Classify each substance in the table as either a metallic, ionic, molecular, or covalent network solid: $$\begin{array}{|c|c|c|c|c|} \hline \text { Substance } & \text { Applatence } & \text { Mature Polut } & \text { Ecetuct Conductive } & \text { Solusing a water } \\ \hline \mathrm{x} & \text { Iustrous, malleable } & 1500^{\circ} \mathrm{C} & \text { high } & \text { insoluble } \\ \hline \mathrm{Y} & \text { sott, yellow } & 113^{\circ} \mathrm{C} & \text { none } & \text { insoluble } \\ \hline \mathrm{z} & \text { hard, white } & 800^{\circ} \mathrm{C} & \text { only if melted/dissolved } & \text { soluble } \\ \hline \end{array}$$

Imagine holding a piece of shiny metal. What you're feeling is the cool touch of a metallic solid. These solids showcase properties that distinguish them from other types of materials. Metallic solids are composed of closely packed metal atoms that share a 'sea of electrons'.

These electrons are not bound to any specific atom, allowing them to move freely. This mobility is what grants metallic solids their characteristic electrical conductivity and thermal conductivity. They are generally malleable and can be shaped or bent without breaking. A high melting point is also typical, reflecting the strong bonds between the metal ions and the electrons.

For students trying to identify metallic solids, look for substances with a lustrous shine, high electrical conductivity, and the ability to be deformed without breaking.

When you scoop up a pile of sugar or snow, you're interacting with molecular solids. These substances consist of atoms or molecules held together by intermolecular forces such as hydrogen bonds, dipole-dipole interactions, or London dispersion forces.

Characteristics of Molecular Solids

These forces are relatively weak compared to other types of chemical bonds, which leads to a lower melting point for molecular solids, making them often soft. An example is ice, which melts at 0°C.

In terms of electrical conductivity, molecular solids typically do not conduct electricity because they lack free charges—ions or electrons—that can move throughout the substance. This trait helps in distinguishing them from metallic or ionic solids, which do conduct electricity under certain conditions.

Have you ever licked a salt crystal? That tangy taste comes from an ionic solid. Ionic solids are comprised of a lattice of positive and negative ions, such as sodium ions (Na+) and chloride ions (Cl-), held together by ionic bonds.

Understanding Ionic Solids

This bond forms when an atom with a low ionization energy loses an electron to an atom with a high electron affinity, creating ions that attract each other. Ionic solids have characteristic high melting points due to the strong interactions between the oppositely charged ions.

While they do not conduct electricity in their solid state, they become conductive when melted or dissolved in water, as the ions are free to move. Furthermore, many ionic solids are soluble in water due to the attraction between the ions and the polar water molecules.

Diamonds, the strongest natural material on earth, are covalent network solids. These solids are distinguished because their atoms are joined by a continuous network of covalent bonds.

Properties of Covalent Network Solids

This extensive bonding network extends throughout the material, resulting in very hard substances with high melting points. These solids usually do not conduct electricity, as all electrons are localized in covalent bonds rather than being free to move as in metallic solids.

Other well-known covalent network solids include graphite and silicon carbide, materials that are both tough and heat resistant - perfect for high-stress industrial applications. When dealing with substances exhibiting extreme hardness and temperature resistance, it's likely you're dealing with a network solid.

Chemical bonding is the very backbone of the material world, dictating the properties of different solids. There are four primary types of chemical bonds:

1. Covalent bonds, which involve sharing of electron pairs between atoms.
2. Ionic bonds, formed between positively and negatively charged ions.
3. Metallic bonds, featuring delocalized electrons in a sea of electron cloud around metal ions.
4. Intermolecular forces, weaker bonds such as hydrogen bonds or van der Waals forces in molecular solids.

A firm grasp of these bond types allows us to predict and explain the properties of a wide array of materials, from the water that we drink to the metals used in our technology.

Solids exhibit a variety of different properties based on the type of bonding holding them together. Understanding these properties can help us choose the right material for the right application.

Exploring the Range of Solid Properties

Melting point, malleability, hardness, electrical and thermal conductivity, and solubility are key characteristics to consider. A metal’s malleability contrasts with the brittleness of an ionic solid. The impressive electrical conductivity of metallic solids contrasts with the insulating nature of covalent network solids.

By examining these properties, we can classify unknown solids and predict how they might behave in various environments. This understanding is essential for professionals ranging from engineers designing spacecraft to chefs looking for the right packaging to keep food fresh.