Elements will form compounds in order for them to become stable. Elements don’t combine to form compounds at r&om, many different factors will come into this, the main & most important one being the electronic configuration, meaning the finally electronic configuration of the atoms must be 2n2, where n=maximum number of electrons a shell can hold (1). Most non-metals that do this In two ways, by reacting with metals, creating ionic compounds. Or by forming molecular compounds with metalloids & other non-metals, known as covalent compounds. (2) in most cases, the full outer shell of electrons will be made up of eight electrons, expect in the case of hydrogen & helium, where a full outer shell of electrons in made up of two electrons. Fluorine is the most reactive element & therefore forms compounds most readily, only two elements don’t form compounds with fluorine, these are helium & neon. (3)
Ionic bonding is the bonding between a metal & a non-metal. This occurs when there is a transfer of electron(s). The transfer occurs due to the ionisation energy & the electron affinity of the two atoms bonding. Ionisation energy is defined as the energy needs to remove one of more electrons from a neutral atom. (4) electron affinity is defined as the energy that is released when one or more electrons is added to a neutral atom.(5) Ionic bonds will happen between two elements where one had a low ionisation energy that is bonded to an element with a has a high electron affinity. This means that the transfer of electrons happens very easily. The atom with the low ionisation energy will easily give up one or maybe more electrons in order for it to achieve a full outer shell & the atom with the high electron affinity will happily accept these electrons in order for it to also achieve a full outer shell. (6) ionic structures typically have high melting & boiling points. They don’t conduct heat when solid but will when in dissolved or molten, due to there being no delocalised electrons when in the crystal structure but when the molecule is melted or molten, the ions making up the structure will become the mobile valence delocalised electrons required to carry charge. Ionic structures are hard but they are also brittle, if a force is applied, a layer of the structure will move, causing repulsion & the structure will spring apart. (7) there are four main types of ionic crystal structure;
- rock salt formation (NaCl)- this is where there Cl- will make up the corners & face centres of the structure & the Na+ make up the body & the edge of centres.
- Zinc blended type (ZnS)- this is where each Zn2+ is surrounded by 4 S2- & vice versa.
- Fluorite type (CaF2)- the is when each Ca2+ is surrounded by 8 F- & each F- is surrounded by 4 Ca2+. Anti-fluorite type- this is where the positive ion Is surrounded by 4 negative ions & the negative ions are surrounded by 8 positive ions.
- Caesium Chloride type (CsCl)- this has Cs+ at the body centre with Cl- at the corners & vice versa. (8)
Metallic bonding is the bonding between a lattice of cations & a sea of mobile valence delocalised electrons. (9) these cations will have an electronic configuration of 2n2. There are two key conditions for metallic bonding;
- a low ionisation energy
- Sufficient valence electrons
The strength of a metallic bond is reliant on three things;
- The number of valence electrons
- Charge on the nucleus
- Shielding
Metallic structures all have very similar properties. They conduct electricity & heat, due to the delocalised elections that move carrying heat & charge. Due to the non-directional bonds within the structure, metals are considered to be malleable & ductile. Non-directional meaning the bonds are not with tied to specific electrons but the collective electrons surrounding.(10)
Localised & delocalised bonds are both present in the formation of covalent bonding. Localised bonds, better known as sigma & pi bonds are the overlapping of two orbitals from two different atoms. A sigma bond is the overlapping of the two orbitals on the internuclear axis & a pi bond is the bonding between unhybridised p-orbitals both above & below a sigma bond, making two pi bonds. sometimes, hybridisation will occur for some atoms to form molecules. When this happens, they will create & occupy newly created orbitals. there are several levels to this, first level is sp.Where one s orbital will overlap with one p orbital, taking on a linear orientation. The next level is sp2 which is between one s orbital & two p orbitals which will create three hybrid orbitals. These orbitals will have 1/3 “s” characteristics & 2/3 “p” characteristics. The third level of hybridisation is sp3, is between one s orbital & three p orbtials & will create four hybrid orbitals each with 1/4 “s” character & 3/4 “p” character. A delocalised bond or electron is one that has resonance & therefore no fixed positioning. The most obvious example of this would be In benzene where there is no double or single bonds in the ring structure but there is instead six delocalised electrons that allow the structure increased stability. This theory is known as molecular orbital theory. (11)
There are three main types of intermolecular forces- ion dipoles, hydrogen bonds & van Der Waals forces. Ion dipoles are the interactions between an ion & a polar molecule. This allows the charge of the ion to line up with the charge of polar molecule, forming an interaction which is typically found in ionic structures. Hydrogen bonding is the strongest force, which is between the lone pair of electron on an electronegative species & a hydrogen atom that is bonded to either N, O or F. This bond is very strong but it will weaken the bond between the hydrogen & its other molecule. (12) van Der Waals are an umbrella term from four different types of forces, dipole-dipole, ion induced dipole, dipole induced dipole & London dispersion. There are in the descending order of strength. London dispersion is present in all molecules but is the only force in non-polar molecules. Van der Waals are what play a part in the state of a substance. (13)
Compounds are formed in order to become stable through many different types of bonding. The strongest of which is the ionic bond. The four different ionic structures, forming different lattices make it much stronger than the next strongest type, covalent. Although valence bond theory is useful in the explanation of basic bonding, it doesn’t give any explanation of resonance which gives rise to molecular orbital theory. the weakest type of intramolecular is metallic bonding, due to its non-directional bonding. This types of bonding can also be explained using the b& theory which I didn’t discuss in my essay but it gives a better explanation of the rise of the metallic properties. Intermolecular forces are at play in the interaction between these different types of structures with one another. These are found widely in various biological system, most notably in the hydrogen bonding within DNA. The hydrogen bonding occurs between two nucleotides that will fold in various different ways depending on how the h-bonds interact. (14).
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