In this article you will be able to understand the important role of intermolecular forces on solids, and how this determines many characteristics they have, and also how they look!
- What is a dipole moment?
- What is Solubility?
- Intermolecular Forces
- Properties of Solids, Liquids and Gases
- Phase Changes
What is a solid?
The solid is one of the three fundamental states of matter. Because the energy of atoms lowers when they occupy an organized, three-dimensional structure, a solid develops from a liquid or gas. Differentiating themselves from liquids and gases, solids have particular properties.For instance, all solids can withstand forces directed either perpendicularly or parallelly to a surface. The solid’s constituent atoms’ characteristics, their arrangement, and the forces that exist between them determine these characteristics.
Types of Solids
The particles that make up a solid, whether they are ionic, molecular, covalent, or metallic, experience powerful attraction forces that keep them in place. Consequently, when discussing solids, we shift the focus away from the movements of atoms, molecules, or ions, and instead, we emphasize their fixed locations in space. A solid can organize its elements in one of two ways: it can form an ordered, repeating three-dimensional structure known as a crystal lattice, thus creating a crystalline solid, or it can aggregate randomly, resulting in an amorphous solid.
Crystalline vs. Amorphous Solids
Two distinctive qualities are present in amorphous substances. When someone cleaves or fractures them, they form pieces with uneven, frequently curved surfaces, and they do not group their parts in a regular way, leading to poorly defined patterns in x-rays. People call an opaque, amorphous substance glass. If someone chills the liquid phase of nearly any material sufficiently quickly, it can solidify in an amorphous state. However, certain substances are naturally amorphous because either their constituent parts don’t fit together well enough to create a stable crystalline lattice or they include impurities that cause the lattice to break.
The particular internal structures of crystals, also known as crystalline solids, give rise to characteristic flat surfaces, or faces. The faces cross at angles that are typical of the material. Each structure also generates a unique pattern when exposed to x-rays, which may be used to identify the substance. The distinctive angles represent the regular repeating arrangement of the constituent atoms, molecules, or ions in space and are independent of the crystal’s size. Repulsive interactions, for instance, lead an ionic crystal to split along fixed planes, resulting in new faces that cross at the same angles as those in the original crystal. The angles at which the faces of a covalent solid, meet are likewise not random; rather, they are governed by how the atoms are arranged in the crystal.
Weak intermolecular forces, specially the Van Der Waals forces, hold together the atoms or molecules that make up molecular solids to form weak crystalline lattices. Each of the three different types of molecular solids correlates to one of the three primary Van Der Waals forces. The primary “point” of van der Waals forces is that a partial negative charge will attract a partial positive charge in order to produce them.
Since neutral molecules make up molecular solids, they are devoid of free electrons. As a result, electrons are poor conductors because they cannot move freely. Due to the weak intermolecular forces, molecular solids are also soft, easy to deform, and have low melting points. Also, since the “bonds” between molecules are lengthy, there is a lot of space between them, making them solids with low density.
One of two factors determines whether a species is non-polar: either the molecule is symmetrical, wiping out any polarity, or the difference in electronegativity is smaller than 0.5.
London dispersion forces hold together non-polar substances. The partly positive end of the instantaneous dipole is being drawn towards by the electrons in the non-polar species. The result is an unbalanced distribution of electrons that eventually forms a dipole.
Dipole-dipole interactions are the glue that holds together polar solids.
Solids with hydrogen bonds are kept together by these bonds. A specific kind of dipole-dipole interaction is hydrogen bonding. Due to the significant difference in electronegativity between H and N, O, or F, it is really considerably stronger.
A solid, which covalent bonds bind together, like a crystal or an amorphous solid, is a covalent network solid. In a network solid, the atoms link together in a continuous network. As a result, people can refer to the solid as a whole as a macromolecule, given that there are no individual molecules.
