Types of Crystalline Solids

Summary

This educational video script delves into the fascinating world of crystalline solids, categorizing them into four types: ionic, covalent network, metallic, and molecular. Ionic solids, exemplified by sodium chloride, feature strong ionic bonds within a crystal lattice, leading to high melting points and brittleness. Covalent network solids, like diamond and quartz, consist of atoms bonded by covalent bonds, resulting in hardness and high melting points. Metallic solids, characterized by a 'sea of electrons' among metal atoms, are known for their conductivity and malleability. Lastly, molecular solids, held together by weaker intermolecular forces, are softer with lower melting points and are non-conductive. The script provides a comprehensive understanding of these solids' properties and behaviors.

Takeaways

• 🔬 Crystalline solids are categorized into four main types: ionic solids, covalent network solids, metallic solids, and molecular solids.
• 🧲 Ionic solids, like sodium chloride, have a crystal lattice structure where ions are held together by strong ionic bonds, resulting in high melting points and brittleness.
• 💧 Covalent network solids are characterized by a giant 3-dimensional array of covalently bonded atoms, which makes them hard and with high melting points, but generally insoluble in solvents.
• 💠 Metallic solids consist of metal atoms bonded by metallic bonds, which allows for high electrical conductivity, malleability, and high melting points.
• 🌡️ Molecular solids are held together by weak intermolecular forces, leading to low melting points, softness, and non-conductivity of electricity.
• 🔍 The smallest repeatable unit in a crystal lattice is known as the unit cell, which varies depending on the ions or atoms involved.
• 🧊 Solubility of ionic compounds in water is due to their ability to form ion-dipole interactions with water molecules.
• 🌐 Covalent network solids, such as diamond and graphite, are allotropes of carbon with different bonding structures, leading to distinct properties.
• 🌟 Metallic bonds involve delocalized valence electrons that contribute to the metallic properties like high electrical and thermal conductivity.
• 🌀 Allotropes are different structural forms of the same element, with oxygen having O2 and O3 as examples, and carbon having diamond and graphite as notable allotropes.

Q & A

What are the two broad categories of solids mentioned in the video?

The two broad categories of solids mentioned are amorphous solids, where atoms are arranged randomly, and crystalline solids, where atoms are arranged in an orderly and repeated fashion.

What are the four main types of crystalline solids discussed in the video?

The four main types of crystalline solids discussed are ionic solids, covalent network solids, metallic solids, and molecular solids.

How are ionic solids structured?

Ionic solids exist in a crystal lattice structure where ions are bonded together with ionic bonds, forming a highly ordered arrangement.

What is a unit cell in the context of ionic solids?

A unit cell is the smallest repeatable unit of a crystal lattice, which when repeated, forms the entire crystal lattice structure.

Why do ionic solids have high melting points?

Ionic solids have high melting points because the ions are held together by strong ionic bonds that require a lot of energy to break.

How do covalent network solids differ from ionic solids in terms of bonding?

Covalent network solids are bonded by covalent bonds, which are shared pairs of electrons, forming a giant 3-dimensional array, as opposed to ionic bonds in ionic solids.

What are some examples of covalent network solids?

Examples of covalent network solids include boron nitride, diamond, graphite, quartz, and silicon carbide.

Why are metallic solids good conductors of electricity?

Metallic solids are good conductors of electricity because the valence electrons are delocalized and can flow around the metal nuclei, facilitating the movement of electrons.

What are the intermolecular forces that hold molecular solids together?

Molecular solids are held together by weak intermolecular forces such as hydrogen bonds, dipole-dipole interactions, and London dispersion forces.

Why do molecular solids have low melting points compared to other crystalline solids?

Molecular solids have low melting points because they are held together by weak intermolecular forces rather than strong covalent, ionic, or metallic bonds.

What is the difference between allotropes of an element?

Allotropes are different structural forms of the same element, bonded in different ways, resulting in different physical and chemical properties. Examples include diamond and graphite, which are both forms of carbon.

Outlines

00:00

🔬 Crystalline Solids Overview

This segment introduces the concept of crystalline solids, contrasting them with amorphous solids. Crystalline solids are characterized by their orderly and repeating atomic arrangements, which are divided into four main categories: ionic, covalent network, metallic, and molecular solids. The video promises an exploration of each category, starting with ionic solids. Ionic solids are defined by their crystal lattice structure where ions, exemplified by sodium chloride, are bonded together with strong ionic bonds. The smallest repeatable unit in this structure is known as the unit cell. Ionic solids have high melting points due to these strong bonds, are generally soluble in water, brittle, and do not conduct electricity in their solid form but do when dissolved in water.

