Q.What are the major differences between metals and non-metals?

Understanding Acid Strength in Aqueous Solution

The strength of an acid in water depends on how readily it donates a proton (H⁺). For hydrogen halides, this means breaking the H-X bond and releasing H⁺ into solution:

The easier this bond breaks, the stronger the acid.

Evaluating the Assertion (A)

Let's verify the acidic strength order: HF < HCl < HBr < HI

1. HF is the weakest acid despite fluorine being the most electronegative halogen. In aqueous solution, HF forms strong hydrogen bonds with water and exists largely as undissociated molecules. It's a weak acid with pKₐ ≈ 3.2.

2. HCl, HBr, and HI are all strong acids in water, but their strength increases down the group. HI is the strongest, followed by HBr, then HCl.

3. The experimental pKₐ values confirm this trend:

   • HF: pKₐ ≈ 3.2 (weak acid)
   • HCl: pKₐ ≈ -7 (strong acid)
   • HBr: pKₐ ≈ -9 (stronger)
   • HI: pKₐ ≈ -10 (strongest)

Assertion (A) is correct.

Evaluating the Reason (R)

The bond dissociation enthalpy (BDE) trend: HF > HCl > HBr > HI

1. Bond strength decreases down the group because:

   • Atomic size increases: F < Cl < Br < I
   • The H-X bond length increases
   • Orbital overlap becomes less effective with larger, more diffuse halogen orbitals

2. Actual BDE values (in kJ/mol):

   • H-F: 568 (strongest bond)
   • H-Cl: 431
   • H-Br: 366
   • H-I: 298 (weakest bond)

Reason (R) is correct.

Does R Explain A?

Now the crucial question: Does bond dissociation enthalpy explain acidic strength?

1. In aqueous solution, the weaker the H-X bond, the easier it breaks, releasing H⁺ more readily.

2. HI has the weakest bond (lowest BDE = 298 kJ/mol), so it breaks most easily → strongest acid.

3. HF has the strongest bond (highest BDE = 568 kJ/mol), so it breaks least easily → weakest acid. …

The key concept here is lattice energy and how it relates to ion size. Superoxides (containing the O₂⁻ ion) are stable only for the largest alkali metals because the superoxide anion is quite large. For a salt to be stable, the lattice energy must be high enough to compensate for the energy required to form the ions. Lattice energy is inversely proportional to the sum of ionic radii: smaller cations give higher lattice energy, but only if the anion is also small. Here, the anion is large, so a small cation actually destabilizes the crystal because the ions don't pack well — the lattice energy becomes too low to offset the formation energy.

1. Recall the trend in alkali metal cation sizes.

   Going down Group 1: Li⁺ < Na⁺ < K⁺ < Rb⁺ < Cs⁺.

   Li and Na are too small to stabilize O₂⁻; they form normal oxides (Li₂O, Na₂O) or peroxides. Only K, Rb, and Cs form stable superoxides.

2. Understand the stability criterion for superoxides.

   The superoxide ion (O₂⁻) is large. For a salt MO₂ (M = alkali metal), the lattice energy is roughly proportional to $\frac{1}{r_+ + r_-}$. Since $r_-$ is fixed, a larger $r_+$ gives a smaller lattice energy — but that seems counterintuitive for stability. The catch: the formation of the superoxide ion itself is energetically costly. A larger cation produces a softer, more polarizable lattice that better accommodates the large, diffuse anion, lowering the overall energy of the crystal.

3. Apply the known experimental trend. …

The key here is to separate two common confusions: solubility versus reactivity with water, and trends in basic strength for alkaline earth metal oxides. Many students mistakenly think that if a compound is "insoluble," it cannot react — but that's not the case here.

Concept and Intuition

The oxides MgO, CaO, SrO, and BaO are ionic compounds of the alkaline earth metals. When placed in water, they do not simply dissolve as neutral molecules; instead, they undergo a chemical reaction with water to form the corresponding metal hydroxides:

This reaction is often called "slaking" (especially for CaO, quicklime). The resulting hydroxides are sparingly soluble to moderately soluble, but the oxides themselves are not "insoluble" in the sense of being inert — they react. In chemistry, "solubility" of an oxide usually refers to whether it reacts with water or not. By that definition, these oxides are soluble (they react). So the assertion (A) is false.

