Class 11 Chemistry – Chapter 10: Periodic Table (FBISE)
This section provides complete, exam-oriented notes for Class 11 Chemistry Chapter 10 – Periodic Table according to the Federal Board (FBISE) syllabus. The content is designed to help students clearly understand periodic classification and trends.
Major topics include modern periodic law, periodic trends, atomic and ionic radii, ionization energy, electron affinity, electronegativity, shielding effect, and anomalies in periodic trends. Concepts are explained with tables, comparisons, and exam-focused points.
For better preparation, students can also access video lectures, MCQs, conceptual tests, and live revision sessions on our YouTube channel and receive regular updates through our WhatsApp channel.
Notes: Periods, Groups, and Blocks of the Periodic Table
1. Basics of the Periodic Table
- Periods: The seven horizontal rows. The period number corresponds to the maximum principal quantum number ($n$) of the valence shell.
- Periods 1–3 are "short periods."
- Periods 4–7 are "long periods."
- Groups: Vertical columns containing elements with similar chemical properties.
- Traditional System: Uses letters A (Representative elements) and B (Transition elements).
- IUPAC System: Groups are numbered 1–18 from left to right.
2. Blocks of Elements
| Block | Groups | Description |
|---|---|---|
| s-block | 1 and 2 (plus He) | Valence electrons are in the $s$ sub-shell. Includes alkali and alkaline earth metals. |
| p-block | 13 to 18 | Valence electrons are in the $p$ sub-shell. Contains metals, non-metals, and metalloids. |
| d-block | 3 to 12 | Valence electrons are in the $d$ sub-shell. These are the transition metals. |
| f-block | Bottom two rows | Valence electrons are in the $f$ sub-shell. Known as lanthanides and actinides (inner transition metals). |
3. Determining Position from Electronic Configuration
- Period Number: Equal to the highest value of $n$ in the electronic configuration.
- Group Number:
- For s-block: Total valence electrons.
- For p-block: Total valence electrons $+ 10$.
Review Questions and Answers
Q1: How do you determine the period number of an element using its electronic configuration?A: The period number is determined by the highest principal quantum number ($n$) of the valence shell. For example, if the valence shell is $3s^2 3p^5$, the element is in Period 3.
Q2: If an element has a valence shell configuration of $2s^2 2p^3$, what is its group number in both the traditional and IUPAC systems?A: Total valence electrons = $2 + 3 = 5$.Traditional System: Group VA.IUPAC System: $5 + 10 =$ Group 15.
Q3: What are the specific names given to Group 1, 2, 16, 17, and 18?A:
- Group 1: Alkali metals
- Group 2: Alkaline earth metals
- Group 16: Chalcogens
- Group 17: Halogens
- Group 18: Noble gases
Q4: Why are d-block elements also called transition elements?A: Because they represent a transition in properties between the highly reactive electropositive s-block metals and the less metallic/non-metallic p-block elements.
Q5: Distinguish between representative and inner transition elements.A: Representative elements (Group A) belong to the s-block and p-block. Inner transition elements belong to the f-block, located at the bottom of the periodic table, consisting of lanthanides and actinides.
Characteristic Properties of Elements in a Group
Chemical Periodicity: The repeating pattern of element properties within the periodic table. Elements in the same group share similar chemical properties due to having similar electronic configurations in their valence shells.
Group Trends (Top to Bottom)
- Atomic Size: Gradually increases. Elements lower in a group have larger atomic radii.
- Ionization Energy: Decreases down the group. Lower elements require less energy to remove an electron.
- Electronegativity: Decreases down the group.
- Metallic Properties: Increase down the group.
- Chemical Reactivity: Generally similar within a group due to identical valence electron counts.
Problem-Solving Strategies
1. Finding Position from Electron Configuration
- Write the full configuration.
- The highest principal quantum number ($n$) of the $s$ or $p$ sub-shell is the Period Number.
- The total number of electrons in the valence shell is the Group Number.
2. Finding Configuration from Position
- Period Number = $n$ value of the valence shell.
- Group Number = Total valence electrons.
- Distribute electrons into $s$ and $p$ sub-shells accordingly.
Self-Assessment & Examples
Example: Finding Position
For Nitrogen (Atomic No. 7):
Configuration: $1s^2 2s^2 2p^3$
- Valence Shell: $2s^2 2p^3$
- Period: $2$
- Group: $2 + 3 = 5$ (Group V-A or Group 15)
Review Questions & Answers
Q1: Find the valence shell electronic configuration of Mg (Magnesium) and Cl (Chlorine) based on their position.
A:
- Mg: Located in Period 3, Group II-A. $n=3$, valence electrons = 2. Configuration: $3s^2$.
- Cl: Located in Period 3, Group VII-A. $n=3$, valence electrons = 7. Configuration: $3s^2 3p^5$.
