6.4 Metallic Bonding
Summary
TLDRThe 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
⚡ 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.
🌟 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.
Mindmap
Keywords
💡Metallic Bonding
💡Sea of Electrons
💡Conductivity
💡Lattice Structure
💡Malleability
💡Ductility
💡Delocalized Electrons
💡Enthalpy of Vaporization
💡Alkali Metals
💡Reflectivity and Luster
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
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