The reason (R) discusses basic strength. As we go down Group 2 (from Mg to Ba), the metal cation becomes larger and its charge density decreases. This makes the M–O bond weaker and the oxide ion more available to accept a proton (or to react with water), so basic strength increases. That trend is correct.

Step-by-step reasoning

1. Check the assertion (A): "MgO, CaO, SrO and BaO are insoluble in water."

   • In reality, all these oxides react vigorously (especially CaO, SrO, BaO) with water to form hydroxides.
   • MgO reacts more slowly but still forms Mg(OH)₂.
   • Therefore, they are not insoluble; they are chemically reactive with water.
   • So (A) is incorrect.

2. Check the reason (R): "In aqueous medium the basic strength of MgO, CaO, SrO and BaO increases with increase in the atomic number of metal."

   • Basic strength of an oxide depends on how easily it donates OH⁻ or accepts H⁺. …

The solubility of ionic compounds in organic solvents like ethanol hinges on one central idea: like dissolves like. Purely ionic compounds dissolve in polar solvents like water, while compounds with significant covalent character dissolve in less polar or organic solvents. The key is understanding what makes LiCl and BeCl₂ behave differently from typical alkali and alkaline earth metal chlorides.

Lithium and beryllium occupy unique positions in their groups. Both have exceptionally small ionic radii (Li⁺ ≈ 76 pm, Be²⁺ ≈ 31 pm) and high charge densities. According to Fajans' rules, a small, highly charged cation with high polarizing power distorts the electron cloud of a large, polarizable anion like Cl⁻, introducing covalent character into what would otherwise be an ionic bond. This covalent character makes LiCl and BeCl₂ soluble in organic solvents.

Now let's examine both statements:

1. Assertion (A): LiCl and BeCl₂ are soluble in ethanol

   This is true. Both compounds exhibit anomalous behavior compared to their group members. LiCl dissolves readily in ethanol, acetone, and pyridine — solvents in which NaCl, KCl, etc., are essentially insoluble. BeCl₂ is even more covalent (beryllium has a +2 charge on an even smaller ion) and dissolves in ethanol and ether. The high covalent character reduces the lattice energy and allows interaction with the organic solvent molecules.

2. Reason (R): Lithium and beryllium are harder than their respective group elements

   This statement is also true. Hardness in metals correlates with small atomic size, high charge density, and strong metallic bonding. Lithium is harder than sodium or potassium; beryllium is harder than magnesium or calcium. The small size leads to stronger electrostatic attraction in the metallic lattice.

3. Does (R) explain (A)? …

Concept and Intuition

This is a fact-checking problem from inorganic chemistry. Each option touches a distinct topic: the composition of calgon (a water softener), the solubility behavior of beryllium halides, the reducing power trends among alkali metals, and the composition of white metal (an alloy). The trick is to recall or reason out the correct facts for each—only one will be false.

Step-by-step reasoning

1. Option (A): Molecular formula of calgon is NaAlSiO₄

   Calgon is a commercial water-softening agent. Its chemical name is sodium hexametaphosphate, with the formula (NaPO₃)₆ or Na₆P₆O₁₈. The formula NaAlSiO₄ corresponds to a feldspathoid mineral (like nepheline), not calgon.

   → This statement is incorrect.

2. Option (B): Beryllium halides are soluble in organic solvents

   Beryllium halides (e.g., BeCl₂) are covalent in nature due to the small size and high charge of Be²⁺. They dissolve in organic solvents like ether, benzene, etc. This is a well-known property.

   → This statement is correct.

3. Option (C): Among alkali metals, the reducing property of sodium is least in aqueous solution

   In aqueous solution, the reducing power of alkali metals depends on their standard electrode potentials (E°). The order of E° values (more negative = stronger reductant) is: Li > K > Ba > Sr > Ca > Na > Mg. Actually, lithium has the most negative E° (−3.04 V), while sodium is −2.71 V, and potassium is −2.93 V. So sodium is not the least; it is intermediate. But the statement says “sodium is least”—that is false. However, note: in aqueous solution, lithium is the strongest reductant, and sodium is weaker than potassium? Wait, check: E°(Li⁺/Li) = −3.04 V, E°(K⁺/K) = −2.93 V, E°(Na⁺/Na) = −2.71 V. So sodium is actually the weakest reductant among Li, Na, K. Yes, because a less negative E° means weaker reducing power. So sodium is indeed the least reducing among Li, Na, K. But the statement says “among alkali metals” — that includes all: Li, Na, K, Rb, Cs. Rb and Cs have even more negative E° (~−2.98 V and −3.03 V). So sodium is still the least? Actually, Cs: −3.03 V, Rb: −2.98 V, K: −2.93 V, Na: −2.71 V, Li: −3.04 V. So sodium has the least negative E°, meaning it is the weakest reductant in aqueous solution. So this statement is correct.

