Ionic Solids: Meaning, Examples & Properties | StudySmarter

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Ionic Solids

Ionic Solids

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.

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True or false: Solids are divided into crystalline and amorphous solids based on their particle arrangement.

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______ have a highly organized arrangement of particles in a 3D structure.

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______ lack a structured arrangement of particles. Particles are randomly arranged and melt over a range of temperatures.

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Ionic solids are a type of ______ solid.

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Ionic solids are made up of ions joined together by _____ bonds.

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________ is a bond between a positive and a negatively charged ion where the transfer of electrons occurs.

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In the crystal lattice structure of ionic solids, the _____ surrounds the bigger _____.

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The stronger the attractive forces between ions, the _____ the melting point of ionic solids.

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The higher the charges, the ______ the electrostatic forces holding the ions together.

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True or false: ionic solids have low melting points.

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______ is the ability of a compound to conduct electricity.

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True or false: Solids are divided into crystalline and amorphous solids based on their particle arrangement.

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______ have a highly organized arrangement of particles in a 3D structure.

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______ lack a structured arrangement of particles. Particles are randomly arranged and melt over a range of temperatures.

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Ionic solids are a type of ______ solid.

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Ionic solids are made up of ions joined together by _____ bonds.

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________ is a bond between a positive and a negatively charged ion where the transfer of electrons occurs.

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In the crystal lattice structure of ionic solids, the _____ surrounds the bigger _____.

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The stronger the attractive forces between ions, the _____ the melting point of ionic solids.

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The higher the charges, the ______ the electrostatic forces holding the ions together.

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Last Updated: 28.05.2022
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10 min reading time

Chemical Analysis
Chemical Reactions
Chemistry Branches
Inorganic Chemistry
Ionic and Molecular Compounds
Kinetics
Making Measurements
Nuclear Chemistry
Organic Chemistry
Physical Chemistry

