Question

The fluorides of the second-period elements and their melting points are: \(\mathrm{LiF}, 845^{\circ} \mathrm{C} ; \mathrm{BeF}_{2}, 800^{\circ} \mathrm{C} ; \mathrm{BF}_{3}\), \(-126.7^{\circ} \mathrm{C} ; \mathrm{CF}_{4},-184^{\circ} \mathrm{C} ; \mathrm{NF}_{3},-206.6^{\circ} \mathrm{C} ; \mathrm{OF}_{2},\) \(-223.8^{\circ} \mathrm{C} ; \mathrm{F}_{2},-219.6^{\circ} \mathrm{C}\). Classify the type(s) of intermolecular forces present in each compound.

Short answer

LiF: ionic bonding; BeF2: covalent network (solids), dispersion; BF3: London dispersion; CF4: London dispersion; NF3: dipole-dipole, dispersion; OF2: dipole-dipole, dispersion; F2: London dispersion.

Step-by-step solution

Analyzing LiF

Lithium fluoride (LiF) is an ionic compound formed between lithium (a metal) and fluorine (a non-metal). Ionic compounds like LiF exhibit strong electrostatic forces between ions. Therefore, the primary intermolecular force present in LiF is ionic bonding.

Analyzing BeF2

Beryllium fluoride (BeF2) forms via covalent bonding, as beryllium shares electrons with fluorine atoms. In the solid state, it forms a network solid, similar to silica, with strong covalent bonds holding the network together. In the molten or gaseous state, dispersion forces and potential dipole-dipole interactions may be present.

Analyzing BF3

Boron trifluoride (BF3) is a covalent compound. It is non-polar due to its symmetrical planar shape, meaning the dipoles of the polar bonds cancel out. The primary intermolecular force in BF3 is London dispersion forces, which are weak.

Analyzing CF4

Carbon tetrafluoride (CF4) is a non-polar molecule due to its symmetric tetrahedral shape. Therefore, the only intermolecular forces present are London dispersion forces, which are weak.

Analyzing NF3

Nitrogen trifluoride (NF3) is a polar molecule because of its asymmetric trigonal pyramidal shape. The dipole-dipole interactions due to the polar N-F bonds are the primary intermolecular forces present, alongside the weaker London dispersion forces.

Analyzing OF2

Oxygen difluoride (OF2) has a bent shape, making it a polar molecule with dipole-dipole interactions as its primary intermolecular forces. It also exhibits London dispersion forces.

Analyzing F2

Fluorine (F2) is a diatomic molecule consisting of two identical fluorine atoms. Because it is non-polar, the primary type of intermolecular force is London dispersion forces.

Key concepts

Ionic Bonding

Ionic bonding occurs when there is a complete transfer of electrons from one atom to another. This typically happens between metal and non-metal atoms. In this type of bond, one atom (usually the metal) loses one or more electrons, becoming positively charged, while the other atom (the non-metal) gains those electrons, becoming negatively charged. The resulting oppositely charged ions attract each other due to electrostatic forces, creating a strong bond.

• Example: Lithium fluoride (LiF) forms ionic bonds.
• Characteristics: High melting and boiling points, typically form crystalline solids.
• Intermolecular Forces: Primarily strong electrostatic forces between ions, known as lattice energy.

This type of bonding is essential for the stability of many compounds and contributes significantly to their physical properties.

Covalent Bonding

Covalent bonding involves the sharing of electron pairs between atoms. It typically occurs between non-metal atoms which have similar electronegativities. Unlike ionic bonds, covalently bonded compounds form discrete molecules rather than a lattice structure.

• Example: Beryllium fluoride (BeF extsubscript{2}) and boron trifluoride (BF extsubscript{3}) exhibit covalent bonding.
• Bond Formation: Atoms share electrons to achieve a full outer electron shell, typically following the octet rule.
• Properties: These compounds can vary widely in their state (solid, liquid, gas) at room temperature, and generally have lower melting and boiling points compared to ionic compounds.

Covalent bonds are vital in organic chemistry and result in the formation of countless molecular structures.

London Dispersion Forces

London dispersion forces are the weakest of all intermolecular forces and are present in all molecules, whether polar or non-polar. They arise due to the temporary fluctuations in electron density within a molecule, which create temporary dipoles. These dipoles can induce a corresponding dipole in a nearby molecule, resulting in an attraction between them.

• Examples: Significant in boron trifluoride (BF extsubscript{3}), carbon tetrafluoride (CF extsubscript{4}), and fluorine (F extsubscript{2}).
• Characteristics: Increase in strength with larger molecular size or molar mass.
• Properties: These forces can account for the liquification of noble gases and the melting and boiling points of organic molecules.

Though weak, London dispersion forces are crucial for understanding interactions in non-polar compounds and noble gases.

Dipole-Dipole Interactions

Dipole-dipole interactions occur between polar molecules. These are molecules with a permanent dipole moment due to an unequal distribution of electrons, leading to partial positive and negative charges within different regions of the molecule. When polar molecules are close to each other, opposite charges align and attract.

• Examples: Notable in nitrogen trifluoride (NF extsubscript{3}) and oxygen difluoride (OF extsubscript{2}).
• Characteristics: Stronger than London dispersion forces but weaker than ionic and covalent bonds.
• Impact: Significantly affect boiling and melting points, making them higher compared to non-polar molecules of similar size and shape.

Understanding dipole-dipole interactions is essential for predicting the physical properties and reactivity of polar compounds.