4. Option (D): White metal is an alloy of Lithium

   “White metal” is a term used for various alloys, but commonly it refers to tin-based or lead-based alloys (e.g., Babbitt metal) used for bearings. It does not contain lithium as a major component. Lithium alloys are not called white metal.

   → This statement is incorrect as well? Wait, we need only one incorrect statement. Let’s double-check.

   Actually, “white metal” can also refer to a lithium-containing alloy? No, standard reference: white metal is an alloy of tin, antimony, and copper (or lead). Lithium is not involved. So (D) is false. But we already found (A) false. So two false? That can’t be—only one option is “not correct.” Let’s re-examine (C) carefully. …

Concept. $E^\circ_{M^+/M}$ is fixed by a thermodynamic cycle: sublimation $+$ ionisation energy $+$ hydration energy. For $\text{Li}^+$ the very large (most negative) hydration energy of the tiny ion dominates, making $\text{Li}$'s $E^\circ$ the most negative even though it is not the most reactive by ionisation energy alone.

Standard values (V vs SHE):

• Lowest (most negative) $=\text{Li }(-3.04)\Rightarrow X=\text{Li}$. …

The key concept here is standard reduction potential $E^\circ$, which measures the tendency of a metal ion $M^+$ to gain an electron and become the metal $M$. For alkali metals, we are looking at the half-reaction:

A more negative $E^\circ$ means the metal is more easily oxidized (i.e., it is a stronger reducing agent). Intuitively, as we go down Group 1, the atomic radius increases, the ionization energy decreases, and the metal should become more reactive — so $E^\circ$ should become more negative. But lithium, the smallest, has the most negative $E^\circ$ of all. Why? Because the reduction potential depends not only on ionization energy but also on hydration enthalpy and sublimation enthalpy in a thermodynamic cycle. Lithium’s tiny size gives it an exceptionally large hydration energy, which stabilizes the ion in solution and shifts the equilibrium toward $Li^+$, making $E^\circ$ more negative.

Let’s work through the reasoning step by step.

1. Recall the trend in reactivity for alkali metals.

   Reactivity increases down the group: Li < Na < K < Rb < Cs. This is because the ionization energy decreases, making it easier to lose an electron. Since a more negative $E^\circ$ corresponds to greater ease of oxidation, we expect $E^\circ$ to become more negative down the group. That would suggest: Li (least negative) < Na < K (most negative). But the actual data show the opposite for Li.

2. Look at the actual standard reduction potentials.

   The values (in volts) are:

So the order from most negative to least negative is: Li (-3.04) < K (-2.93) < Na (-2.71). That means the increasing order (from most negative to least negative) is: Li < K < Na. But wait — the question asks for increasing order of $E^\circ$, which means from smallest (most negative) to largest (least negative). That gives Li < K < Na. However, none of the options match that exactly. Let’s check the options again:

• (A) Na < Li < K
• (B) K < Na < Li
• (C) Li < Na < K
• (D) Li < K < Na

Option (D) is Li < K < Na, which matches the actual data. But is that the correct answer? Let’s verify carefully.

3. Re-examine the data — a common pitfall. Many textbooks and standard tables list: $\begin{aligned} \text{Li: } &-3.04\ \text{V} \\ \text{Na: } &-2.71\ \text{V} \\ \text{K: } &-2.93\ \text{V} \end{aligned}$ …

The question asks for the correct order of Lewis acidic character of boron trihalides. Lewis acidity here means the tendency of the boron atom to accept a pair of electrons from a Lewis base. In $\mathrm{BX_3}$ molecules, boron is $sp^2$ hybridized with an empty $p_z$ orbital perpendicular to the molecular plane. This empty orbital is what accepts electrons, making the molecule a Lewis acid.

Now, intuition might suggest that the most electronegative halogen (fluorine) would pull the most electron density away from boron, making it the most electron-deficient and thus the strongest Lewis acid. That would give $\mathrm{BF_3 > BCl_3 > BBr_3 > BI_3}$, which is option (D). But this is wrong — and the reason is a classic pitfall.