05:02

💠 Covalent Network Solids

Covalent network solids are composed of nonmetal or metalloid atoms linked by covalent bonds, forming a vast 3-dimensional network. Examples include boron nitride, diamond, graphite, quartz, and silicon carbide. These solids have strong covalent bonds, leading to high melting points and hardness. Diamond and graphite, both pure carbon forms, differ in bonding and properties, demonstrating the concept of allotropes—compounds of the same element with different bonding and properties. Covalent network solids are generally hard, have high melting points, and are not soluble in most solvents. Some can conduct electricity, depending on their structure.

10:05

🔩 Metallic Solids

Metallic solids are made up of metal atoms bonded together by metallic bonds, where valence electrons are delocalized and shared among the metal ions. This electron cloud creates a strong bond, resulting in high melting and boiling points. Metals are good conductors of electricity due to the free movement of these electrons. They are also malleable and ductile, allowing them to be shaped. The video notes that while most metals are hard, there are exceptions like mercury, which is liquid at room temperature. Metallic bonds are responsible for the general properties of metals, such as high melting points and electrical conductivity.

15:08

🌡️ Molecular Solids

Molecular solids differ from the others as they are held together by weak intermolecular forces, such as hydrogen bonds, dipole-dipole interactions, and London dispersion forces, rather than strong covalent, ionic, or metallic bonds. Examples include ice, dry ice (solid CO2), sugar, and frozen ammonia. These solids have low melting and boiling points because the intermolecular forces are easily overcome. They do not conduct electricity, even in the liquid or molten state, due to the lack of free-moving ions or electrons. The video concludes by summarizing the properties of ionic, molecular, covalent network, and metallic solids, emphasizing the importance of recognizing the type of compound and its properties based on the chemical formula.

Keywords

💡Crystalline Solids

Crystalline solids are a type of solid where the atoms, ions, or molecules are arranged in a highly ordered, repeating pattern extending in all three spatial dimensions. This orderly arrangement is what gives crystalline solids their characteristic properties, such as fixed melting points and anisotropic physical properties. In the video, crystalline solids are contrasted with amorphous solids, where the atomic arrangement is random. The video discusses four main categories of crystalline solids: ionic, covalent network, metallic, and molecular solids.

💡Ionic Solids

Ionic solids are composed of ions held together by ionic bonds, which are electrostatic attractions between oppositely charged ions. These solids typically form a crystal lattice structure. An example given in the video is sodium chloride (table salt), where sodium (Na+) and chloride (Cl-) ions form a regular lattice. Ionic solids are characterized by high melting points due to the strong ionic bonds, and they often exhibit brittleness and are soluble in water due to their ability to form ion-dipole interactions with water molecules.

💡Unit Cell

The unit cell in crystallography is the smallest repeating unit that can be used to construct the entire crystal lattice by translation. It represents the fundamental building block of the crystal structure. In the context of the video, the unit cell is used to describe how ionic solids like sodium chloride are organized, with each unit cell containing a specific arrangement of ions that repeats throughout the crystal.

💡Covalent Network Solids

Covalent network solids are composed of atoms linked together by covalent bonds, forming a continuous network throughout the material. These solids are characterized by very strong covalent bonds that result in high melting points and hardness. Examples from the video include diamond, where each carbon atom is covalently bonded to four others in a tetrahedral structure, and quartz, which is a form of silicon dioxide with a continuous silicon-oxygen network.

💡Allotropes

Allotropes are different structural forms of the same element, in which the atoms are bonded together in different ways. The video mentions carbon allotropes such as diamond and graphite, which are both pure forms of carbon but have different bonding arrangements (sp3 for diamond and sp2 for graphite), leading to vastly different physical properties. The concept is also extended to other elements like oxygen, which can form O2 (molecular oxygen) or O3 (ozone), each with distinct properties.

💡Metallic Solids

Metallic solids are composed of metal atoms bonded together by metallic bonds, which involve a 'sea' of delocalized electrons that are shared among the positively charged metal ions. This electron sharing gives rise to the characteristic properties of metals, such as high electrical and thermal conductivity, malleability, and ductility. The video notes that metallic bonds are strong, leading to high melting and boiling points, and that metals are generally good conductors of electricity.

💡Molecular Solids

Molecular solids consist of molecules held together by intermolecular forces, which are weaker than the intramolecular covalent bonds that hold the atoms within the molecules together. These forces include hydrogen bonds, dipole-dipole interactions, and London dispersion forces. The video explains that molecular solids, such as ice, dry ice (solid CO2), and sugar, have lower melting and boiling points compared to other crystalline solids because the intermolecular forces are weaker. They do not conduct electricity because the electrons are not free to move between molecules.