Absolute Entropy And Entropy Change
Acid Dissociation Constant
Acid-Base Indicators
Acid-Base Reactions and Buffers
Acids and Bases
Alkali Metals
Allotropes of Carbon
Amorphous Polymer
Amount of Substance
Application of Le Chatelier's Principle
Arrhenius Equation
Arrhenius Theory
Atom Economy
Atomic Structure
Autoionization of Water
Avogadro Constant
Avogadro's Number and the Mole
Beer-Lambert Law
Bond Enthalpy
Bonding
Born Haber Cycles
Born-Haber Cycles Calculations
Boyle's Law
Brønsted-Lowry Acids and Bases
Buffer Capacity
Buffer Solutions
Buffers
Buffers Preparation
Calculating Enthalpy Change
Calculating Equilibrium Constant
Calorimetry
Carbon Structures
Cell Potential
Cell Potential and Free Energy
Chalcogens
Chemical Calculations
Chemical Equations
Chemical Equilibrium
Chemical Thermodynamics
Closed Systems
Colligative Properties
Collision Theory
Common-Ion Effect
Composite Materials
Composition of Mixture
Constant Pressure Calorimetry
Constant Volume Calorimetry
Coordination Compounds
Coupling Reactions
Covalent Bond
Covalent Network Solid
Crystalline Polymer
De Broglie Wavelength
Determining Rate Constant
Deviation From Ideal Gas Law
Diagonal Relationship
Diamond
Dilution
Dipole Chemistry
Dipole Moment
Dissociation Constant
Distillation
Dynamic Equilibrium
Electric Fields Chemistry
Electrochemical Cell
Electrochemical Series
Electrochemistry
Electrode Potential
Electrolysis
Electrolytes
Electromagnetic Spectrum
Electron Affinity
Electron Configuration
Electron Shells
Electronegativity
Electronic Transitions
Elemental Analysis
Elemental Composition of Pure Substances
Empirical and Molecular Formula
Endothermic and Exothermic Processes
Energetics
Energy Diagrams
Enthalpy Changes
Enthalpy For Phase Changes
Enthalpy of Formation
Enthalpy of Reaction
Enthalpy of Solution and Hydration
Entropy
Entropy Change
Equilibrium Concentrations
Equilibrium Constant Kp
Equilibrium Constants
Examples of Covalent Bonding
Factors Affecting Reaction Rates
Finding Ka
Free Energy
Free Energy Of Dissolution
Free Energy and Equilibrium
Free Energy of Formation
Fullerenes
Fundamental Particles
Galvanic and Electrolytic Cells
Gas Constant
Gas Solubility
Gay Lussacs Law
Giant Covalent Structures
Graham's Law
Graphite
Ground State
Group 3A
Group 4A
Group 5A
Half Equations
Heating Curve for Water
Heisenberg Uncertainty Principle
Henderson-Hasselbalch Equation
Hess' Law
Hybrid Orbitals
Hydrogen Bonds
Ideal Gas Law
Ideal and Real Gases
Intermolecular Forces
Introduction to Acids and Bases
Ion And Atom Photoelectron Spectroscopy
Ion dipole Forces
Ionic Bonding
Ionic Product of Water
Ionic Solids
Ionisation Energy
Ions: Anions and Cations
Isotopes
Kinetic Molecular Theory
Lattice Structures
Law of Definite Proportions
Le Chatelier's Principle
Lewis Acid and Bases
London Dispersion Forces
Magnitude Of Equilibrium Constant
Mass Spectrometry
Mass Spectrometry of Elements
Maxwell-Boltzmann Distribution
Measuring EMF
Mechanisms of Chemical Bonding
Melting and Boiling Point
Metallic Bonding
Metallic Solids
Metals Non-Metals and Metalloids
Mixtures and Solutions
Molar Mass Calculations
Molarity
Molecular Orbital Theory
Molecular Solid
Molecular Structures of Acids and Bases
Moles and Molar Mass
Nanoparticles
Neutralisation Reaction
Oxidation Number
Partial Pressure
Particulate Model
Partition Coefficient
Percentage Yield
Periodic Table Organization
Phase Changes
Phase Diagram of Water
Photoelectric Effect
Photoelectron Spectroscopy
Physical Properties
Polarity
Polyatomic Ions
Polyprotic Acid Titration
Prediction of Element Properties Based on Periodic Trends
Pressure and Density
Properties Of Equilibrium Constant
Properties of Buffers
Properties of Solids
Properties of Water
Quantitative Electrolysis
Quantum Energy
Quantum Numbers
RICE Tables
Rate Equations
Rate of Reaction and Temperature
Reacting Masses
Reaction Quotient
Reaction Quotient And Le Chateliers Principle
Real Gas
Redox
Relative Atomic Mass
Representations of Equilibrium
Reversible Reaction
SI units chemistry
Saturated Unsaturated and Supersaturated
Shapes of Molecules
Shielding Effect
Simple Molecules
Solids Liquids and Gases
Solubility
Solubility Curve
Solubility Equilibria
Solubility Product
Solubility Product Calculations
Solutes Solvents and Solutions
Solution Representations
Solutions and Mixtures
Specific Heat
Spectroscopy
Standard Potential
States of Matter
Stoichiometry In Reactions
Strength of Intermolecular Forces
The Laws of Thermodynamics
The Molar Volume of a Gas
Thermodynamically Favored
Trends in Ionic Charge
Trends in Ionisation Energy
Types of Mixtures
VSEPR Theory
Valence Electrons
Van der Waals Forces
Vapor Pressure
Water in Chemical Reactions
Wave Mechanical Model
Weak Acid and Base Equilibria
Weak Acids and Bases
Writing Chemical Formulae
pH
pH Change
pH Curves and Titrations
pH Scale
pH and Solubility
pH and pKa
pH and pOH

The Earths Atmosphere

Contents

Chemical Analysis
Chemical Reactions
Chemistry Branches
Inorganic Chemistry
Ionic and Molecular Compounds
Kinetics
Making Measurements
Nuclear Chemistry
Organic Chemistry
Physical Chemistry