The key concept is back-bonding (or π-bonding). Each halogen has lone pairs of electrons. The empty $p_z$ orbital on boron can accept electron density from a filled $p$ orbital on the halogen, forming a $\pi$ bond. This back-donation partially fills the empty orbital on boron, making it less eager to accept electrons from a Lewis base. The effectiveness of this back-bonding depends on the size and energy match between the halogen's $p$ orbital and boron's $p$ orbital.

Let's work through the reasoning step by step.

1. Identify the two opposing effects.

   The inductive effect (through σ-bonds) pulls electron density away from boron, making it more Lewis acidic. The back-bonding effect (through π-bonds) donates electron density to boron, making it less Lewis acidic. The actual Lewis acidity is the net result of these two.

2. Analyze the inductive effect.

   Electronegativity decreases down the group: $\mathrm{F > Cl > Br > I}$. So the inductive withdrawal of electrons from boron is strongest for $\mathrm{BF_3}$ and weakest for $\mathrm{BI_3}$. If only this mattered, the order would be $\mathrm{BF_3 > BCl_3 > BBr_3 > BI_3}$.

3. Analyze the back-bonding effect.

   For effective π-bonding, the halogen's $p$ orbital must be similar in size and energy to boron's $2p$ orbital. Fluorine has a $2p$ orbital — the same principal quantum number as boron's $2p$ — so the overlap is excellent. This makes back-bonding in $\mathrm{BF_3}$ very strong. As we go down the group, the halogen's $p$ orbitals become larger ($3p$, $4p$, $5p$) and their energy levels differ more from boron's $2p$, so the overlap becomes poorer. Thus, back-bonding decreases in the order: $\mathrm{BF_3 > BCl_3 > BBr_3 > BI_3}$. …

When ions have the same number of electrons but different nuclear charges, a beautiful pattern emerges: the more protons in the nucleus, the tighter the electron cloud is pulled inward. This is the key to comparing isoelectronic species.

Let me first verify that these ions are indeed isoelectronic:

All four have the neon configuration: $1s^2 2s^2 2p^6$.

Now the ranking becomes straightforward:

1. Identify the nuclear charges: $\mathrm{O}$ has 8 protons, $\mathrm{F}$ has 9, $\mathrm{Na}$ has 11, and $\mathrm{Mg}$ has 12. …

The question asks which metal has the least reactivity with molecular oxygen. Reactivity with oxygen is essentially a measure of how easily a metal gets oxidised — that is, how readily it loses electrons to form an oxide. This is directly linked to the metal's position in the electrochemical series or its standard reduction potential.

Metals like lithium and magnesium are highly electropositive — they have a strong tendency to lose electrons and form stable oxides (Li$_2$O and MgO). Iron also reacts with oxygen, especially under moist conditions, to form rust (Fe$_2$O$_3$). Platinum, on the other hand, is a noble metal. It has a very high reduction potential, meaning it does not easily lose electrons. In fact, platinum is famously inert and is often used as a catalyst or electrode precisely because it resists oxidation.

Let’s break it down step by step.

1. Understand the trend in reactivity with oxygen.

   Reactivity with oxygen increases as you go down Group 1 and Group 2 (Li, Mg are very reactive) and decreases for transition metals as you move toward the right of the periodic table (Fe is moderately reactive, Pt is very unreactive). The most reactive metals form oxides readily and even burn in oxygen; the least reactive ones (like gold and platinum) do not oxidise under normal conditions.

2. Compare the given metals.

   • Li (Lithium): An alkali metal. It reacts vigorously with oxygen to form lithium oxide ($4\text{Li} + \text{O}_2 \rightarrow 2\text{Li}_2\text{O}$). Very high reactivity.
   • Mg (Magnesium): An alkaline earth metal. It burns in oxygen with a brilliant white flame to form magnesium oxide ($2\text{Mg} + \text{O}_2 \rightarrow 2\text{MgO}$). High reactivity.
   • Fe (Iron): A transition metal. It reacts slowly with oxygen in the presence of moisture to form rust (hydrated iron(III) oxide). Moderate reactivity. …

The concept here is straightforward: a hydride’s reducing ability depends on how easily it can donate hydrogen (or electrons) to another species. For group 15 hydrides ($\mathrm{NH_3, PH_3, AsH_3, SbH_3, BiH_3}$), the key factor is the strength of the $\mathrm{E-H}$ bond. A weaker bond means the hydride can more readily release hydrogen, making it a stronger reducing agent.