💡Melting Point

The melting point of a solid is the temperature at which it changes from a solid to a liquid. In the context of the video, melting points are discussed in relation to the strength of the bonds within the solid. Ionic, covalent network, and metallic solids tend to have high melting points due to the strong bonds between their constituent particles. In contrast, molecular solids have lower melting points because the intermolecular forces are weaker.

💡Conductivity

Conductivity in the context of the video refers to the ability of a material to conduct electricity. Ionic solids do not conduct electricity in their solid form because the ions are tightly packed and cannot move. However, when dissolved in water, they can conduct electricity as the ions are free to move. Metallic solids are good conductors due to the delocalized electrons that can move freely. Molecular solids, on the other hand, do not conduct electricity because the electrons are not free to move between molecules.

💡Brittleness

Brittleness is a mechanical property of a material that indicates its tendency to break or shatter under stress rather than deform. The video mentions that ionic solids, due to the strong ionic bonds holding the ions together in a rigid lattice, are typically brittle. This property is demonstrated by the ability to crush table salt, which breaks along smooth planes, indicating a lack of plastic deformation before failure.

Highlights

Solids are categorized into amorphous and crystalline solids, with the latter having orderly and repeated atomic arrangements.

Crystalline solids are divided into four main types: ionic, covalent network, metallic, and molecular solids.

Ionic solids form a crystal lattice structure with ions bonded by strong ionic bonds, exemplified by sodium chloride.

The unit cell is the smallest repeatable unit in a crystal lattice, crucial for forming the lattice structure.

Ionic solids have high melting points due to the strength of ionic bonds.

Ionic compounds are generally soluble in water due to ion-dipole interactions.

Mechanically, ionic solids are brittle and exhibit smooth cleavage when broken.

Ionic solids do not conduct electricity in solid form but do when dissolved in water due to ion mobility.

Covalent network solids are characterized by a giant 3-dimensional array of covalently bonded atoms.

Examples of covalent network solids include diamond, graphite, quartz, and silicon carbide.

Diamond's tetrahedral structure with sp3 hybridization gives it exceptional hardness.

Graphite's layered structure with sp2 hybridization allows it to form sheets and slide over each other.

Allotropes are different structural forms of the same element, such as diamond and graphite.

Covalent network solids have high melting points and are generally hard and not soluble in most solvents.

Metallic solids are bonded by metallic bonds, where valence electrons are delocalized among metal atoms.

Metallic solids have high melting and boiling points and are good conductors of electricity.

Metals are malleable and ductile due to the nature of metallic bonding.

Molecular solids are held together by weak intermolecular forces, such as hydrogen bonds and London dispersion forces.

Examples of molecular solids include ice, dry ice, sugar, and frozen ammonia.

Molecular solids have low melting points and do not conduct electricity due to the weakness of intermolecular forces.

A summary of crystalline solids includes recognizing the type of bonding and properties like melting points and conductivity.