Covalent bonding within covalent network solids has a major role in how they behave. Which are: Hardness, elevated melting point, high or low conductivity and minimal solubility. This hardness results from covalent bonds, which are extremely powerful and challenging to break. Because it is challenging to dissolve network solids, their melting points are high. The type of bonding affects a network solid’s conductivity. Molecules holding their sheets together with intermolecular force exhibit high conductivity, while materials that are merely covalently bound show low conductivity.
Alternating ion lattices make up ionic solids. Because ions would have to come into contact, it is not possible to densely pack them.
The tight bonding of electrons to the ions frequently causes ionic compounds to be insulators. These compounds are dense and have high melting points due to the strong ionic bonding. Though they are brittle, the crystal will shatter if harmed or if the wrong ions come into touch. Ionic bonding is based on the fundamental concept of lattice energy. If you contrast the electron affinity for a non-metal with the ionization energy for a metal, you’ll find that producing ions is endothermic. Gases are the only type of gas when considering the definitions of electron affinities and ionization energies. The formation of ionic solids from the elements is exothermic because getting the ions into the crystal lattice produces heat.
We know them as metallic solids, which consist entirely of metal atoms held together by metallic bonds. A lattice of positive ions and delocalized electrons interact to form metallic bonds, a sort of intramolecular force of attraction.
Alkali metals, for example, have low melting points despite the fact that metallic solids generally have high melting temperatures. Each metal has a different melting point for metallic solids. Because their delocalized electrons may readily flow and transmit electrical charges, metallic substances are excellent electrical and thermal conductors. Metals are regarded as ductile and malleable substances. Because delocalized electrons may reflect light, metals are shiny.
Types of Solids Practice Problems
Identify the type of crystal lattice for each of the following substances:
- a) Sodium chloride (NaCl)
- b) Diamond (C)
- c) Iron (Fe)
- d) Graphite (C)
- e) Silicon (Si)
Explain the main difference between amorphous and crystalline solids. Provide one example of each type of solid.
Classify each of the following solids as either Ionic, Covalent, Metallic, or Molecular:
- a) Sodium chloride (NaCl)
- b) Diamond (C)
- c) Iron (Fe)
- d) Water (H2O)
- e) Copper (Cu)
Explain why metals are good conductors of electricity, referring to their electronic band structure and the behavior of free electrons.
Compare the melting point of a molecular solid with an ionic solid, assuming both have the same molecular weight. Justify your answer.
Types of Solids Practice Problem Solutions
- a) Sodium chloride (NaCl) – Ionic crystal lattice
- b) Diamond (C) – Covalent crystal lattice
- c) Iron (Fe) – Body-centered cubic (BCC) crystal lattice
- d) Graphite (C) – Hexagonal crystal lattice
- e) Silicon (Si) – Diamond cubic crystal lattice
Amorphous solids lack long-range order in their atomic arrangement, meaning their particles are arranged randomly without a defined crystal lattice. Crystalline solids, on the other hand, have a well-defined repeating pattern in their atomic arrangement, forming a crystal lattice. Example of amorphous solid: Glass. Example of crystalline solid: Sodium chloride (NaCl).
- a) Sodium chloride (NaCl) – Ionic solid
- b) Diamond (C) – Covalent solid
- c) Iron (Fe) – Metallic solid
- d) Water (H2O) – Molecular solid
- e) Copper (Cu) – Metallic solid
Metals are good conductors of electricity due to their unique electronic band structure. In metals, the valence and conduction bands overlap, resulting in the presence of free electrons that are not confined to individual atoms. These delocalized electrons are highly mobile and can move through the metal lattice with minimal resistance, enabling efficient electrical conductivity.
Molecular solids generally have lower melting points compared to ionic solids with the same molecular weight. This is because molecular solids are held together by relatively weak intermolecular forces, such as Van Der Waals forces or hydrogen bonding, which require less energy to overcome during melting. In contrast, ionic solids have strong electrostatic forces between ions, which necessitate higher energy input to break the ionic bonds and achieve melting.