Atomic Structure: Periodicity, Bonding, and Reactivity (26CMA102)

Flashcards testing periodic trends, ionisation energy, electron attachment, electronegativity scales, molecular orbital theory, bond orders, and hybridization from lecture 26CMA102.

What chemical equation defines first ionisation energy?

$$M(g) \rightarrow M^+(g) + e^-$$

Why does the first ionisation energy decrease when moving down Group 1?

The atomic radius increases down the group, placing the outer electron farther away from the nucleus. Consequently, the electrostatic force holding it weakens, making it easier to remove and requiring less energy.

What is the general trend in first ionisation energy across Period 2 from lithium to neon, and what causes it?

First ionisation energy generally increases because atomic radii decrease across the period, bringing outer electrons closer to the nucleus. The resulting stronger attractive force makes them harder to remove, requiring more energy.

Why is the first ionisation energy of boron lower than that of beryllium?

Ionisation of boron involves removing an electron from a $2p$ orbital rather than a $2s$ orbital. The $2p$ orbital is situated farther out from the nucleus and is therefore less tightly bound.

Why is the first ionisation energy of oxygen ($2s^2 2p^4$) lower than that of nitrogen ($2s^2 2p^3$)?

Ionising oxygen requires removing an electron from a doubly occupied $2p$ orbital, where electron-electron pairing repulsion makes the electron easier to remove compared to nitrogen's half-filled subshell.

What chemical equation defines electron attachment (gain) energy, and how is it related to electron affinity?

Electron attachment energy is the energy change for the process $M(g) + e^- \rightarrow M^-(g)$. Electron affinity is defined as the negative of this value.

Why does electron attachment energy increase (become less negative and less favourable) going down Group 1?

The radii of the elements increase down the group, meaning the incoming electron is placed farther away from the nucleus. The attractive nuclear force holding it is weaker, making addition less energetically favourable.

Why is electron attachment less favourable for beryllium than for lithium, and less favourable for nitrogen than for carbon?

For beryllium, the added electron must enter a $2p$ orbital rather than a $2s$ orbital (which is farther out). For nitrogen ($2s^2 2p^3$), the electron must enter an orbital already occupied by another electron, encountering electrostatic repulsion.

Why does neon have the highest (most unfavourable) electron attachment energy in Period 2?

Neon has a complete, filled shell ($2s^2 2p^6$), so adding an electron requires placing it into a new, higher energy shell.

How is electronegativity ($\chi$) defined?

The power of an atom to attract electrons to itself when part of a compound.

How did Robert Mulliken formulate electronegativity mathematically?

Electronegativity is estimated as $\frac{1}{2}(IE - EA)$, meaning elements with high ionisation energy and strongly negative electron attachment energy are the most electronegative.

What expression did Allred and Rochow propose to calculate electronegativity?

Electronegativity is proportional to the electric field at the surface of an atom: $\frac{Z_{\text{eff}}}{R^2}$, where $Z_{\text{eff}}$ is the effective nuclear charge and $R$ is the covalent radius.

According to Linus Pauling, why is an $X-Y$ bond stronger than the average of $X-X$ and $Y-Y$ bonds?

The bond has additional strength stemming from an electrostatic attraction between the partial charges ($\delta+$ and $\delta-$) created by the difference in electronegativity between $X$ and $Y$.

What is the key rule regarding the number of molecular orbitals formed in LCAO theory?

If $N$ atomic orbitals combine, exactly $N$ molecular orbitals are formed.

What famous phrase did physicist N. David Mermin use in 1989 regarding quantum mechanical wavefunctions?

"Shut up and calculate!"

What defines a sigma ($\sigma$) molecular orbital in terms of rotational symmetry?

It has cylindrical symmetry about the internuclear ($z$) axis, appearing unchanged regardless of the angle of rotation about that axis.

How is a pi ($\pi$) molecular orbital defined from the overlap of $p$ orbitals?

It is formed by side-on overlap where the wavefunction lies above and below the internuclear axis; a rotation of $180^\circ$ around the internuclear axis maps one phase onto the other.

What is the formula used to calculate bond order from a molecular orbital diagram?

$$\text{Bond Order} = \frac{\text{number of electrons in bonding orbitals} - \text{number of electrons in antibonding orbitals}}{2}$$

What are the calculated bond orders for $H_2$ and the interstellar cation $H_2^+$?

$H_2$ has a bond order of $1$, while $H_2^+$ has a bond order of $0.5$.

Why is the energy ordering of the $\sigma_p$ and $\pi$ molecular orbitals reversed in lighter diatomics ($Li_2$ through $N_2$) compared to $O_2$ and $F_2$?

In lighter diatomics, significant interaction ($s-p$ mixing) between the $2s$ and $2p$ atomic orbitals occurs, pushing the $\sigma_p$ orbital higher in energy than the $\pi$ orbitals.

What property of dioxygen ($O_2$) is demonstrated by liquid oxygen being attracted and suspended between the poles of a magnet as shown below?

Paramagnetism, which results from two unpaired electrons residing in separate degenerate $\pi^*$ antibonding orbitals in accordance with Hund's rule.

How do the bond lengths and bond orders compare across $O_2^+$, $O_2$, and $O_2^-$?

$O_2^+$ has a bond order of $2.5$ and length of $1.12\,\text{\AA}$; $O_2$ has a bond order of $2$ and length of $1.21\,\text{\AA}$; $O_2^-$ has a bond order of $1.5$ and length of $1.28\,\text{\AA}$. Higher bond order yields stronger and shorter bonds.

What creates a non-bonding interaction when atomic orbitals interact?

Zero net overlap occurs because positive (in-phase) overlap is entirely cancelled out by negative (out-of-phase) overlap.

Why was rhodium (Rh) essential in automotive catalytic converters for destroying nitric oxide ($NO$) while platinum (Pt) and palladium (Pd) were not?

On Pt and Pd surfaces, $NO$ donates electron density from an antibonding orbital to the metal, which strengthens the $N-O$ bond. Rh donates electron density into the antibonding orbital of $NO$, weakening the bond and facilitating its destruction.

What geometries and exemplar molecules correspond to $sp^3$, $sp^2$, and $sp$ orbital hybridization in carbon compounds?

$sp^3$ gives a tetrahedral shape (e.g., methane, $CH_4$); $sp^2$ gives a trigonal planar geometry with one unhybridized $p$ orbital forming a double bond (e.g., ethene, $CH_2=CH_2$); $sp$ gives a linear geometry with two unhybridized $p$ orbitals (e.g., alkynes).