Absolute Entropy And Entropy Change
Acid Dissociation Constant
Acid-Base Indicators
Acid-Base Reactions and Buffers
Acids and Bases
Alkali Metals
Allotropes of Carbon
Amorphous Polymer
Amount of Substance
Application of Le Chatelier's Principle
Arrhenius Equation
Arrhenius Theory
Atom Economy
Atomic Structure
Autoionization of Water
Avogadro Constant
Avogadro's Number and the Mole
Beer-Lambert Law
Bond Enthalpy
Bonding
Born Haber Cycles
Born-Haber Cycles Calculations
Boyle's Law
Brønsted-Lowry Acids and Bases
Buffer Capacity
Buffer Solutions
Buffers
Buffers Preparation
Calculating Enthalpy Change
Calculating Equilibrium Constant
Calorimetry
Carbon Structures
Cell Potential
Cell Potential and Free Energy
Chalcogens
Chemical Calculations
Chemical Equations
Chemical Equilibrium
Chemical Thermodynamics
Closed Systems
Colligative Properties
Collision Theory
Common-Ion Effect
Composite Materials
Composition of Mixture
Constant Pressure Calorimetry
Constant Volume Calorimetry
Coordination Compounds
Coupling Reactions
Covalent Bond
Covalent Network Solid
Crystalline Polymer
De Broglie Wavelength
Determining Rate Constant
Deviation From Ideal Gas Law
Diagonal Relationship
Diamond
Dilution
Dipole Chemistry
Dipole Moment
Dissociation Constant
Distillation
Dynamic Equilibrium
Electric Fields Chemistry
Electrochemical Cell
Electrochemical Series
Electrochemistry
Electrode Potential
Electrolysis
Electrolytes
Electromagnetic Spectrum
Electron Affinity
Electron Configuration
Electron Shells
Electronegativity
Electronic Transitions
Elemental Analysis
Elemental Composition of Pure Substances
Empirical and Molecular Formula
Endothermic and Exothermic Processes
Energetics
Energy Diagrams
Enthalpy Changes
Enthalpy For Phase Changes
Enthalpy of Formation
Enthalpy of Reaction
Enthalpy of Solution and Hydration
Entropy
Entropy Change
Equilibrium Concentrations
Equilibrium Constant Kp
Equilibrium Constants
Examples of Covalent Bonding
Factors Affecting Reaction Rates
Finding Ka
Free Energy
Free Energy Of Dissolution
Free Energy and Equilibrium
Free Energy of Formation
Fullerenes
Fundamental Particles
Galvanic and Electrolytic Cells
Gas Constant
Gas Solubility
Gay Lussacs Law
Giant Covalent Structures
Graham's Law
Graphite
Ground State
Group 3A
Group 4A
Group 5A
Half Equations
Heating Curve for Water
Heisenberg Uncertainty Principle
Henderson-Hasselbalch Equation
Hess' Law
Hybrid Orbitals
Hydrogen Bonds
Ideal Gas Law
Ideal and Real Gases
Intermolecular Forces
Introduction to Acids and Bases
Ion And Atom Photoelectron Spectroscopy
Ion dipole Forces
Ionic Bonding
Ionic Product of Water
Ionic Solids
Ionisation Energy
Ions: Anions and Cations
Isotopes
Kinetic Molecular Theory
Lattice Structures
Law of Definite Proportions
Le Chatelier's Principle
Lewis Acid and Bases
London Dispersion Forces
Magnitude Of Equilibrium Constant
Mass Spectrometry
Mass Spectrometry of Elements
Maxwell-Boltzmann Distribution
Measuring EMF
Mechanisms of Chemical Bonding
Melting and Boiling Point
Metallic Bonding
Metallic Solids
Metals Non-Metals and Metalloids
Mixtures and Solutions
Molar Mass Calculations
Molarity
Molecular Orbital Theory
Molecular Solid
Molecular Structures of Acids and Bases
Moles and Molar Mass
Nanoparticles
Neutralisation Reaction
Oxidation Number
Partial Pressure
Particulate Model
Partition Coefficient
Percentage Yield
Periodic Table Organization
Phase Changes
Phase Diagram of Water
Photoelectric Effect
Photoelectron Spectroscopy
Physical Properties
Polarity
Polyatomic Ions
Polyprotic Acid Titration
Prediction of Element Properties Based on Periodic Trends
Pressure and Density
Properties Of Equilibrium Constant
Properties of Buffers
Properties of Solids
Properties of Water
Quantitative Electrolysis
Quantum Energy
Quantum Numbers
RICE Tables
Rate Equations
Rate of Reaction and Temperature
Reacting Masses
Reaction Quotient
Reaction Quotient And Le Chateliers Principle
Real Gas
Redox
Relative Atomic Mass
Representations of Equilibrium
Reversible Reaction
SI units chemistry
Saturated Unsaturated and Supersaturated
Shapes of Molecules
Shielding Effect
Simple Molecules
Solids Liquids and Gases
Solubility
Solubility Curve
Solubility Equilibria
Solubility Product
Solubility Product Calculations
Solutes Solvents and Solutions
Solution Representations
Solutions and Mixtures
Specific Heat
Spectroscopy
Standard Potential
States of Matter
Stoichiometry In Reactions
Strength of Intermolecular Forces
The Laws of Thermodynamics
The Molar Volume of a Gas
Thermodynamically Favored
Trends in Ionic Charge
Trends in Ionisation Energy
Types of Mixtures
VSEPR Theory
Valence Electrons
Van der Waals Forces
Vapor Pressure
Water in Chemical Reactions
Wave Mechanical Model
Weak Acid and Base Equilibria
Weak Acids and Bases
Writing Chemical Formulae
pH
pH Change
pH Curves and Titrations
pH Scale
pH and Solubility
pH and pKa
pH and pOH

