Have you ever used Epsom salt to help with muscle soreness and stress? Epsom salt is made up of magnesium and sulfate ions that join together to form magnesium sulfate (MgSO4), and it is believed that Epsom salt has many benefits, from helping with muscle pain to relieving anxiety! Epsom salt (MgSO4) is a type of crystalline solid, and more specifically, it is considered an ionic solid.
Now, you might be wondering what exactly ionic solids are and what their properties are. So, let's dive into the properties of ionic solids!
What is the definition of an Ionic Solid?
Solids are divided into crystalline and amorphous solids based on their particle arrangement.
Crystalline solids have a highly organized arrangement of particles in a 3D structure. In crystalline solids, all bonds between particles have equal strength. Crystalline solids also have distinct melting points.
Amorphous solids lack a structured arrangement of particles. Particles are randomly arranged and melt over a range of temperatures.
Crystalline solids can be further divided into ionic solids, molecular solids, covalent network solids, and metallic solids.
If you want to learn more about the difference between these types of solids, read the article "Properties of Solids".
In this article, we will focus solely on ionic solids. For starters, let's define ionic solids!
Ionic solids are made up of ions joined together by ionic bonds. "Ionic bonding" is a type of chemical bond between a positive and a negatively charged ion where the transfer of electrons occurs.
Let's look at our first example!
Which is the following compounds is considered an ionic solid? Ag, SO3 and CaO.
We defined ionic solids are compounds that contain a cation and an anion. Ag is by itself so it cannot be a compound. SO3 is made up of two nonmetals. CaO is made up of a metal cation (Ca2+) and a nonmetal anion (O2-). So, CaO is the ionic solid in this example.
Structure of Ionic solids
Because ionic solids are considered a type of crystalline solid, they have a well-structured, 3D arrangement of particles which we call a crystal lattice. The basic structure of a crystal lattice is shown below.
Basic crystal lattice structure, Isadora Santos - Vaia Originals
You can learn more about the structure of solids in the article "Solids"!
In ionic solids, the metal cations have smaller sizes compared to the nonmetal anions. So, as the anions line up, the cations accommodate in a way that they surround the anions. This arrangement helps to maximize the electrostatic attraction between the ions. The easiest way to understand this is by looking at the crystal lattice structure of potassium chloride.
Actually, there are a couple of exemptions to this rule. In K, Rb, Cs fluorides and, Rb, Cs oxides the anion is the bigger one. Of course, this does not matter since then the anions will arrange to accommodate the cations.
Furthermore, if the metal cation is not one metal atom but multiple ones (Hg2+2) joined together or have different atoms coordinated it will vary in size vastly. And just to complicate things, there are cations that are not metals, for example, NH4+.
When we analyze the structure of potassium chloride (KCl) we know that potassium ion (K) has a charge of +1 so we called it the cation, whereas chlorine ion has a charge of -1 and is called the anion. So, in the crystal lattice structure of KCl, the cation surrounds the anion on all sides. After that comes the next "layer" (or cube to be specific) of anions, then cations, and so on to build a crystal.
Crystal Lattice Structure of KCl, Isadora Santos - Vaia Originals.
Properties and Characteristics of Ionic Solids
When describing the properties of ionic solids, we need to consider the following characteristics: melting point, hardness, conductivity, solubility, lattice energy, and strength of electrostatic interactions.
Electrostatic interactions
As we saw before, ionic solids are made up of ions held together by ionic bonds. These ionic bonds are considered a strong electrostatic interaction, and it affects the melting point of ionic solids. Having strong ionic bonds means that a lot of kinetic energy is needed to break the bond between ions.
The charges of ions also determine the strength of the attractive forces. So, the higher the charges, the stronger the electrostatic forces holding the ions together.
Think about the charges of MgO and NaCl. Mg has a +2 charge, and O has a -2 charge, whereas Na has a charge of +1 and Cl has a charge of -1. So, based on the charge, we can say that MgO will have stronger attractive forces, and therefore a higher melting point.
This is actually true, as the melting point of magnesium oxide (MgO) and sodium chloride is 2852 °C and 801°C respectively.
Did you know that ionic bonds are related to Coulombs Law? According to Coulombs Law, the strength of the ionic bond is directly proportional to the charges on the ions. Basically, the higher the charge of ions, the stronger the attraction between them and the larger the coulombic forces.
Melting point and Hardness of ionic solids
Because of the strength of their ionic bonds, ionic solids have very high melting points. Ionic solids are also considered very hard because of the forces of attraction between ions. However, ionic solids are brittle, meaning that their crystal structure is easily shattered into pieces.
The conductivity of ionic solids
Are ionic solids able to conduct electricity? The answer is yes, but only when their ions are mobile, which occurs when an ionic solid is molten or dissolved in an aqueous solution!
