ICSE Chemistry Chapter 1: Periodic Table, Periodic Properties, and Variations

Syllabus and Introduction to Periodicity in MARCH 2020 Scope

• Syllabus Overview:

  • The study focuses on periodic properties and their variations in groups and periods, specifically: atomic size, metallic character, non-metallic character, ionisation potential, electron affinity, and electronegativity.

  • The basis for periodicity is the atomic number of elements.

  • The scope includes the complete Modern Periodic Table, though examination questions are generally limited to elements up to Period $3$ ($Ar$ - Argon).

  • Explanations must be based on nuclear charge and electron shells, not orbitals.

  • Special reference is given to alkali metals (Group $1$) and halogens (Group $17$).

• IUPAC Notation Updates:

  • Groups are numbered from $1$ to $18$.

  • This replaces older notations like $IA..VIIA, VIIIB, IB...VIIB$, and $0$.

  • Both notations are accepted for examination purposes.

  • Example correlation:

  • $IA$ = Group $1$

  • $IIA$ = Group $2$

  • $IIIA$ = Group $13$

  • $IVA$ = Group $14$

  • $VA$ = Group $15$

  • $VIA$ = Group $16$

  • $VIIA$ = Group $17$

  • $0$ = Group $18$

• The Need for Classification:

  • Elements were grouped into families based on maximum resemblance to simplify study.

  • Properties were observed to reappear at regular intervals when arranged in systematic tables.

Historical Development of Element Classification

• Early Chemists:

  • Arranged elements based on valency, metallic, and non-metallic characters.

  • The method was discarded because elements frequently showed variable valency and dual character.

• Dobereiner ($1815$) - Law of Triads:

  • Elements were arranged in increasing order of atomic weights in groups of three (triads).

  • The atomic weight of the middle element was generally the average of the other two.

  • Discarded because it did not hold true for all known elements.

• Newland ($1864$) - Law of Octaves:

  • Arranged elements in increasing order of atomic weights in series of eight.

  • Found that the properties of every eighth element were a repetition of the first element.

  • Discarded because it failed to leave space for undiscovered elements.

• Mendeleeff ($1869$):

  • Arranged elements in increasing order of atomic weights in the Mendeleeff's Periodic Table.

  • Stated the Periodic Law: "Properties of elements are periodic functions of their atomic weights."

  • Could not justify the position of certain elements, rare earths, and isotopes.

• Moseley ($1912$):

  • Modified the classification by using atomic numbers instead of atomic weights.

  • Arranged elements in the Modern Periodic Table.

  • Stated the Modern Periodic Law: "Properties of elements are periodic functions of their atomic numbers."

  • Removed most defects of Mendeleeff’s table.

The Modern Periodic Table (Long Form)

• Basis of Classification:

  • Physical and chemical properties depend on the number of electrons and their arrangement.

  • Atomic number ($Z$) is equal to the number of electrons in energy shells and is the fundamental property of an element.

• Salient Features:

  • Elements are arranged in increasing order of atomic numbers.

  • Seven horizontal rows are called Periods.

  • Eighteen vertical columns are called Groups.

  • Logical Completion: Each period begins with an element having one electron in its outermost shell and ends with a Group $18$ (zero group) element having a completely filled outer shell.

  • Character Transition: Across a period, there is a transition from metallic to non-metallic character.

  • Group Properties: Each vertical column contains elements with the same outer electronic configuration and similar properties.

• Element Grouping:

  • Normal Elements: Groups $1, 2$, and $13$ to $17$ ($IA$ to $VIIA$).

  • Transition Elements: Groups $3$ to $12$ ($IB$ to $VIIB$ and $VIII$).

  • Noble/Inert Gases: Group $18$ ($0$ group) on the extreme right.

  • Reactive Metals: Placed in Groups $1$ ($IA$) and $2$ ($IIA$).

  • Non-metals: Placed in the upper right corner.

• Periodicity:

  • Recurrence in properties occurs after differences in atomic numbers of $2, 8, 18$, or $32$.

  • This is due to the recurrence of similar valence shell electronic configurations.

Periods in the Modern Periodic Table

• Period Characteristics:

  • The period number ($1, 2, 3$, etc.) signifies the number of electron shells of an element.

  • Period $1$: $H$ has $1$ shell.

  • Period $2$: $Li$ has $2$ shells.

  • Period $3$: $Na$ has $3$ shells.

• Classification of the Seven Periods:

  • Period 1 (Shortest): $2$ elements ($H$ to $He$), Atomic Nos. $1$ to $2$.