The Earths Atmosphere

Contents

Fact Checked Content
Last Updated: 28.05.2022
10 min reading time

Content creation process designed by Lily Hulatt
Content sources cross-checked by Gabriel Freitas
Content quality checked by Gabriel Freitas

Jump to a key chapter

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.

Ionic solids Crystal lattice structure Ionic solid charactestics StudySmarter Basic crystal lattice structure, Isadora Santos - StudySmarter 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.

Ionic solids crystal lattice structure of KCl Ionic solids examples StudySmarter Crystal Lattice Structure of KCl, Isadora Santos - StudySmarter 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 stronger the attractive forces between ions, the higher the melting point of ionic solids.

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.

Solubility Rules

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 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

Crystalline solids have a highly organized arrangement of particles in a 3D structure. Ionic solids are a type of crystalline solid.
Ionic solids are made up of ions joined together by ionic bonds. Ionic bonds are bonds between a positive and a negatively charged ion where the transfer of electrons occurs.
The stronger the attractive forces between ions, the higher the melting point of ionic solids.
Ionic solids can only conduct electricity as molten ionic compounds or in aqueous solutions.
The lattice energy of ionic solids depends on the charge of the ions and on the size of the ions.
(⇧ charge of the ions = ⇧ lattice energy, ⇩ size of the ions = ⇧ lattice energy)

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

Flashcards in Ionic Solids

18 Start learning

True or false: Solids are divided into crystalline and amorphous solids based on their particle arrangement.

True

______ have a highly organized arrangement of particles in a 3D structure.

Crystalline solids

______ lack a structured arrangement of particles. Particles are randomly arranged and melt over a range of temperatures.

Amorphous solids

Ionic solids are a type of ______ solid.

Crystalline

Ionic solids are made up of ions joined together by _____ bonds.

Ionic

________ is a bond between a positive and a negatively charged ion where the transfer of electrons occurs.

Ionic bonding

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Frequently Asked Questions about Ionic Solids

What are ionic solids?

Ionic solids are made up of ions joined 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.

What are the physical characteristics of ionic solids?

The characteristics of ionic solids are high melting points, very hard, brittle, and strong forces of attraction.

What are the structure and properties of ionic solids?

The structure of ionic solids consists of a well-arranged pattern of cations and anions being held together by ionic bonds. The properties of ionic solids are high melting points, very hard, brittle, and strong electrostatic interactions.

What is the melting point of Ionic Solids?

Ionic solids have very high melting points because of the strength of the electrostatic interactions between ions.

What are examples of Ionic Solids?

Some common examples of ionic solids are NaCl, LiI, CsBr, and MgO.

Save Article Test your knowledge with multiple choice flashcards

True or false: Solids are divided into crystalline and amorphous solids based on their particle arrangement.

A. False B. True

______ have a highly organized arrangement of particles in a 3D structure.

A. Crystalline solids B. Amorphous solids

______ lack a structured arrangement of particles. Particles are randomly arranged and melt over a range of temperatures.

A. Amorphous solids B. Crystalline solids

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Tag » What Is An Ionic Solid

Ionic Solids: Meaning, Examples & Properties | StudySmarter

What is the definition of an Ionic Solid?

Structure of Ionic solids

Properties and Characteristics of Ionic Solids

Electrostatic interactions

Melting point and Hardness of ionic solids

The conductivity of ionic solids

Solubility of ionic compounds

Lattice energy

Ionic Solids - Key takeaways

References

Similar topics in Chemistry

Related topics to Physical Chemistry

Flashcards in Ionic Solids

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Frequently Asked Questions about Ionic Solids

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