6.4 Metallic Bonding

Summary

The video explains metallic bonding, emphasizing how metals differ from ionic and covalent bonds. Unlike those, metals have free-moving electrons, forming a 'sea of electrons' that allows for efficient conductivity of electricity and heat. Metals can absorb a wide range of light frequencies due to many free electrons, making them shiny. Their malleability and ductility are attributed to their lattice structure, allowing atoms to slide past each other easily. The bond strength in metals varies and is measured by enthalpy of vaporization, which reflects the energy needed for a metal to change from liquid to gas.

Takeaways

• 🔗 Metallic bonding is neither ionic nor covalent, and metals have unique properties different from these types of bonds.
• ⚡ Metals are excellent conductors of electricity because their electrons can move freely across large sheets of metal.
• 🔬 In metallic bonds, electrons are delocalized, forming a 'sea of electrons' that can flow freely, contributing to their conductive properties.
• 📊 Alkali metals, alkali earth metals, and D-block elements have free orbitals that allow their electrons to disassociate and move freely.
• 💡 The lattice of metals, like ionic lattices, remains intact while the free electrons conduct electricity and heat.
• 🔥 Metals can transfer heat efficiently because free electrons quickly transfer momentum across the material.
• ✨ Metals are shiny and reflective because their free electrons can absorb and release a wide range of light frequencies.
• 🔧 Metals are malleable and ductile because their atomic lattice allows layers of atoms to slide past each other without breaking the structure.
• 📏 Ionic compounds lack malleability and ductility because of strong ionic repulsion, which can cause them to shear when stressed.
• 🌡️ The strength of metallic bonds varies between metals and can be measured using enthalpy of vaporization, indicating how much energy is needed for phase change from liquid to gas.

Q & A

What is metallic bonding?

Metallic bonding is a type of chemical bonding that is neither ionic nor covalent. It involves a 'sea' of delocalized electrons that are free to move across a metal lattice, allowing for properties such as electrical conductivity and malleability.

Why are metals good conductors of electricity?

Metals are good conductors of electricity because the electrons in a metallic bond are delocalized and can move freely across the metal lattice, facilitating the flow of electric current.

How do the properties of metallic bonding differ from ionic and covalent bonding?

In ionic bonding, electrons are bound to one atom, typically an anion, while in covalent bonding, electron pairs are shared between atoms but still localized. In contrast, metallic bonding allows electrons to be delocalized and move freely across the entire metal.

Which elements on the periodic table are known for their metallic bonding?

Alkali metals and alkaline earth metals, which fill the s orbital and have the p orbital completely free, are known for their metallic bonding. Additionally, most d-block elements have both their p and d orbitals free, contributing to metallic bonding.

What is the significance of the sea of electrons in metallic bonding?

The sea of electrons in metallic bonding refers to the delocalized electrons that surround the positively charged metal ions. This sea of electrons allows for the metal's high electrical conductivity, thermal conductivity, and malleability.

Why are metals shiny and reflective?

Metals are shiny and reflective because the free electrons can absorb a wide range of light frequencies, then quickly return to their ground state, releasing the absorbed energy as photons of light.

What is the role of malleability and ductility in the shaping of metals?

Malleability allows metals to be bent into thin sheets, while ductility allows them to be drawn into thin wires. These properties arise from the ability of metal atoms to slide past each other due to the lack of repulsion between them, unlike in ionic compounds.

How does the strength of metallic bonds vary between different metals?

The strength of metallic bonds varies due to factors such as atomic size and nuclear charge. These factors affect the bond strength, which can be measured by the enthalpy of vaporization.

What is the enthalpy of vaporization, and how does it relate to metallic bond strength?

The enthalpy of vaporization is the amount of energy required for a metal to transition from a liquid to a gaseous phase. It is related to metallic bond strength because this transition involves separating the metal atoms from the sea of electrons.

How does the structure of a metal lattice differ from that of an ionic compound?

A metal lattice is less robust than an ionic compound due to the absence of repulsion between metal atoms. This allows metal atoms to slide past each other easily, which is not possible in ionic compounds where such repulsion would cause layers to shear and break bonds.

What happens to the electrons in a metal when it is heated?

When a metal is heated, the electrons can quickly transfer the thermal energy (momentum) across the metal lattice, distributing the heat evenly and affecting the internal structure of the metal.

Outlines

00:00

⚡ Understanding Metallic Bonding and Its Properties

In this section, the concept of metallic bonding is introduced as distinct from ionic and covalent bonding. Metallic bonds allow for free-flowing electrons across a sheet of metal, unlike in ionic and covalent compounds where electrons are more restricted. This freedom of electron movement is key to explaining why metals are excellent conductors of electricity and heat. The alkali and alkaline earth metals, as well as d-block elements, have free p and d orbitals that allow electrons to become delocalized, forming a 'sea of electrons' around positively charged metal ions. This interaction between the lattice of metal ions and the electron sea explains the structural integrity of metals and their ability to conduct electricity and heat efficiently.

