Covalent Compounds
Covalent bonds form between non-metal atoms. In order to reach stability, non-metal atoms want to gain electrons due to high electronegativities, so instead of stealing electrons, as in ionic bonding, or giving them away as metals do, non-metals share. How sweet, they are so kind.
Atoms share electrons so that they have a full valence shell. Remember Nobel Gases won't form covalent bonds/molecules as they already have full valence shells.
Discrete Covalent Molecules:
Exist as a 2 or 3D layer lattice where each atom os covalently bonded to other atoms. Their properties vary greatly due to variation in structure formed.
Covalent Layer (2D) Lattice:
each layer is held together by covalent bonds. The layers are weakly held by dispersion forces. - Graphite is an example.
Covalent Network Lattice:
Giant structures of countless atoms in which no individual molecules exist, covalently bonded into a 3D shape. - Diamond and quartz are examples.
Showing Covalent Bonding:
Electron Dot Diagrams (Lewis Diagrams):
Outer shell electrons are represented by dots, to show how the covalent bond has occurred.
Here we have Chlorine, with its symbol in the middle, and 7 valence electrons on the outside.
The electrons in pairs are known as non-bonding electrons or lone pairs, and the ones that are single are bonding electrons, as these are the ones that will be shared.
Here's a nice little table that gives some dot diagrams of various elements.
This is ammonia, a covalent compound of one nitrogen, bonded to three hydrogen molecules. It is also shown as a space filling model, and stick diagram/structural diagram.
Multiple Bonds:
While it is very common for two atoms to be joined by just one bond (one set of shared electrons) it is also possible for two atoms to share two or even three pairs of electrons, forming double and triple bonds. These bonds are stronger and shorter.
Molecular Shape:
Electron dot diagrams only give information of the molecules shape on a 2D level, however, these molecules are 3D. The shape of covalent compounds is affected by the location of valence electrons. As electrons have like charges, they repel. This is known as Valence Shell Electron Pair Repulsion Theory.
Naming Covalent Compounds:
1. The first element is named in full.
2. The second element as if it is an ion (shortened + ide)
3. The number of each is identified by a prefix
4. if the name of the element begins in a vowel, then the o or a is dropped from the prefix.
Example: Ph2O5 becomes diphosphorus pentoxide
If the covalent compound is made up of only one type of atom/element, then they are known simply by their elemental name.
Example: H2 is just hydrogen and N2 is just nitrogen.
Polar Molecules:
The strength - due to an atom's electronegativity - of the pull on electrons by an atom can cause polar bonds and polar molecules to be formed. Shared electrons are pulled towards the more electronegative atoms, however this can be off set if the molecule is symmetrical.
Water is polar, but CO2 is non-polar. Although both contain polar bonds, carbon dioxide's symmetrical shape balances it out.
Intramolecular Forces:
Forces within a molecule, covalent bonds are examples of this.
Intermolecular Forces:
Forces between different molecules. This allows for the formations of liquids and solids.
Dispersion (Van der Waals) Forces:
Between every molecule, even the atoms of Nobel gases.
They are the weakest intermolecular force, and they re the attraction of electrons by nuclei of neighbouring molecules/atoms.
Strength is dependent on the number of atoms in the molecule (more = stronger) and shape (the closer the molecules can get to each other, the stronger the dispersion force)
Instantaneous dipoles: electrons may nr evely distributed, but ofr a fraction of a second, randomly and constantly, the electrons end up at one end, forming dipoles (charged ends)
Induced Dipoles: Electrons are evenly spaced but moved due to being repelled by the dipole next to it.
Dipole - Dipole Interactions:
2nd weakest, and occur only between polar molecules: the positive end of one, attracts the negative end of another. Substances that contain dipole forces have higher melting and boiling points that those with only dispersion forces, as more energy is required to break these bonds.
Hydrogen Bonding:
While it is actually a type of dipole-dipole interactions, hydrogen bonds are far stronger, and classified on their own. When hydrogen bonds to a more electronegative atom, its lone electron is drawn away, leaving its nucleus exposed. A strong dipole forms as other molecules can approach close. This is especially evident when H bonds with nitrogen, oxygen or fluorine (NOF) the most electronegative elements. It is hydrogen bonding that causes many of the physical properties of water, and why it has a high melt/boil point compared to other molecules of similar size, and why it expands upon freezing.