Transcripts

in this video we are going to examine

the different types of crystalline

solids as we know solids can be placed

into very broad categories amorphous

solids where the atoms are arranged very

randomly and crystalline solids where

the atoms are arranged in a very neat

orderly a repeated fashion crystalline

solids can be broken into four large

categories ionic solids covalent Network

solids metallic solids and molecular

solids so in this video we're going to

look at each of these categories of

crystalline solids and talk about some

of their properties let's start off with

ionic solids ionic compounds exist in

what is called a crystal lattice

structure where the ions in this case

let's look at a maybe a sodium chloride

crystal lattice you have a sodium and

chloride ions all bonded together in

this orderly crystal lattice structure

of course these ions are bonded together

with ionic bonds which are very strong

so each ion is connected with an ionic

bond which you can see represented with

the kind of stick the smallest

repeatable unit of this crystal lattice

is what is called the unit cell so here

you can see this is what would be a unit

cell for sodium chloride and if you

place a bunch of these together you

would form the crystal lattice so here

you can see it in a slightly different

way here we are showing you the chloride

and the sodium atoms each of those

bonded together and the smallest unit of

that would be the unit cell now the type

of unit cell there are many different

types of unit cells and that varies

depending on the ions involved and their

charges so we'll look at that too a

little later some of the properties of

ionic solids because all of the ions are

held together by very strong ionic bonds

it takes a lot of energy to separate

those so you have to have the

temperature very high so ionic solids

tend to have very high melting points

because of the strong ionic bonds

holding them together

solubility ionic compounds tend to be

soluble in water because they can form

ion dipole interactions which are fairly

favorable with water molecules and of

course not all ionic compounds are

soluble in water this is a very general

statement but in general ionic compounds

tend to be fairly soluble in water it's

a mechanical properties of ionic sauce

they tend to be brittle and break apart

of course if you've taken some sodium

chloride or table salt you can crush it

it's very brittle and if you strike it

with a hammer shown here you'll get a

generally a smooth cleavage point

between the ions so you can see that

represented here in conductivity ionic

solids tend not to conduct electricity

because the ions are very tightly packed

and they cannot move so they don't

conduct the electrons they do however

contend to conduct electricity when they

are dissolved in water because they're

the ions are able to move and disperse

throughout the water and carry the

current through the water let's look at

our second class of crystalline solids

covalent Network solids in covalent

Network solids the atoms which tend to

be

nonmetal or metalloids

they're all linked together by covalent

bonds into a giant 3-dimensional array

of course ionic compounds we saw were

held together by ionic bonds covalent

Network all of the atoms are bonded

together with very strong covalent bonds

which are shared pairs of electrons you

can see some examples here on the left

side of the screen boron nitride a

diamond graphite those are two very good

examples quartz silicon carbide

and here you can see for example diamond

all of the carbon atoms are bonded

together in these giant three

dimensional arrays and everything is

linked together with covalent bonds

graphite is similar they tend to be

arranged slightly different in these

sheets you can see silica which is

silicon oxide represented here if we

look at quartz which is a form of

silicon dioxide you can see that you

have silicon atoms covalently bonded to

oxygen atoms and then those oxygen atoms

are covalently bonded to neighboring

silicon atoms so all the atoms are

bonded in these with these strong

covalent bonds you can see the unit cell

here in the dashed lines and so again

all of these unit cells combine together

to make various these very large crystal

lattice structures so again I just want

to stress that all the atoms are bonded

together with covalent bonds which are

very strong let's look at diamond and

graphite we mentioned those earlier

diamond and graphite are both composed

of pure carbon atoms however the bonding

is slightly different and diamond each

of these little spheres here would be a

carbon atom and the stick is a covalent

bond to another one you can see that

each carbon atom let's look at this one

right here this carbon atom is bonded to

one two three four other carbon atoms so

we would say that carbon is sp3

hybridized and has tetrahedral geometry

so because it's bonded in this

tetrahedral shape it has a very strong

structure and that's what gives diamonds

its properties a graphite again is pure

carbon but the bonding of the graphite

the carbons in graphite are slightly

different here you can see let's look at

this carbon here this carbon is bonded

to one two three other carbons so this

carbon would be sp2 hybridized and so

it's bonding of these carbons is

different than the bonding and diamonds

so graphite and diamonds we can have

very different properties graphite tends

to form these sheets which are bonded

together but again the bonding of the

carbon atoms is very different and so we

call those allotropes compounds that

have are composed of the same types of

atoms but their bonding is different so

they have very different chemical

properties there are actually a lot of

different allotropes of carbon you can

see here a this represents diamond and

here and B that represents graphite but

there are lots of others there's

actually eight allotropes of carbon only

a diamond and graphite form are

considered covalent Network solids but

there are several others we have C 60 C

540 we have nanotubes perhaps you've

heard of those we also have Amorphis

carbon so all of these are pure carbon

but the bonding is different which gives

them different shapes and structures and

therefore physical and chemical

properties now there are also a low

troupes of other compounds other than

just carbon silicon dioxide if it's

bonded in a certain way it forms quartz

and we looked at that if it's bonded

other ways you'd form sand in quartz

glass

oxygen molecular oxygen o2 that's an

ozone o3 those are allotropes again both

composed of pure oxygen but they're