  • Period 2 (Short): $8$ elements ($Li$ to $Ne$), Atomic Nos. $3$ to $10$.

  • Period 3 (Short): $8$ elements ($Na$ to $Ar$), Atomic Nos. $11$ to $18$.

  • Period 4 (Long): $18$ elements ($K$ to $Kr$), Atomic Nos. $19$ to $36$.

  • Period 5 (Long): $18$ elements ($Rb$ to $Xe$), Atomic Nos. $37$ to $54$.

  • Period 6 (Longest): $32$ elements ($Cs$ to $Rn$), Atomic Nos. $55$ to $86$. Includes Lanthanide Series (Rare Earth Elements).

  • Period 7 (Incomplete): $26$ elements ($Fr$ to $Uub$), Atomic Nos. $87$ to $112$. Includes Actinide Series (Radioactive Elements).

• Bridge Elements (Period 2):

  • Show similar properties diagonally with the period of the next group.

  • Pairs include: $Li$ and $Mg$, $Be$ and $Al$, $B$ and $Si$.

• Typical Elements (Period 3):

  • These elements ($Na, Mg, Al, Si, P, S, Cl$) represent the properties of their respective groups.

Property Trends Across Period 2 and Period 3

• Similarity: Number of electron shells remains the same across the period.

• Transitions:

  • Valence electrons increase by one.

  • Non-metallic character increases.

• Period 2 Data:

  • Elements: $Li, Be, B, C, N, O, F, Ne$

  • Atomic Numbers: $3, 4, 5, 6, 7, 8, 9, 10$

  • Configuration: $2,1$ to $2,8$

  • State: Metals ($Li, Be$), Metalloid ($B$), Non-metals ($C$ to $F$), Noble gas ($Ne$).

  • Valency: $1, 2, 3, 4, 3, 2, 1, 0$.

• Period 3 Data and Compound Chemistry:

  • Elements: $Na, Mg, Al, Si, P, S, Cl, Ar$

  • Configuration: $2,8,1$ to $2,8,8$

  • Chlorides:

  • Formulas: $NaCl, MgCl_2, AlCl_3, SiCl_4, PCl_3/PCl_5, S_2Cl_2, Cl_2$

  • Bonding: Ionic ($Na, Mg$), Ionic/Covalent ($Al$), Covalent ($Si$ to $Cl$).

  • State: Solid ($Na$ to $Al$), Liquid ($Si$), Liquid/Solid ($P, S, Cl$).

  • Oxides:

  • Formulas: $Na_2O, MgO, Al_2O_3, SiO_2, P_2O_5, SO_3, Cl_2O_7$

  • Character: Strongly basic ($Na_2O$), Basic ($MgO$), Amphoteric ($Al_2O_3$), Weakly acidic ($SiO_2$), Acidic ($P_2O_5, SO_3$), Strongly acidic ($Cl_2O_7$).

  • Hydroxides and Oxy-acids:

  • Formulas: $NaOH, Mg(OH)_2, Al(OH)_3, H_2SiO_3, H_3PO_4$ or $HPO_3$, $H_2SO_4, HClO_4$

  • Character: Strong base ($NaOH$) to Strong acid ($HClO_4$).

  • Hydrides:

  • Formulas: $NaH, MgH_2, AlH_3, SiH_4, PH_3, H_2S, HCl$

  • Character: Strong base to Strong acid.

Groups in the Modern Periodic Table

• Group Characteristics:

  • Vertical columns; group number signifies the number of valence electrons.

  • Transition elements also generally have $2$ valence electrons, e.g., $_{26}Fe [2, 8, 14, 2]$ and $_{30}Zn [2, 8, 18, 2]$.

• Specific Groups:

  • Group 1 (IA): Alkali Metals ($Li$ to $Fr$) - Light metals.

  • Group 2 (IIA): Alkaline Earth Metals ($Be$ to $Ra$) - Light metals.

  • Group 3 to 12: Transition Elements ($Sc$ to $Zn$) - Heavy metals. Properties lie between electropositive metals and non-metals. Electrons fill inner orbitals.

  • Inner Transition Elements: Lanthanides ($Ce$ to $Lu$) and Actinides ($Th$ to $Lr$). Placed at the bottom.

  • Group 13 to 16: Post-Transition Elements (includes metals, metalloids, and non-metals).

  • Group 17 (VIIA): Halogens ($F$ to $At$).