As you go down the group from nitrogen to bismuth, the atomic size increases and the $\mathrm{E-H}$ bond length grows. This makes the bond weaker — the bond dissociation enthalpy decreases steadily. So $\mathrm{BiH_3}$ has the weakest $\mathrm{Bi-H}$ bonds and is the strongest reducing agent, while $\mathrm{NH_3}$ has the strongest $\mathrm{N-H}$ bonds and is the weakest.

Let’s work through the reasoning step by step.

1. Identify the trend in bond strength.

   Down group 15, the central atom’s size increases. The $\mathrm{E-H}$ bond becomes longer and therefore weaker. The bond dissociation enthalpy (energy needed to break one $\mathrm{E-H}$ bond) follows:

   $\mathrm{NH_3 > PH_3 > AsH_3 > SbH_3 > BiH_3}$.

2. Link bond strength to reducing ability.

   A reducing agent donates electrons or hydrogen. For these hydrides, the reaction often involves:

   $\mathrm{2EH_3 \rightarrow 2E + 3H_2}$

   or they reduce other compounds by losing hydrogen. The easier it is to break the $\mathrm{E-H}$ bond, the more readily the hydride acts as a reducing agent. So weaker bonds mean stronger reducing power.

3. Apply the trend.

   Since $\mathrm{BiH_3}$ has the weakest bonds, it is the strongest reducing agent. $\mathrm{SbH_3}$ comes next, then $\mathrm{PH_3}$, and finally $\mathrm{NH_3}$ is the weakest. This gives the order: …

The question asks us to evaluate an assertion and a reason, and then determine the relationship between them. This requires understanding the chemical properties of lithium and magnesium, particularly their diagonal relationship with elements of the next group, and how these properties influence the nature of their compounds and the physical properties of the metals themselves.

Concept: Diagonal Relationship, Covalent Character, and Metallic Hardness

1. Diagonal Relationship: Lithium (Li) in Group 1 and Magnesium (Mg) in Group 2 exhibit a diagonal relationship. This means they show similarities in properties with elements diagonally opposite to them in the periodic table (Li with Mg, and Be with Al). This similarity arises primarily due to their comparable ionic sizes and charge-to-radius ratios (charge density), leading to similar polarizing power.
2. Polarizing Power and Covalent Character:
   • Polarizing power is the ability of a cation to distort the electron cloud of an anion.
   • Small cations with high charge (high charge density) have high polarizing power.
   • $Li^+$ and $Mg^{2+}$ ions are relatively small and have high charge densities.
   • High polarizing power leads to a significant degree of covalent character in their compounds, even with highly electronegative elements like chlorine.
   • Covalent compounds are generally soluble in organic solvents (like ethanol) and less soluble in water, unlike purely ionic compounds.
3. Metallic Hardness:
   • The hardness of a metal is related to the strength of its metallic bonding.
   • Metallic bonding strength depends on factors like atomic size, the number of valence electrons, and the effectiveness of electron delocalization.
   • Generally, smaller atomic size and a greater number of valence electrons contribute to stronger metallic bonds and thus harder metals.

Step-by-step Evaluation:

1. Evaluate Assertion (A): LiCl and MgCl$_2$ are soluble in ethanol.

   • Lithium ($Li^+$) and Magnesium ($Mg^{2+}$) ions are very small and have high charge densities.
   • According to Fajan's rules, small cations with high charge have a high polarizing power. This means they can significantly distort the electron cloud of an anion (like $Cl^-$), leading to a considerable degree of covalent character in their compounds.
   • LiCl and MgCl$_2$ are not purely ionic; they possess significant covalent character due to the high polarizing power of $Li^+$ and $Mg^{2+}$ ions.
   • Covalent compounds are generally soluble in organic solvents such as ethanol.
   • Therefore, Assertion (A) is true.

2. Evaluate Reason (R): Lithium and magnesium are harder than their respective group elements.

   • For Group 1 (Alkali Metals): Lithium is the first element. As we move down the group (Li, Na, K, Rb, Cs), the atomic size increases, and the metallic bond strength decreases. This is because the valence electron is further from the nucleus and less tightly held, leading to weaker metallic bonding. Consequently, lithium is significantly harder than other alkali metals. …