bonded together differently

forming different compounds so there are

lots of different allotropes and here

again just a representative sample so

back to the covalent Network solids some

properties of those the covalent Network

solids as we mentioned they're all the

atoms are bonded together with strong

covalent bonds so that would give them a

very high melting point takes a lot of

energy to break

atoms apart or disrupt them enough to go

into the liquid phase they tend to be

very hard again because of the bonding

some but not all are able to conduct

electricity and in general they're not

soluble in polar or nonpolar solvent so

again that's a very general statement

but they tend not to be very soluble so

let's look at our third class of

crystalline solids metallic solids

metallic solids of course are composed

of a group of metal atoms which are

bonded together and the metal atoms are

bonded together with a unique type of

bond called a metallic bond in a

metallic bond which we're representing

in the figure here you basically have

all of the metal nuclei so let's say

that this represents iron here's an iron

nucleus in iron nucleus and iron nucleus

and then the valence electrons of all

the iron atoms are kind of free to flow

around all the nucleus or the valence

electrons it's almost as if you have all

of these nuclei kind of floating in a

sea of electrons of valence electrons so

we say those electrons are I tend to be

delocalized and because all these

electrons are shared that causes them to

be bond all of these atoms to be bonded

together some properties of metallic

solids the metallic bond again like an

ionic and covalent is a very strong bond

and so it takes a lot of energy to

disrupt or weaken those so the metallic

compounds tend to have very high melting

and boiling points you can see some of

these transition metals the boiling

points are seeing the melting points of

those are very very high so around 1,500

to 2,000 degrees Celsius some are very

low there's always exceptions like

mercury is actually a little liquid at

room temperature but in general metallic

solids tend a very high melting and

board

points because of the nature of the

metallic bond the electrons valence

electrons are delocalized so they're

able to kind of flow around the nuclei

that makes the metals very good

conductors of electricity they also tend

to be fairly malleable and ductile 'ti

so they can be shaped generally they can

be heated and bent and again due to the

nature of the metallic bond they tend to

be generally fairly hard again there's

always exceptions to them our final

class of crystalline solids are

molecular solids now these are a little

different than the other three molecular

solids the molecules are held together

not by strong ionic or covalent or

metallic bonds but the molecules are

held together by fairly weak

intermolecular forces things like

hydrogen bonds dipole-dipole

interactions London dispersion forces

and relative to true bonds which are

covalent metallic ionic these

intermolecular forces tend to be fairly

weak some examples of molecular solids

are water in the solid phase of course

ice solid carbon dioxide dry ice sugar

frozen ammonia again all of these are

molecules as you can see and so if we

cool them enough then we can get them

into the solid phase and their atoms do

a line into these crystalline patterns

though so a few properties of these are

Maalik molecular solids of course are

composed of molecules which are two or

more nonmetals which are covalently

bonded together so here I think we're

looking at a image of water here here

it's important to appreciate that

we have the oxygen atom here in red

the oxygen atom is covalently bonded to

the two hydrogen's so we have the intra

molecular forces which hold the the

molecule together hold the different

atoms together I should say and those

are very strong covalent bonds so the

atoms in the molecule are held together

by strong covalent bonds

we call those intra molecular forces the

molecules themselves in the solid phase

are held to one another by much weaker

intermolecular forces and of course

these are formed when we cool these

molecules low enough for these

intermolecular forces to kind of grab

hold and solidify the substance some

properties of molecular solids because

the molecules are had to get held

together by weak interactions they tend

to have low melting and boiling points

relative to the other compounds or the

other crystalline solids we discussed

when you melt these the molecules

themselves are not broken up they just

move apart from each other so when you

melt a molecule again you're not

breaking any covalent bonds you're just

disrupting the intermolecular forces

molecular solids attend not to conduct

electricity even when they're in the

liquid or molten phase because there are

generally no free ions and the ions are

would conduct the electrons and

therefore the electricity so let's wrap

up with just a very brief summary of our

different types of crystalline solids

ionic compounds the atoms or ions are

held together by very strong ionic bonds

they tend to be brittle hard high

melting points you can see some examples

here with an ionic compound you want to

look for a metal and then an one or more

nonmetal so typically an ionic compound

we have a metal cation and one or more

nonmetal anions molecular solids we just

discussed those the atoms are held

together by covalent bonds but

molecules are held together by weak

intermolecular forces

Londyn dispersed and dipole-dipole or

hydrogen bonds they tend to be soft low

melting and not conduct electricity you

can see here some examples you notice

that all of these are composed of

nonmetals and so that's a hallmark of

molecular compounds a covalent Network

compounds all the atoms are held

together by covalent bonds they tend to

be very hard high melting points we

showed you some examples graphite

diamond silicon dioxide such as quartz

and metallic compounds metallic solids

the atoms the metal atoms are held

together by metallic bonds and we

discussed those they tend to be hard but

there's a variation a lot variation in

that they have high melting points and

they're good conductors of electricity

in metals of course of things like

sodium zinc copper iron platinum gold

silver all of those so be familiar with

these and and you want to be sure that

you can recognize the type of compound

and therefore the type of solid if you

are given the molecular formula or the

chemical formula so be sure that you're

familiar with these types of crystalline

solids and their general properties