Conductivity is the ability of a compound to conduct electricity.
Electrolytes are referred to as compounds whose aqueous solutions or molten state is able to conduct electricity. Ionic solids are considered strong electrolytes.
When ionic solids are heated until in their molten (liquid) state, they become good conductors of electricity because, in this state, the ions are mobile. Aqueous solutions of ionic solids are also good at conducting electricity because the presence of ions in the aqueous solution allows electricity to pass through!
Solubility of ionic compounds
Remember the rule: Like dissolved like! So, Ionic solids can dissolve in polar solvents, such as water.
Solubility is referred to as the ability of a solute to dissolve in a solvent to form a solution.
However, keep in mind that not all ionic solids are soluble. To know whether a compound will be soluble or insoluble, let's review the solubility rules for ionic compounds.
|
| Ionic compounds containing NO3- and CH3COO- are soluble. |
| Ionic compounds containing Cl-, Br- and I- are soluble, except for when Ag+, Hg22+, and Pb2+ are present. |
| Ionic compounds containing SO42- are soluble, except for Sr2+, Ba2+, Hg22+, and Pb2+. |
Ionic compounds containing S2- are insoluble, except for NH4+, Li+, Na+, K+, Rb+, Cs+, Ca2+, Sr2+, and Ba2+. |
Ionic compounds containing CO32- are insoluble, except for NH4+, Li+, Na+, K+, Rb+, and Cs+. |
Ionic compounds containing CO32- are insoluble, except for NH4+, Li+, Na+, K+, Rb+, and Cs+. |
Ionic compounds containing PO43- are insoluble, except for NH4+, Li+, Na+, K+, Rb+, and Cs+. |
Ionic compounds containing OH- are insoluble, except for NH4+, Li+, Na+, K+, Rb+, Cs+, Ca2+, Sr2+, and Ba2+. |
The temperature and strength of the ionic bonds also affect the solubility of solids. The temperature and strength of the ionic bonds are directly proportional to the solubility of solids.
Lattice energy
The lattice structure of ionic solids consists of a well-arranged pattern, with cations surrounding the nonmetal anion. The stability of the crystal lattice depends on the strength of the attractive forces between the oppositely charged ions. The stronger the attraction, the higher the stability of the crystal lattice and the greater the lattice energy.
Lattice energy is defined as the energy released when an ionic solid is formed from its ions.
Lattice energy can be used to estimate the strength of ionic bonds. The greater the lattice energy, the stronger the ionic bond.
The lattice energy of ionic solids depends on the charge of the ions and on the size of the ions.
The greater the charge of the ions, the greater the lattice energy (⇧ charge of the ions = ⇧ lattice energy)
The smaller the size of the ions, the greater the lattice energy ( ⇩ size of the ions = ⇧ lattice energy)
Let's look at an example!
Which of the following ionic solids would have the largest lattice energy? CsBr, LiI, ZnO
First, look at the charge of the ions present in each ionic compound:
CsBr: Cs+1 and Br -1
LiI: Li+1 and I-1
ZnO: Zn+2 and O-2
The higher the charges, the higher the lattice energy. So, the ionic solid with the largest lattice energy would be ZnO.
Now, if two ionic solids have the same charge, how would you determine which would have the greatest lattice energy? You would have to look at the ionic radius of each ion. The ionic compound with the smallest ion size will have the greatest amount of lattice energy. The trend for ionic radius is that ionic radius increases as the number of electrons increases. The periodic trend for ionic radius is shown below:
Which of the following compounds has the largest lattice energy? NaF, NaCl, NaBr or NaI?
By looking at the periodic table, we can see that F- has the smallest ionic radius. So, NaF has the highest lattice energy.
Now, you should be able to recognize ionic solids, their basic structure, and also some of their properties!
Ionic Solids - Key takeaways
References
- Swanson, J. W. (2020). Everything you need to Ace Chemistry in one big fat notebook. Workman Pub.
- Malone, L. J., Dolter, T. O., & Gentemann, S. (2013). Basic concepts of Chemistry (8th ed.). Hoboken, NJ: John Wiley & Sons.
- Moore, J. T., & Langley, R. (2021). McGraw Hill: AP Chemistry, 2022. New York: McGraw-Hill Education.
- Salazar, E., Sulzer, C., Yap, S., Hana, N., Batul, K., Chen, A., . . . Pasho, M. (n.d.). Chad's general chemistry Master course. Retrieved May 4, 2022, from https://courses.chadsprep.com/courses/general-chemistry-1-and-2
- AP Chemistry course and exam description, effective fall 2020. (n.d.). Retrieved May 28, 2022, from https://apcentral.collegeboard.org/pdf/ap-chemistry-course-and-exam-description.pdf?course=ap-chemistry
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