Q2: Find the position of K (Atomic No. 19) and S (Atomic No. 16) in the periodic table.
A:
- K: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1$. Highest $n=4$, valence electrons = 1. Period 4, Group 1 (I-A).
- S: $1s^2 2s^2 2p^6 3s^2 3p^4$. Highest $n=3$, valence electrons = 6. Period 3, Group 16 (VI-A).
Q3: Which two elements make up more than 98% of the normal matter in the universe?
A: Hydrogen (approx. 73%) and Helium (approx. 25%).
1. Classification of Elements
- Metals: Located on the left side of the periodic table. They are characterized as good conductors, malleable, and ductile.
- Non-metals: Located on the right side of the periodic table (except Hydrogen). They are poor conductors, non-malleable, and non-ductile.
- Metalloids: Also known as semi-metals, these elements border the "leader line" separating metals and non-metals and possess properties of both.
2. Periodicity of Properties
Physical and chemical properties of elements vary in a periodic manner because they depend on the electronic configuration of the atoms.
- Groups: Elements in the same group have similar valence shell configurations, leading to similar chemical properties. Atomic size changes gradually from top to bottom.
- Periods: Moving from left to right, the number of valence electrons increases gradually, causing a gradation in chemical and physical properties.
3. Electron Affinity
Definition: The amount of energy released when an electron is added to the valence shell of an isolated gaseous atom to form a uni-negative ion.
General Equation: $X_{(g)} + e^- \longrightarrow X^-_{(g)} + \text{electron affinity}$
Factors Affecting Electron Affinity:
- Nuclear Charge: As nuclear charge increases (more protons), the attraction for electrons increases.
- Trend: Increases from left to right in a period.
- Atomic Size: In larger atoms, the valence shell is farther from the nucleus, weakening the attraction for incoming electrons.
- Trend: Decreases from top to bottom in a group.
- Shielding Effect: Inner shell electrons shield the nucleus's attraction from outer electrons.
- Trend: An increase in shielding effect (down a group) leads to a decrease in electron affinity.
Relevant Questions and Answers
Q1: Why do elements in the same group exhibit similar chemical properties?
A: Chemical properties depend on the electronic configuration of the valence shell. Since all elements in a specific group have a similar valence shell electronic configuration, they react in similar ways.
Q2: Explain the trend of electron affinity across a period.
A: Electron affinity tends to increase from left to right across a period. This is because the nuclear charge increases and the atomic size decreases, allowing the nucleus to bind an additional electron more tightly, releasing more energy.
Q3: Why does electron affinity decrease as you move down a group?
A: As you move down a group, atomic size increases and the shielding effect becomes stronger due to more inner electron shells. Consequently, the nucleus has a weaker attraction for an added electron, and less energy is released.
Q4: What is the significance of the "blue leader line" in the periodic table?
A: The blue leader line serves as a boundary to separate metals (on the left) from non-metals (on the right). Elements adjacent to this line are typically metalloids.
Q5: Based on the provided data, which group of elements has an electron affinity of $0$?
A: According to the table of main group elements, the Noble Gases (He, Ne, Ar, Kr, Xe, Rn) and some alkaline earth metals like Be, Mg, and Ca have an electron affinity of $0$, as they have stable electronic configurations and do not easily accept extra electrons.
Metallic and Non-Metallic Behavior
Metallic Character: Metals are found on the left of the periodic table. They tend to lose valence electrons to form cations because they have low ionization energy and low electronegativity.
Non-Metallic Character: Non-metals are on the right side. They tend to gain or share electrons to achieve stability, characterized by high ionization energy and high electronegativity.
Trends in the Periodic Table
- Down a Group: Metallic character increases (e.g., Group 1) because atomic radii increase, making it easier to lose electrons. Conversely, non-metallic character decreases (e.g., Group 17).
- Across a Period: Metallic character decreases and non-metallic character increases as you move from left to right.
Electronegativity and Bond Types
The difference in electronegativity ($\Delta EN$) determines the bond nature:
- Ionic Bond: $\Delta EN > 1.8$ (e.g., $NaCl$ where $\Delta EN = 2.23$).
- Polar Covalent: $0.4 < \Delta EN < 1.8$ (e.g., $H-F$ where $\Delta EN = 1.78$).
- Non-polar Covalent: $\Delta EN < 0.4$ (e.g., $C-H$ where $\Delta EN = 0.35$).
Chemical Properties and Oxides
The nature of oxides changes across Period 3 ($Na$ to $Cl$):
| Oxide Type | Examples | Behavior |
|---|---|---|
| Basic | $Na_{2}O, MgO$ | React with water to form bases; react with acids to form salt and water. |
| Amphoteric | $Al_{2}O_{3}$ | Reacts with both acids and bases. |
| Acidic | $CO_{2}, P_{4}O_{10}, SO_{3}, Cl_{2}O_{7}$ | React with water to form acids. |
Oxidation Numbers
In Period 3 oxides, the maximum oxidation number increases from $+1$ to $+7$:
$Na(+1) \rightarrow Mg(+2) \rightarrow Al(+3) \rightarrow Si(+4) \rightarrow P(+5) \rightarrow S(+6) \rightarrow Cl(+7)$
Questions and Answers
Q1: Why does metallic character increase down a group?
