Atomic Bonding

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1.1
The atoms, during bond formation, may lose or gain electrons (valence electrons) in order to achieve a stable state, or technically speaking, a stable electron configuration. Usually metal atoms lose electrons and non-metals gain electrons in order to achieve electron stability. When dealing with bond formation (Ionic bonding for example) we need to analyse the outer shell of the atom. Metals usually present 1, 2 or 3 electrons in their outer shell therefore they have to give them away to achieve stability. Vice versa non-metals have 5, 6 or 7 electrons in their outer shell and they need to receive more electrons in the outer shell to be stable. 1.2

Ionic bonding occurs between metal and non-metal atoms and consists in gaining and losing electrons to reach electron stability. Metal atoms lose electrons (negative charge) becoming ions, positively charged. Non-metal atoms gain electrons to achieve electron stability. Sodium has one valence electron in its outer shell, so it needs to give that electron away in order to achieve stability. Vice versa, Chlorine atom has 7 valence electrons in its outer shell and it needs to gain one electron (from sodium atom) to achieve electron stability.

1.3
Covalent bond occurs between non-metals and consists in a shared pair of electrons, from the outer shell, between two atoms. Each of the two atoms, involved in the bonding, provides an electron. Usually non-metals present 4 or more electrons in their outer shell and a lot of energy is needed to remove electrons in order to form bonds. That’s one of the reasons why covalent bonding is used between non-metals. The electrons shared are called molecules. When the pair of electrons is shared from only one of the two atoms involved in the reaction, then it’s called dative covalent bond. It is represented by a short arrow going from the electron giving both electrons towards the one providing neither.

2.1

Metallic bond is a strong electrostatic attraction between a sea of delocalised electrons and a lattice of positive ions, called cations. This strong electrostatic attraction holds the structure together, avoiding that positively charged ions repel the negatively charged electrons. Because of the fact that electrons are free to move metals have the characteristic to be good conductor of heat and electricity. The metals are also ductile, which it means that they can be stretched out to make a wire. They are malleable too. This means that they can pressed and beaten to change their shape.

http://www.ndt-ed.org/EducationResources/CommunityCollege/Materials/Structure/metallic.htm 3.1
Before defining the van der Waals’ forces we need to talk about intermolecular forces. These are forces holding molecules together and to melt or boil a substance we have to overcome these forces. The strongest these forces are, and the more energy will be required to break the substance’s bonds in order to reach its melting or boiling point. The weakest of the intermolecular forces are called van der Waals’ forces, which are forces of attraction between a temporary dipole on one molecule, and a induced dipole on an adjacent molecule. Temporary diploes are a surplus of electrons in one of the atoms, making it to appear asymmetrical in only a very short time because the electrons are in constant motion. If we consider an adjacent molecule then we can notice that the electrons are repelled by the negative side of the dipole and attracted by the positive side.

“Mill Hill County High School”, Scholarly articles on Atomic Bonding found on Google.

3.2

Hydrogen Bonding is the strongest intermolecular force and it can be defined as an attractive force between the Hydrogen (bonded to oxygen, nitrogen, or fluorine) and an electronegative atom of an adjacent molecule. This happens because The Hydrogen has almost null electron density and this make it able to bond with electronegative atoms on neighbouring molecules,...
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