- 1 Introduction to Bonding
- 2 4.1 Ionic Bond
- 3 4.2 Covalent Bond
- 4 4.3 Intermolecular forces
- 5 4.4 Metallic bonding
- 6 4.5 Physical Properties
- 7 HL Material
Introduction to Bonding
Put simply, chemical bonds join atoms together to form more complex structures (like molecules or crystals). Bonds can form between atoms of the same element, or between atoms of different elements. There are several types of chemical bonds which have different properties and give rise to different structures.
Ionic bonds form between positive and negative ions (atoms). In an ionic solid, the ions arrange themselves into a rigid crystal lattice. NaCl (common salt) is an example of an ionic substance.
Covalent bonds are formed when atoms share electrons with each other. This gives rise to two structures: molecules and covalent network solids. Methane (CH4) is a covalent molecule and glass is a covalent network solid.
Whether two atoms form a covalent or ionic bond can be predicted from the atoms' electronegativities:
|Type of bond||Difference in atoms' electronegativities||Example|
|Non-polar covalent bond:||0.0-0.4||F2, CH4|
|Slightly polar bond:||0.5-0.9||Cl2O, NH3|
|Moderately polar bond:||1-1.3||CO2, SiCl4|
|Highly polar bond:||1.4-1.7||H2O, Al2Cl6|
|Slightly ionic bond:||1.8-2.2||NaCl, Al2O3|
|Ionic Bond:||2.3+||Na2O, CsF|
Metallic bonds occur between metal atoms. In a metallically bonded substance, the atoms' outer electrons are able to freely move around - they are delocalised. Iron is a metallically bonded substance.
Chemical bonding is one of the most crucial concepts in the study of Chemistry. In fact, the properties of materials are basically defined by the type and number of atoms they contain and how they are bonded together.
4.1 Ionic Bond
4.1.1 : Ionic bond - +ve (cations) and -ve (anions) ions are attracted to each other and form a continuous ionic lattice.
What are ions?
Ions are atoms or molecules which are electrically charged. Cations are positively charged and anions carry a negative charge. Ions form when atoms gain or lose electrons. Since electrons are negatively charged, an atom that loses an electron will become positively charged (similarly an atom that gains one or more electrons becomes negatively charged).
Description of Ionic Bonds
Ionic bonds form between positive and negatively charged ions. These oppositely charged ions attract each other and remain close together - they become ionically bonded. The law of electrostatic explains why this happens: opposite charges attract and like charges repel. When many ions attract each other, they form into large, orderly crystal lattices in which each ion is surrounded by ions of the opposite charge. When a metal forms an ionic bond with a non-metal, electrons are transferred from metal element to non-metal. When Ions are formed it is called Ionization
A diagram of an ionic solid should go here
Ionic bonds form when metals and non-metals chemically react. By definition, a metal is relatively stable if it loses electrons to form a complete valence shell and becomes positively charged. Likewise, a non-metal is relatively happy to gain electrons to complete its valence shell and become negatively charged. When a metal and a non-metal come into contact, the metal loses electrons by transferring them to the non-metal, which gains them. Consequently, ions are formed, which instantly attract each other and form an ionic bond.
4.1.2 : Group 1 metals form +1 ions, group 2 metals form +2 ions, metals in group 3 form +3 ions. Examples : Li+, Mg2+, Al3+...Greater ease of ionisation Li->Cs is due to the increased electron shielding of the nuclear attraction caused by additional inner shells of electrons. The easier atoms are to ionise, the more reactive they will be because less energy is required to ionise them, and so they react faster.
4.1.3 : Group 6 ions will form 2- ions, Group 7 ions will form 1- ions. Examples : O2-, Cl-...
4.1.4 : The transitions metals (elements from Ti to Cu, ignore Sc and Zn) can form multiple ions (ie Fe2+, Fe3+) (due to proximity of 4s and 3d shells)
4.1.5 : The ionic or covalent nature of the bonding in a binary compound is a result in the difference between their electronegativity...NaCl(s) is ionic, HCl(g) is (polar) covalent (also, covalent molecules tend to be gases/liquids, ionic tends to be solid...except network covalent which will be solid). In general, if the difference between electronegativities is greater than 1.7, the bond will be more than 50% ionic.
4.1.6 : Take the name of the group 1,2, or 3 metal and add...fluoride, chloride, bromide, iodide etc , oxide, sulfide etc...Nitride and phosphide...how exciting :)
4.2 Covalent Bond
One useful model of covalent bonding is called the Valence Bond model. It states that covalent bonds form when atoms share electrons with each other in order to complete their valence (outer) electron shells. They are mainly formed between non-metals (i.e. chlorine, sulfur, carbon etc.).
The Valence Bond Model
An example of a covalently bonded substance is hydrogen gas (H2). A hydrogen atom on its own has one electron – there is room for two to complete its valence shell. When two hydrogen atoms bond, each one shares its electron with the other, i.e. the electrons are attracted by two nuclei instead of just one and so releasing energy. Both atoms now have access to two electrons: they become a stable H2 molecule joined by a single covalent bond.
Double and Triple Bonds
Covalent bonds can also form between other non-metals, for example chlorine. A chlorine atom has 7 electrons in its valence shell — it needs 8 to complete it. Two chlorine atoms can share 1 electron each to form a single covalent bond. They become a Cl2 molecule.