05:01

🌟 Metallic Properties: Ductility, Malleability, and Shiny Appearance

This section highlights the physical properties of metals, explaining why they are ductile, malleable, and shiny. Metals can be shaped into wires or sheets due to the flexibility in their crystal lattice, which lacks the rigid structure found in ionic compounds. Unlike ionic compounds, metal atoms can slide past each other without repelling, allowing for reshaping without breaking. The shiny appearance of metals is attributed to their ability to absorb and re-emit a wide range of light frequencies due to the numerous free electrons and available orbitals, which makes them reflective and lustrous. Additionally, the concept of metallic bond strength is briefly introduced, varying between metals based on atomic size and nuclear charge, which can be measured by the enthalpy of vaporization.

Keywords

💡Metallic Bonding

Metallic bonding is a type of chemical bonding distinct from ionic and covalent bonds. It occurs in metals, where electrons are free to move across a 'sea of electrons' shared by all atoms. This mobility of electrons gives metals their unique properties such as conductivity, malleability, and ductility. In the video, metallic bonding is explained as the main focus, describing how metals differ from other types of bonded compounds.

💡Sea of Electrons

The 'sea of electrons' refers to the delocalized electrons in a metal that are not bound to any specific atom. These free electrons can move throughout the metal, allowing it to conduct electricity and heat efficiently. In the video, this concept is used to explain how metallic bonding works, with electrons flowing freely over large areas, providing flexibility and conductivity.

💡Conductivity

Conductivity is the ability of a material to transfer electricity or heat. In metals, conductivity is particularly high due to the free movement of electrons in the 'sea of electrons.' The video contrasts this with ionic and covalent compounds, where electrons are more tightly bound, explaining why metals are better conductors.

💡Lattice Structure

A lattice structure in metals refers to the organized arrangement of atoms held together by metallic bonds. This structure allows the electrons to move freely around the nuclei, which are fixed in place, making the metal solid yet flexible. The video compares this structure to ionic compounds' lattices, which are more rigid and less flexible.

💡Malleability

Malleability is a property of metals that allows them to be hammered or pressed into thin sheets without breaking. This is possible because the lattice structure in metals allows atoms to slide past each other without repulsive forces breaking the material. The video highlights this property as a key difference between metals and ionic compounds.

💡Ductility

Ductility is the ability of a material to be stretched into thin wires. Metals are ductile because their atoms can slide past each other within the lattice structure without breaking the bonds. The video mentions this property to explain why metals can be shaped into wires, unlike ionic compounds that would shatter under similar stress.

💡Delocalized Electrons

Delocalized electrons in metals are not confined to a single atom but can move freely across the entire structure. These electrons are crucial to the properties of metallic bonding, allowing metals to conduct electricity and heat. The video emphasizes the importance of delocalized electrons in explaining why metals behave differently from ionically or covalently bonded materials.

💡Enthalpy of Vaporization

The enthalpy of vaporization is the amount of energy required to convert a liquid into a gas. In the context of metals, it is used to measure the strength of the metallic bond by determining how much energy is needed to break the bond between the atoms and the 'sea of electrons.' The video uses this concept to illustrate variations in metallic bond strength across different metals.

💡Alkali Metals

Alkali metals are elements found in Group 1 of the periodic table, characterized by having one electron in their outermost s-orbital. These metals, such as sodium and potassium, are highly reactive and form metallic bonds. The video refers to alkali metals when discussing the periodic table and how their electronic structure affects metallic bonding.

💡Reflectivity and Luster

Reflectivity and luster refer to a metal's ability to reflect light, making it shiny. This property arises from the movement of free electrons, which can absorb a wide range of light frequencies and then emit photons when they return to a lower energy state. The video explains how this is a direct consequence of the delocalized electrons in metallic bonding.

Highlights

Metallic bonding is neither ionic nor covalent and differs significantly in properties from ionically bonded and covalently bonded compounds.

Metals are excellent conductors of electricity, even better than molten ionic compounds, due to the freedom of electron movement.

In metallic bonding, electrons are delocalized and can move freely across a sheet of metal, unlike ionic or covalent bonds where electrons are bound.

Alkali metals and alkali earth metals have their p orbital free, allowing the delocalization of electrons across metal sheets.

Electrons in metals can disassociate from their host atoms and form a 'sea of electrons,' allowing them to flow freely and conduct electricity and heat.

The 'sea of electrons' enables metals to be highly conductive as electrons can transfer momentum quickly, leading to efficient heat and electrical transfer.

Metals absorb a wide range of light frequencies due to the availability of many free orbitals that electrons can move up to.

The reflection and shiny appearance of metals are caused by excited electrons returning to their ground state, releasing photons of light.

Metals are ductile and malleable, meaning they can be shaped into thin sheets and wires due to the flexibility of the metallic bond.

Unlike ionic compounds, metals do not experience strong repulsion between ions, allowing atoms to slide past each other and form various shapes.

