Network Solids and Carbon: Crash Course Chemistry #34

CrashCourse
14 Oct 201308:19
EducationalLearning
32 Likes 10 Comments

TLDRThis video explains the atomic structure of the network solids graphite and diamond, both composed of carbon. Despite their identical composition, the arrangement of carbon atoms in layers or a 3D lattice results in vastly different material properties like hardness, electrical conductivity, and heat conduction. For instance, strong covalent pi bonds between carbon atoms in 2D sheets give graphite a layered structure that makes it soft and lubricating but electrically conductive, whereas 3D bonding makes diamond extremely hard and an electrical insulator. The video elucidates how subtle differences in chemical bonding patterns can profoundly impact bulk material qualities.

Takeaways
  • 😀 Diamond and graphite are both made of carbon, but have very different properties due to their atomic structure
  • 🔬 Graphite is composed of sheets of bonded carbon atoms that can slide past each other, making it soft and lubricating
  • 💎 Diamond has a 3D network of carbon bonds, making it extremely hard and durable
  • ⚡ Graphite conducts electricity while diamond does not due to differences in their bonding
  • 🔥 Both graphite and diamond conduct heat well due to their strong covalent bonds
  • 📝 The weak bonds between graphite sheets allow it to rub off onto paper, making it useful in pencils
  • 🛠 Diamond's hardness makes it useful for cutting, drilling, and grinding other materials
  • 💸 Graphite is more abundant and cheaper than diamond, allowing mass production of items like pencils
  • 🚨 Applying very high heat and pressure can convert graphite into diamond, but the reverse is virtually impossible
  • 🤯 The arrangement of the same carbon atoms into different bonding patterns completely changes the properties of the resulting material
Q & A

  • What are the two different ways that carbon atoms can bond to form network solids?

    -Carbon atoms can bond in two different ways to form the network solids graphite and diamond. In graphite, the carbon atoms form sheets with hexagonal structures bonded by strong pi bonds. In diamond, the carbon atoms form a 3D tetrahedral network with sp3 hybridized orbitals.

  • How does the bonding in graphite and diamond affect their physical properties?

    -The sheet structure of graphite allows the layers to slide past each other, making graphite soft and a good lubricant. The 3D network of diamond makes it extremely hard. Graphite conducts electricity while diamond is an insulator.

  • Why is graphite a good conductor of electricity but diamond is not?

    -Graphite contains pi bonds between the carbon atom sheets. These allow electrons to move freely and conduct electricity. Diamond contains strong sigma bonds that do not allow electron movement, making diamond an electrical insulator.

  • Why can both graphite and diamond conduct heat well?

    -They both contain strong covalent bonds. When an atom vibrates from heat energy, these bonds cause surrounding atoms to vibrate as well, spreading the heat through the entire structure.

  • What conditions are required to turn graphite into diamond?

    -Graphite can be turned into diamond through intense heat (around 3000°C) combined with extreme pressure (around 15 million kPa), which provides enough energy to break and reform the atomic bonds into a 3D diamond structure.

  • Why is it so difficult for diamond to turn into graphite naturally?

    -The covalent bonds in diamond have very high activation energy, so their breakdown rate at normal temperatures is essentially zero. The conditions required to break down diamond into graphite do not occur naturally on Earth.

  • How does the directionality of bonding affect the properties of graphite and diamond?

    -Graphite's two-dimensional sheet structure makes it strong in those directions but weak between layers. Diamond's three-dimensional bonding makes it hard and resistant to stress from any direction.

  • Why can graphite be used as pencil lead but diamond cannot?

    -Graphite's weak inter-layer bonding allows sheets to slide off as pencil lead deposits it onto paper. Diamond is too hard and rigid to rub off onto surfaces.

  • Why are diamonds used for cutting and drilling while graphite is used for lubricating?

    -Diamond's extreme hardness allows it to cut through other materials. Graphite's soft layered structure allows surfaces to slide smoothly over it, making it an excellent lubricant.