A1: Down a group, the atomic radius increases. This reduces the nuclear attraction on valence electrons, making it easier for the atom to lose electrons and form cations.
Q2: Predict the bond type for $MgO$ given $EN$ of $Mg = 1.31$ and $O = 3.44$.
A2: $\Delta EN = 3.44 - 1.31 = 2.13$. Since the difference is greater than $1.8$, the bond is ionic.
Q3: What is meant by an "amphoteric" oxide? Provide an example.
A3: An amphoteric oxide is one that displays both acidic and basic properties, reacting with both acids and bases to form salts. An example is Aluminium oxide ($Al_{2}O_{3}$).
Q4: Write the chemical equation for Sodium reacting with water.
A4: $2Na + 2H_{2}O \rightarrow 2NaOH + H_{2}$
Q5: Why do Phosphorus and Sulphur show variable oxidation numbers?
A5: They can expand their octet. For example, Sulphur can have an oxidation state of $+4$ in $SO_{2}$ and $+6$ in $SO_{3}$.
Chemistry Study Notes: Periodicity and Element Identification
1. Identifying Unknown Elements
The position and properties of an element in the periodic table can be determined using its atomic number and electronic configuration.
- Determining Period: Indicated by the highest principal quantum number ($n$) in the electronic configuration.
- Determining Group: Indicated by the number of electrons in the valence (outermost) shell.
| Atomic Number ($Z$) | Electronic Configuration | Valence Shell | Period | Group | Element |
|---|---|---|---|---|---|
| 17 | $1s^2, 2s^2, 2p^6, 3s^2, 3p^5$ | $3s^2, 3p^5$ | 3 | 17 (VIIA) | Chlorine (Cl) |
| 19 | $1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^1$ | $4s^1$ | 4 | 1 (IA) | Potassium (K) |
| 35 | $1s^2, 2s^2, 2p^6, 3s^2, 3p^6, 4s^2, 3d^{10}, 4p^5$ | $4s^2, 4p^5$ | 4 | 17 (VIIA) | Bromine (Br) |
2. Reactions of Chlorides with Water
The nature of the solution formed when chlorides react with water depends on the type of chemical bonding (ionic vs. covalent).
- Ionic Chlorides: (e.g., $NaCl, MgCl_2$) Dissolve in water to form neutral solutions ($pH = 7$).
- Covalent Chlorides: (e.g., $AlCl_3, SiCl_4, PCl_5$) React vigorously with water to form acidic solutions ($pH < 7$).
Key Reactions:
- $AlCl_3 + 3H_2O \rightarrow Al(OH)_3 + 3HCl$ ($pH = 3$)
- $SiCl_4 + 4H_2O \rightarrow Si(OH)_4 + 4HCl$ ($pH = 0$)
- $PCl_5 + 4H_2O \rightarrow H_3PO_4 + 5HCl$ ($pH = 0$)
3. Periodic Trends (Group 1 & 17)
Ionization Energy (Group 1): Decreases down the group because the increase in atomic radii and shielding effect reduces the nuclear attraction on valence electrons.
Electron Affinity (Group 17): Generally decreases down the group. Exception: Fluorine has a lower electron affinity than Chlorine because F's small atomic size causes high electron-electron repulsion for the incoming electron.
Review Questions and Answers
Q1: An element has an atomic number of 35. Predict its chemical nature and molecular form.
A: Based on the atomic number 35, the electronic configuration ends in $4s^2 4p^5$. It is a non-metal belonging to Group 17 (Halogens). It exists as a diatomic molecule ($Br_2$), forms salts with metals, and forms acidic oxides.
Q2: Why does Potassium (Atomic Number 19) react violently with water?
A: Potassium is an alkali metal (Group 1). Due to its large atomic radius and low ionization energy, it easily loses its single valence electron, making it highly reactive with water to produce hydrogen gas and a basic solution.
Q3: Compare the pH of a solution of $NaCl$ versus $SiCl_4$ in water.
A: $NaCl$ is an ionic chloride and forms a neutral solution with a $pH = 7$. $SiCl_4$ is a covalent chloride that undergoes hydrolysis to produce $HCl$, resulting in a highly acidic solution with a $pH = 0$.
Q4: Why is the electron affinity of Fluorine lower than that of Chlorine?
A: Although Fluorine is more electronegative, its atomic size is very small. This results in high electron density, which causes the incoming electron to be repelled by the existing electrons more strongly than in Chlorine.