The strength of metallic bonds varies between metals due to differences in atomic size and nuclear charge.

Enthalpy of vaporization measures the metallic bond strength, as it indicates the energy absorbed when a metal transitions from liquid to gas.

In metallic bonding, metals form a structured lattice where the sea of electrons flows around positively charged atomic cores.

The freedom of electron movement explains why metals can conduct current and heat efficiently, making them essential in electrical applications.

Electrons in metals are loosely associated with their atoms, enabling them to become delocalized and contribute to the overall conductivity and malleability of metals.

Transcripts

so in this video we're going to be

discussing chapter 6 section 4 which

covers metallic bonding and the first

thing you need to know is that metallic

bonding is neither uh

ionic nor Cove valent I know earlier it

may have seemed that there was sort of

an absolute to well there's not the

thing is uh metals are very different

just in their properties from ionically

bonded and coal bonded compounds uh for

example they're very excellent

conductors of electricity I by the way

is the symbol for electric current in

physics and

uh they're even better at conducting

than molten ionic compounds and this is

because uh in ionic and calent compounds

what happens is that at least in the

ionic compounds the electrons are sort

of bound to one atom either the cation

or the Antion usually the Antion and and

then in calent bonds electron pairs are

shared however uh they're still bound

to however many uh atoms are within that

molecule in which they're shared they

have no freedom to sort of roam across

an entire material however in metals

electrons can flow freely across a whole

sheet of metal several

meters uh across so to explain this

Behavior we of course have to look to

the periodic table and the first thing

you'll notice is that the alkali metals

and The Alkali earth metals over here

which fill the uh s orbital have the p

orbital completely free and all of the D

Block

elements have both their P orbitals free

and the majority of

them will tend to have much of the D

Block free as well so what you end up

having is that um many electrons in

these

Metals uh in sheets of these metals are

sort of loosely associated with their

atoms so much so that they uh are able

to

disassociate from their host atom and

become uh delocalized and now what this

means is that the atoms or the uh

electrons rather can flow freely across

a sheet of metal completely leaving

their host atoms in a sort of sea of

electrons meanwhile the uh inner shells

of the metals and their nuclei are

attached in a lattice much like the salt

lattice I discussed in my lat last

video but they are uh bonded by positive

and negative charge to the Sea of

electrons and this sort of uh

relationship between a very structured

ladder forms a sheet or Rod or car spoke

or what have you of metal and the Sea of

electrons flowing around these uh Center

Parts uh is what's called metallic

bonding so the freedom of electrons in

this sort of electron C to move freely

through the

metal uh mostly unimpeded is what allows

them to conduct

current as well as heat uh so well

because the electrons can quickly

transfer

momentum from one end of the metal if we

draw a sheet of metal down here very

poorly if you heat up one

end what you'll find is that the

electrons from over here can quickly

transfer that momentum across the sheet

of metal evenly to distribute the this

momentum uh caused by the heat across

this whole C which in turn will affect

the uh internal lattice of atoms also

because there are so many free electrons

in metals what you'll find is that

there's a wide range of

frequencies that the metals can absorb

because there are so many free

orbitals that electrons can occupy and

then be moved up to moved up to and as

we know you can only absorb uh light

within certain wavelengths so that you

can move up a specific uh amount of

energy and so what ends up happening in

metals is that because there's so many

electrons with so many options they can

absorb metals can absorb a wide range of

frequencies of light and then these

electrons which are then excited will

quickly go back down to their ground

state dissipating this energy as a

photon of light that comes away from

metal and this is why metals tend to be

uh shiny very reflective and lustrous

now metals are also very ductile they're

able to be uh formed into different

shapes and this is because of two

properties of metals the first is Mal

ability which is a material especially

Metals ability to be bent into thin flat

sheets and the second property

is ductility

which is the

ability of a material to be sort of

extruded and forced into thin wires and

the reason metal can do this whereas

ionic compounds cannot is because the

uh uh crystal

lattice of metal isn't made of such a uh

robust structure as it is in ionic comp

compounds so in metals because there's

no uh repulsion between certain

ions as there are in ionic compounds

these atoms can easily slide past each

other in order to form whatever shape

you like whereas in an ionic compound if

you try to slide say chlorine past

another chlorine or they're both ionic

this repulsion will cause the layers to

Shear breaking the comp finally the

bonds between

uh different metals and their sea of

electrons that is the metallic bond

strength uh varies from metal to metal

and this is because of size of the atom

various nuclear charge

Etc and both the uh effect of the

changing nuclear charge and its bond to

the electron C can be measured by

property called the

enthalpy

of vaporization now I know that sounds

really complicated but it really isn't

what it is

is the amount of energy in

kles that a metal observe absorbs rather

uh when it uh goes from a from a liquid

phase to a gaseous

phase and the reason you can measure uh

enthalpy of vaporization in order to get

its metallic bond strength is because

when it goes from a liquid to a gas it

separates from a she sea of electrons

and becomes its own independent

atom