  • How does the arrangement of the same material into different atomic networks produce such contrasting properties?

    -Even though graphite and diamond are both made of pure carbon, the different two- and three-dimensional bond networks change the physical, chemical, electrical, and thermal characteristics dramatically.

Outlines
00:00
🎬 Description of Graphite and Diamond Structures

This paragraph introduces graphite and diamond as two forms of carbon with different properties due to their atomic structure. It explains that graphite has a layered sheet structure held together by weak van der Waals forces, allowing the layers to slide over each other. This makes graphite useful for writing and lubrication. Diamond has a rigid 3D network structure, making it incredibly hard and resistant to stress.

05:00
🔬 Differences in Properties between Graphite and Diamond

This paragraph highlights key differences in properties between graphite and diamond due to their differing atomic structures. Graphite conducts electricity due to mobile pi electrons but diamond does not, acting as an insulator. Both conduct heat well due to strong covalent bonds. Graphite can be converted to diamond under conditions of extreme heat and pressure, but the reverse transformation is virtually impossible on Earth.

Mindmap
Keywords
💡Network solids
A type of atomic solid where the atoms are linked together in a network structure. This makes them strong and stable. Examples from the script: 'Network solids are one of the three types of atomic solid...', 'Like all networks, this makes the arrangement both strong and stable.'
💡Graphite
A form of carbon with a sheet-like structure held together by strong pi bonds within the sheets. The sheets slide past each other easily. Used in pencils. Examples: 'You know this carbon network by its street name, graphite.', 'The sheets are only held to other sheets with really weak van der Waals forces. They allow the sheets to slide on top of each other so they can easily be removed layer by layer. That's why graphite is so fantastic for writing with.'
💡Diamond
A form of carbon with a 3D network structure formed by SP3 hybridized orbitals. This makes diamonds incredibly hard. Examples: 'In diamond, the bonding electrons in each atom are arranged in SP3 orbitals. Four lobes that are as far away from each other as possible.', 'This stability is what makes diamonds so famously, incredibly hard.'
💡Covalent bonds
A type of strong chemical bond where electron pairs are shared between atoms. Determines structure and properties of network solids. Examples: 'they're all bonded together with covalent bonds,, so what's the difference?', 'These type bonds require that when one atom vibrates, all the atoms around it vibrate, in exactly the same way.'
💡Atomic orbitals
Regions of space with a high probability of finding an electron. Their arrangement (hybridization) affects network solid properties. Examples: 'If you've seen our episode on atomic orbitals you know the way the electron in a chemical bond orient themselves,, that is the way the orbitals hybridize, has a big effect on the properties of a substance.'
💡SP2 hybrid orbital
A type of hybrid orbital with 2 lobes. Forms planar, 2D sheet structures like in graphite. Example: 'the outer electrons of each atom are arranged into SP2 hybridized orbitals, with two other atoms for a total of three lobes in a flat trigonal arrangement. This creates a sheet with a hexagonal structure'
💡SP3 hybrid orbital
A tetrahedral hybrid orbital with 4 lobes pointing to corners of a tetrahedron. Bonds carbon atoms in 3D diamond structure. Example: 'In diamond, the bonding electrons in each atom are arranged in SP3 orbitals. Four lobes that are as far away from each other as possible.'
💡Pi bonds
Bonds formed from unhybridized p orbitals perpendicular to the plane of hybrid orbitals. Give graphite strength within its sheets. Example: 'These orbitals form pi bonds that merge into an extensive network on their own, and this gives the structure its real strength.'
💡Sigma bonds
Main covalent bonds formed from overlap of hybrid orbitals. Do not allow free electron movement so diamond does not conduct electricity. Example: 'Diamonds are electrical insulators, not conductors. This is because the carbon atoms in diamonds share sigma bonds, not pi bonds,, which doesn't allow the electrons much freedom of movement.'
💡Cleavage plane
A plane between atoms along which a crystal or solid tends to break. Example: 'And these are the places that the diamond cutters attack first.'
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Transcripts

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