period: 3 (n=3 for valence electrons)
• Group: 14 (2+2=4 valence electrons: ns²np²)
• Silicon is in 3rd period, 14th group (Carbon family).
(ii) Identify the electronic configuration of the outermost shell of a transition metal.
Show ExplanationExplanation: Transition metals (d-block elements, groups 3-12) have general configuration: [Noble gas] (n-1)d¹⁻¹⁰ ns¹⁻².
• The outermost shell typically has ns² configuration
• Inner d-orbitals are being filled
• Example: Fe: [Ar] 3d⁶ 4s² (outermost: 4s²)
• nd¹⁰ns² represents complete d-subshell + s² (e.g., Zn: [Ar] 3d¹⁰ 4s²)
(iii) Which is the softest metal?
Show ExplanationExplanation: Among given options, Sodium (Na) is the softest.
• Alkali metals (Group 1) are very soft, can be cut with knife
• Softness order: Na (very soft) > K > Rb > Cs (softest)
• Ca (Group 2) is harder than Na
• Al and Zn are much harder metals
Reason: Large atomic size + weak metallic bonding in alkali metals.
(iv) A yellow solid element exists in allotropic forms which is also present in fossil fuels. Indicate the name.
Show ExplanationExplanation: Sulphur (S) fits all descriptions:
1. Yellow solid at room temperature
2. Allotropic forms: Rhombic, monoclinic, plastic sulfur
3. Present in fossil fuels: Coal and petroleum contain sulphur compounds
4. Group 16, period 3 element
Carbon is black/gray, iodine is purple-black solid, aluminum is silvery-white.
(v) How many electrons can nitrogen accept in its outermost shell?
Show ExplanationExplanation: Nitrogen (N) electronic configuration: 1s²2s²2p³ or [He]2s²2p³.
• Valence shell: n=2
• Valence electrons: 5 (2s²2p³)
• Needs 3 electrons to complete octet (8 electrons total)
• Can form N³⁻ ion by gaining 3 electrons
• Examples: NH₃ (shares 3 electrons), Na₃N (ionic with N³⁻)
(vi) Which element is the most reactive element?
Show ExplanationExplanation: Fluorine (F) is the most reactive element.
• Highest electronegativity (4.0 on Pauling scale)
• Smallest atomic size in period 2
• Strongest tendency to gain electron
• Reacts with almost all elements (except some noble gases)
• Even reacts with glass, gold, platinum!
• Among non-metals: Reactivity decreases down group 17 (F > Cl > Br > I)
(vii) Which element has the highest melting point?
Show ExplanationExplanation: Among alkali metals (Group 1), melting point decreases down the group:
• Li: 180°C (highest in group 1)
• Na: 98°C
• K: 63°C
• Rb: 39°C
• Cs: 28°C (lowest, melts in hand!)
Reason: As size increases, metallic bonding weakens → lower melting point.
But note: Cs has the lowest melting point, not highest! The question seems misleading – among given options, Na has highest MP.
(x) The element having less value of ionization energy and less value of electron affinity is likely to belong to:
Show ExplanationExplanation: Group 1 elements (alkali metals) have:
1. Low ionization energy: Easy to remove outer electron (ns¹ configuration)
2. Low electron affinity: Don’t tend to gain electrons (positive EA values)
• Group 1: Very low IE, positive/very low EA
• Group 17: High IE (hard to remove e⁻), high negative EA (tend to gain e⁻)
• Group 13: Moderate IE, some have positive EA
• Group 16: Moderate-high IE, moderate negative EA
i. Why was atomic number chosen to arrange the elements in the periodic table?
Answer:
Atomic number (Z = number of protons) was chosen because:
1. Fundamental property: Atomic number determines chemical identity of element
2. Moseley’s discovery (1913): Showed that properties correlate better with atomic number than atomic mass
3. Resolves anomalies: Fixes problems in Mendeleev’s table (e.g., Ar-K, Co-Ni, Te-I positions)
4. Electronic configuration: Atomic number = number of electrons in neutral atom → determines electron configuration → determines chemical properties
5. Periodic law: “Properties are periodic function of atomic number” (modern periodic law)
ii. What is the significance of the word periodic?
Answer:
The word “periodic” means “recurring at regular intervals.” In context of periodic table:
1. Periodic repetition: Properties of elements repeat after regular intervals when arranged by atomic number
2. Periodic law: “Properties of elements are periodic functions of their atomic numbers”
3. Periods: Horizontal rows where properties change gradually
4. Periodicity: Recurrence of similar electronic configurations at regular intervals
Example: Li, Na, K all have ns¹ configuration and similar properties, occurring at atomic numbers 3, 11, 19.
iii. Why does the size of a period increase as we move down the periodic table?
