Send the link below via email or IMCopy
Present to your audienceStart remote presentation
- Invited audience members will follow you as you navigate and present
- People invited to a presentation do not need a Prezi account
- This link expires 10 minutes after you close the presentation
- A maximum of 30 users can follow your presentation
- Learn more about this feature in our knowledge base article
Do you really want to delete this prezi?
Neither you, nor the coeditors you shared it with will be able to recover it again.
Make your likes visible on Facebook?
Connect your Facebook account to Prezi and let your likes appear on your timeline.
You can change this under Settings & Account at any time.
Giant covalent structures
Transcript of Giant covalent structures
State covalent structures
Illustrate their structures and state their physical properties
Explain their properties and uses
Giant covalent structures contain a lot of non-metal atoms, each joined to adjacent atoms by covalent bonds.
Carbon can bond to millions of other carbon atoms in different ways creating different types of giant covalent structures. Two of these structures are diamond and graphite. They are called allotropes of carbon.
Allotropes have the same chemical properties; as they have the same number of electrons. However, different physical properties as the
electrons are shared in different ways with other atoms.
PROPERTIES AND USES OF GIANT COVALENT STRUCTURES
Diamond is giant lattice; where one carbon atom is bonded to four other carbon atoms forming a tetrahedral arrangement; its structure is called a giant covalent lattice.
Diamond is the hardest substance on earth due to its arrangement of strong covalent bonds; which makes it a very useful material i.e. diamond coated surgical instruments used for delicate eye operations.
High melting point and boiling point due to the millions of strong covalent bonds within the giant covalent lattice; which requires allot of energy to separate the atoms.
Insoluble in water, as the particles are not charged (unlike ionic compounds), so water molecules are not attracted to them.
Does not conduct electricity as there are no free ions or electrons to carry charge.
Lustrous (shiny), colourless and clear (transparent).
Graphite also has a giant covalent structure; however its properties are very different to other giant covalent structures; as one carbon atom is covalently bonded to three other carbon atoms; arranged in hexagonal layers.
The hexagonal layers are held together by weak forces called intermolecular forces; which allow the layers to slide over each other. Hence, graphite is much softer than diamond. So, if you touch a lump of graphite it feels smooth and slippery.
Silica, also known as silicon dioxide; commonly known as sand has a similar structure to that of diamond; giant covalent structure.
It contains a silicon atom covalently bonded to four oxygen atoms. Each oxygen atom is covalently bonded to two silicon atoms.
last but not least.....
They do not conduct electricity...
as they don't have many atoms they cannot
extend like giant structures.
The atoms are usually arranged into giant regular lattices, which are extremely strong as many bonds are involved.
Carbon can form a maximum of four covalent bonds; as it has an electronic arrangement of 2, 4. It can form 4 covalent bonds with itself.
Silicon dioxide is very hard but lesser than diamond as the covalent bonding arrangement is between carbon atoms only in diamond. Compared to silicon dioxide where silicon and oxygen atoms are involved.
Silicon dioxide has a very high melting point (1,610 °C) and boiling point (2,230 °C). This property results from millions of very strong covalent bonds that hold the silicon and oxygen atoms together in the giant covalent structure.
It is insoluble in water, and does not conduct electricity.
Silicon dioxide is found as in granite, and is the major compound in sandstone. The sand on a beach is made mostly of silicon dioxide.
development of new computers, new catalysts, new coatings, highly selective sensors and stronger and lighter construction materials.
Diamond consists of a carbon atom bonded to four other carbon atoms. So, all the outer electrons in carbon are involved in bonding (full shell). However, in graphite one carbon atom is bonded to three other carbon atoms, which leaves an un - bonded electron.
The one un-bonded electron is therefore found above and below the plane of the layers; they are held loosely to the carbon atoms so can drift along the layers of graphite. The electrons are called 'delocalised' electrons, which make graphite a good conductor of electricity.
forces between the
layers hold the
How can graphite conduct electricity and diamond cannot?
'DELOCALISED' electrons move along the layers
In 1985, another allotrope of carbon was discovered.
It is a molecule made up of 60 carbon atoms joined
together, named Buckminster fullerene.
Buckminster fullerene is one type of fullerene; it is a black solid.
The atoms in fullerenes are folded around to make a hallow sphere.
It is not a giant covalent structure, but a giant molecule in which the carbon atoms form pentagons and hexagons - in a similar way to a leather football.
Therefore they have a high melting and boiling point because there are allot of covalent bonds to be broken.
It is also used in lubricants.
Because of the presence of millions of strong covalent bonds in the giant covalent lattice; graphite has a high melting and boiling point. As allot of energy is required to overcome these bonds and separate the atoms.
Because of its smooth nature, graphite is used in pencils (pencil lead). As you move across a paper with your pencil you leave a trail of carbon atoms. It is also used as a lubricant.
Graphite can also conduct electricity.
More Properties and uses of Graphite.....