  • Group 18 (0): Noble/Inert Gases ($He$ to $Rn$). Stable configurations, unreactive.

Property Trends Down a Group (Subgroup)

• Similarity: Valence electrons remain the same. Chemical properties are similar.

• Transition:

  • Number of shells increases by one with each step down.

  • Metallic character (electropositive character) increases.

• Comparison: Alkali Metals (Group 1) vs. Halogens (Group 17):

  • Valency: Alkali Metals = $1$ (electron donors); Halogens = $1$ (electron acceptors, $7$ valence electrons).

  • Nature: Alkali Metals = Highly electropositive, soft metals; Halogens = Highly electronegative non-metals.

  • Conductivity: Alkali Metals = Good conductors; Halogens = Bad/non-conductors.

  • Reducing/Oxidising: Alkali Metals = Strong reducing agents; Halogens = Strong oxidising agents.

  • Hydrides: Alkali Metals form ionic hydrides (e.g., $LiH, NaH$); Halogens form covalent hydrides (e.g., $HF, HCl$).

Periodic Properties: Definitions and Variations

1. Atomic Radius:

   • Definition: Distance between the center of the nucleus and the outermost shell of the atom.

   • Trends: Decreases across a period (due to increasing nuclear charge); Increases down a group (due to increasing number of shells).

2. Ionisation Potential (I.P.) or Ionisation Energy:

   • Definition: Amount of energy required to remove a loosely bound electron from the outermost shell of an isolated gaseous atom.    $M \rightarrow M^+ + e^-$

   • Units: Electron volt ($eV$).

   • Trends: Increases across a period; Decreases down a group.

3. Electron Affinity (E.A.):

   • Definition: Amount of energy released when an atom in the gaseous state accepts an electron to form an anion.    $X + e^- \rightarrow X^- + \text{energy released}$

   • Trends: Increases across a period; Decreases down a group.

   • Note: Noble gases have zero $E.A.$ because their shells are full.

4. Electronegativity (E.N.):

   • Definition: Tendency of an atom to attract electrons to itself when combined in a compound.

   • Factors: Same as $E.A.$ (Atomic size and nuclear charge).

   • Trends: Increases across a period; Decreases down a group.

   • Extreme Elements: Fluorine ($F$) is the most electronegative ($4.0$); Caesium ($Cs$) is the least.

5. Metallic and Non-Metallic Character:

   • Metallic Character: Tendency to lose electrons (Electropositive character).

   • Non-Metallic Character: Tendency to gain electrons (Electronegative character).

   • Trends: Across a period, metallic character decreases while non-metallic increases. Down a group, metallic character increases while non-metallic decreases.

Factors Affecting Trends

• Number of Shells:

  • Increase in shells increases atomic size.

  • This decreases nuclear attraction on outer electrons, reducing I.P., E.A., and E.N.

• Nuclear Charge:

  • Positive charge on the nucleus (equivalent to Atomic Number).

  • Increase in nuclear charge decreases atomic size (electrons are pulled closer).

  • This increases I.P., E.A., and E.N. because electrons are held more firmly.

Summary of Trends and Relationships

• Across a Period (Left to Right):

  • Atomic Size: Decreases.

  • Ionisation Potential: Increases.

  • Electron Affinity: Increases.

  • Electronegativity: Increases.

  • Metallic Character: Decreases.

  • Non-Metallic Character: Increases.

• Down a Group:

  • Atomic Size: Increases.

  • Ionisation Potential: Decreases.

  • Electron Affinity: Decreases.

  • Electronegativity: Decreases.

  • Metallic Character: Increases.

  • Non-Metallic Character: Decreases.

• Special Note on Cations and Anions:

  • A Cation (e.g., $Na^+$) is smaller than its parent atom ($Na$) because electrons are attracted more strongly by the nucleus after losing one.

  • An Anion (e.g., $Cl^-$) is larger than its parent atom due to electronic repulsion and decreased nuclear hold.

• Stability and Radioactive Elements:

  • Elements with a neutron/proton ($n/p$) ratio around $1$ are stable (e.g., Light metals like $Na$ and $K$).

  • Elements with an $n/p$ ratio above $1.5$ are considered radioactive/unstable (e.g., heavy metals like Uranium, $_{92}^{236}U$, with a ratio of $1.52$).

• Physical Properties Periodicity:

  • Across a period: Density and melting points generally increase and then decrease.

  • Down a group: Density generally increases; melting/boiling points of Group $1$ metals decrease gradually.