Answer:
The size (number of elements) in periods increases down the table because:
1. Increasing electron capacity: Higher shells can accommodate more electrons
• Period 1 (n=1): 1s subshell holds 2 electrons → 2 elements
• Period 2 (n=2): 2s²2p⁶ holds 8 electrons → 8 elements
• Period 3 (n=3): 3s²3p⁶ holds 8 electrons → 8 elements
• Period 4 (n=4): 4s²3d¹⁰4p⁶ holds 18 electrons → 18 elements
• Period 6 (n=6): 6s²4f¹⁴5d¹⁰6p⁶ holds 32 electrons → 32 elements
2. Additional subshells: d and f subshells appear in higher periods
3. Electron filling order: According to Aufbau principle and energy levels
iv. In a group, the elements have the same number of electrons in the outermost shell. Why is it so?
Answer:
Elements in same group have same valence electrons because:
1. Group number definition: For main groups (1,2,13-18), group number = number of valence electrons
• Group 1: ns¹ (1 valence electron)
• Group 2: ns² (2 valence electrons)
• Group 13: ns²np¹ (3 valence electrons)
• Group 17: ns²np⁵ (7 valence electrons)
2. Similar electronic configuration: Same outer shell configuration leads to similar chemical properties
3. Periodic arrangement: Elements with same valence electrons are placed vertically
4. Chemical behavior: Valence electrons determine bonding and reactivity
v. Do you expect calcium to be more reactive than sodium? Give the reason of your answer.
Answer:
No, calcium is less reactive than sodium.
Reasons:
1. Group position: Na (Group 1), Ca (Group 2)
2. Ionization energy: Ca has higher IE than Na
• Na: IE₁ = 496 kJ/mol (easier to remove 1 electron)
• Ca: IE₁ = 590 kJ/mol (harder to remove 1st electron) + IE₂ = 1145 kJ/mol (needs to remove 2 electrons)
3. Metallic bonding: Ca has stronger metallic bonds (2 valence electrons)
4. Reaction with water:
• Na: Violent reaction, moves on water surface, produces H₂ quickly
• Ca: Reacts slowly, sinks, produces H₂ slowly
5. General trend: Group 1 metals > Group 2 metals in reactivity with same period
vi. Which element has the maximum atomic radius and which element has the minimum atomic radius in third period?
Answer:
In third period (Na to Ar):
Maximum atomic radius: Sodium (Na) ~186 pm
Minimum atomic radius: Argon (Ar) ~71 pm
Explanation:
1. Trend across period: Atomic radius decreases from left to right
2. Reason: Increasing nuclear charge pulls electrons closer
3. Order (decreasing radius): Na > Mg > Al > Si > P > S > Cl > Ar
4. Values (approximate):
• Na: 186 pm, Mg: 160 pm, Al: 143 pm, Si: 117 pm
• P: 110 pm, S: 104 pm, Cl: 99 pm, Ar: 71 pm
vii. Why are the most electronegative elements present in sixth and seventh groups?
Answer:
The most electronegative elements are in groups 16 and 17 (specifically period 2):
Reasons:
1. Small atomic size: F (group 17, period 2) and O (group 16, period 2) have smallest sizes in their groups
2. High nuclear charge: Strong pull on electrons
3. Need to complete octet: Group 17 needs 1 electron, group 16 needs 2 electrons
4. Electronegativity order: F (4.0) > O (3.5) > N (3.0) ≈ Cl (3.0)
5. Top-right corner: Highest electronegativity at top-right of periodic table (excluding noble gases)
Note: Noble gases (group 18) have no electronegativity values as they don’t form bonds under normal conditions.
viii. The first ionization energy value of magnesium is less than the second one. Give reason.
Answer:
Reason: After removing first electron, magnesium ion (Mg⁺) has stable noble gas configuration (Ne-like), making second electron harder to remove.
Detailed explanation:
1. Mg atom: 1s²2s²2p⁶3s² (12 electrons)
2. First ionization: Mg → Mg⁺ + e⁻ (IE₁ = 737 kJ/mol)
• Removes 3s electron
• Mg⁺ configuration: 1s²2s²2p⁶3s¹ (like Na)
3. Second ionization: Mg⁺ → Mg²⁺ + e⁻ (IE₂ = 1450 kJ/mol)
• Removes electron from 3s orbital of Mg⁺
• Mg²⁺ configuration: 1s²2s²2p⁶ (like Ne, noble gas configuration)
4. Why IE₂ > IE₁:
• Mg²⁺ has stable noble gas configuration
• Increased effective nuclear charge on remaining electrons
• Electron being removed feels stronger nuclear attraction