Atoms share electrons so that they have a full valence shell. Remember Nobel Gases won't form covalent bonds/molecules as they already have full valence shells.
Discrete Covalent Molecules:
- Exist as individual molecules: CO2, H2O, C6H12O6, HCl
- Many are gases / liquids at room temperature
- Generally low melting/boiling point
- poor conductors of electricity - no free charged particles, unless dissolved in water
- like dissolves in like (polar molecules will dissolve in polar molecules, same as non-polar will in non-polar)
Exist as a 2 or 3D layer lattice where each atom os covalently bonded to other atoms. Their properties vary greatly due to variation in structure formed.
Covalent Layer (2D) Lattice:
each layer is held together by covalent bonds. The layers are weakly held by dispersion forces. - Graphite is an example.
Covalent Network Lattice:
Giant structures of countless atoms in which no individual molecules exist, covalently bonded into a 3D shape. - Diamond and quartz are examples.
Showing Covalent Bonding:
Electron Dot Diagrams (Lewis Diagrams): Outer shell electrons are represented by dots, to show how the covalent bond has occurred.
Here we have Chlorine, with its symbol in the middle, and 7 valence electrons on the outside.
The electrons in pairs are known as non-bonding electrons or lone pairs, and the ones that are single are bonding electrons, as these are the ones that will be shared. Here's a nice little table that gives some dot diagrams of various elements.
This is ammonia, a covalent compound of one nitrogen, bonded to three hydrogen molecules. It is also shown as a space filling model, and stick diagram/structural diagram.
Multiple Bonds:
 |
| Carbon dioxide contains two sets of double bonds |
Molecular Shape:
Electron dot diagrams only give information of the molecules shape on a 2D level, however, these molecules are 3D. The shape of covalent compounds is affected by the location of valence electrons. As electrons have like charges, they repel. This is known as Valence Shell Electron Pair Repulsion Theory.
Naming Covalent Compounds:
1. The first element is named in full.
2. The second element as if it is an ion (shortened + ide)
3. The number of each is identified by a prefix
4. if the name of the element begins in a vowel, then the o or a is dropped from the prefix.
Example: Ph2O5 becomes diphosphorus pentoxide
If the covalent compound is made up of only one type of atom/element, then they are known simply by their elemental name.
Example: H2 is just hydrogen and N2 is just nitrogen.
Polar Molecules:
The strength - due to an atom's electronegativity - of the pull on electrons by an atom can cause polar bonds and polar molecules to be formed. Shared electrons are pulled towards the more electronegative atoms, however this can be off set if the molecule is symmetrical.
Water is polar, but CO2 is non-polar. Although both contain polar bonds, carbon dioxide's symmetrical shape balances it out.
Intramolecular Forces:
Forces within a molecule, covalent bonds are examples of this.
Intermolecular Forces:
Forces between different molecules. This allows for the formations of liquids and solids.
Dispersion (Van der Waals) Forces:
Between every molecule, even the atoms of Nobel gases.
They are the weakest intermolecular force, and they re the attraction of electrons by nuclei of neighbouring molecules/atoms.
Strength is dependent on the number of atoms in the molecule (more = stronger) and shape (the closer the molecules can get to each other, the stronger the dispersion force)
Instantaneous dipoles: electrons may nr evely distributed, but ofr a fraction of a second, randomly and constantly, the electrons end up at one end, forming dipoles (charged ends)
Induced Dipoles: Electrons are evenly spaced but moved due to being repelled by the dipole next to it.
Dipole - Dipole Interactions:
2nd weakest, and occur only between polar molecules: the positive end of one, attracts the negative end of another. Substances that contain dipole forces have higher melting and boiling points that those with only dispersion forces, as more energy is required to break these bonds.
Hydrogen Bonding:
While it is actually a type of dipole-dipole interactions, hydrogen bonds are far stronger, and classified on their own. When hydrogen bonds to a more electronegative atom, its lone electron is drawn away, leaving its nucleus exposed. A strong dipole forms as other molecules can approach close. This is especially evident when H bonds with nitrogen, oxygen or fluorine (NOF) the most electronegative elements. It is hydrogen bonding that causes many of the physical properties of water, and why it has a high melt/boil point compared to other molecules of similar size, and why it expands upon freezing.
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