(i) Explain why BeH2 molecule has zero dipole moment although the Be-H bonds are polar. [1 mark] (ii) By writing molecular orbital electronic configuration, compare the relative stability of the following species and indicate their magnetic properties: O2, O2+, O2 2- (Peroxide).

Question

(i) Explain why BeH2 molecule has zero dipole moment although the Be-H bonds are polar. [1 mark] (ii) By writing molecular orbital electronic configuration, compare the relative stability of the following species and indicate their magnetic properties: O2, O2+, O2 2- (Peroxide). [4 marks]

**OR**

(i) Describe the hybridisation in case of PCl5. Why are axial bonds longer as compared to equatorial bonds? [3 marks] (ii) Distinguish between Sigma (σ) and Pi (π) bond. [2 marks]

Accepted Answer

>

> [!TLDR]
> BeH2's linear (sp) geometry makes its two polar bond dipoles cancel exactly; comparing bonding vs antibonding electron counts for O2, O2+, and O2(2-) gives bond orders 2, 2.5, and 1 respectively, with O2 and O2+ paramagnetic and the peroxide ion diamagnetic.
>
> **(i) Why BeH$_2$ has zero net dipole moment despite polar Be-H bonds:**
>
> Beryllium (Be) has 2 valence electrons and forms 2 sigma bonds to the 2 hydrogen atoms in $BeH_2$. Be undergoes **sp hybridisation**, giving the molecule a perfectly **linear** shape ($H-Be-H$, bond angle $180°$).
>
> Each individual Be-H bond IS polar (Be and H have different electronegativities, so each bond has a small dipole moment vector pointing along that bond). However, because the molecule is linear and symmetric, the two bond dipole vectors are **equal in magnitude but point in exactly opposite directions** (180° apart). Vector addition of two equal-and-opposite vectors gives a **net resultant of zero**:
>
> $$\vec{\mu}_{net} = \vec{\mu}_1 + \vec{\mu}_2 = 0 \quad (	ext{since } \vec{\mu}_1 = -\vec{\mu}_2)$$
>
> So although the bonds themselves are polar, the overall molecular geometry cancels their effects, giving $BeH_2$ a net dipole moment of zero (a nonpolar molecule made of polar bonds).
>
> **(ii) MO electronic configurations, stability, and magnetism:**
>
> Molecular orbital filling order (for O$_2$-like diatomics, above N$_2$): $\sigma1s^2\ \sigma^*1s^2\ \sigma2s^2\ \sigma^*2s^2\ \sigma2p_z^2\ \pi2p_x^2{=}\pi2p_y^2\ \pi^*2p_x^{}\ \pi^*2p_y^{}$ (the last two, antibonding $\pi^*$, partially filled).
>
> - **$O_2$** (16 electrons): $KK\,\sigma2s^2\,\sigma^*2s^2\,\sigma2p_z^2\,\pi2p_x^2\,\pi2p_y^2\,\pi^*2p_x^1\,\pi^*2p_y^1$
>   - $N_b = 10, N_a = 6$; Bond order $= (10-6)/2 = 2$.
>   - Two unpaired electrons (one each in $\pi^*2p_x$, $\pi^*2p_y$) $\Rightarrow$ **paramagnetic**.
>
> - **$O_2^+$** (15 electrons, one less than O$_2$; removed from the highest-energy $\pi^*$ orbital): configuration same as O$_

Q.(i) Explain why BeH2 molecule has zero dipole moment although the Be-H bonds are polar. [1 mark]

Concept understanding — Bond Order

Bond Order: From Intuition to Precision

The Precise Definition

What Bond Order Tells You

A Quick Example: $\text{H}_2$

Another Example: $\text{He}_2$ …

Q.(i) Explain why BeH2 molecule has zero dipole moment although the Be-H bonds are polar. [1 mark]

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