Oxygen can also form covalent bonds; however, it needs a further 2 electrons to complete its valence shell (it has 6). Two oxygen atoms must share 2 electrons each to complete each other's shells, making a total of 4 shared electrons. Because twice as many electrons are shared, this is called a double covalent bond.
Furthermore, nitrogen has 5 valence electrons (it needs a further 3). Two nitrogen atoms can share 3 electrons each (6 in total) to make a N2 molecule joined by a triple covalent bond.
Electron Sharing and Orbitals
Carbon, contrary to the trend, does not share four electrons to make a quadruple bond. The reason for this is that the fourth pair of electrons in carbon cannot physically move close enough to be shared. The valence bond model explains this by considering the orbitals involved. Also more energy can be released by making 4 single bonds to 4 other carbon atoms to form a diamond structure.
Recall that electrons exist as clouds of electron density (orbitals) in atoms. The valence bond model works on the principle that orbitals on different atoms must overlap to form a bond. There are several different ways that the orbitals can overlap, forming several distinct kinds of covalent bonds.
The Sigma Bond
The first, and simplest kind of overlap is when two s orbitals come together. It is called a sigma bond (sigma, or 'σ', is the greek equivalent of 's'). Sigma bonds can also form between two p orbitals that lie pointing towards each other.
picture of sigma bonds- please press the no.1O sigma bond.jpg
The Pi Bond
The second, and equally important kind of overlap is between two parallel p orbitals. Instead of overlapping, head-to-head (as in the sigma bond), they join side-to-side, forming two areas of electron density above and below the molecule(delocalization). This type of overlap is referred to as a pi ('π', from the greek equivalent of p) bond.
4.2.1 : Covalent bonds are where two atoms each donate 1 electron to form a pair held between the two atoms...Such bonds are generally formed by atoms with little difference in electronegativity...ie C, H and O in organic chemistry.
4.2.2 : All electrons must be paired...Lewis diagrams are the element symbol with the outer (valence) shell of electrons left over and spare electrons pair up...in general C forms 4 bonds, N forms 3, O forms 2, halogens form 1, H forms 1...(Li would form 1, Be 2, and B 3 but they don't usually...metallic or ionic bonding)
4.2.3 : Electronegativity values range from 0.7 to 4...from bottom left to top right respectively (hydrogen falls B and C with a electronegativity of 2.1...
4.2.4 : When covalent molecules have a difference in electronegativity (between the two bonding atoms) then the pair will be held closer to the more electronegative atom...resulting in a small -ve charge on the more electronegative atom, and a small +ve charge on the other...results in polar bonds
4.2.5 : Shape of molecule with 4 electron pairs depends on number of lone pairs.
3 lone pairs -> linear, 2 lone pairs -> bent, 1 lone pair -> trigonal pyramid, No lone pairs -> tetrahedral
4.2.6 : The polarity of a molecule depends on both the shape and the polarity of the bonds...1) if there are no polar bonds, it's not polar. 2) if there are polar bonds, but the shape is symmetrical, it's not polar (think about it like 3D vector addition...if they add to zero, then it's not polar). 3) if there are polar bonds, and it's not symmetric, then the molecule is polar
4.3 Intermolecular forces
4.3.1 : van der Waal's forces -- Electrons will not be evenly spread around an atom/molecule at any given time, meaning the molecule will have a slight postitive charge on one end, and a negative at the other. This temporary state may cause attraction between two molecules, pulling them together (also known as London Dispersion Forces). Polar molecules, when properly oriented, will attract each other as a result of Dipole-Dipole forces. Dipole-Dipole forces are stronger than van der Waal's forces. Hydrogen bonding is when hydrogen is bonded to nitrogen, oxygen or fluorine, and a very strong dipole is formed, making the hydrogen very strongly positive. This hydrogen is then attracted to the lone pairs on other similar molecules--nitrogen, oxygen and fluorine all have lone pairs--forming a hydrogen bond, which is stronger than van der Waal's or dipole-dipole forces, but weaker than covalent bonding.
4.3.2 : Structural features -- Nonpolar molecules have van der Waal's forces only. This is also present in all other molecules, though its strength is often insignificant compared to the others. Polar molecules have dipole-dipole forces, which arise from polar bonds and asymmetry in molecules. Hydrogen bonds result from strongly delta positive hydrogen. This results in molecules with hydrogen bonding exhibiting stronger intermolecular forces, i.e. higher boiling/melting points etc. For example, H2O has a higher bp then H2S due to hydrogen bonding. Neutral molecules don't conduct electricity but some polar molecules exchange protons to form ions e.g. 2H2O makes H3O+ and OH-
4.4 Metallic bonding
4.4.1 : The metal atoms lose their outer electrons which then become delocalized, and free to move throughout the entire metal. These -ve delocalized electrons hold the metal cations together strongly. Since these electrons can flow, atoms with metallic bonding exhibit high electrical conductivity. The number of valence electrons involved in the bonding and the strength of the nucleus charge determines the strength. Unlike ionic bonding, distorting the atoms does not cause repulsion so metallic substances are ductile (can be stretched into wires) and malleable (can be made into flat sheets). The free moving electrons also allow for high thermal conductivity, and the electrons can carry the heat energy rather than it being transferred slowly through atoms vibrating.
Metallic bonds occur among metal atoms. A sea of valence electrons surrounds positive metal ions. The electrons are free to move throughout the resulting crystal. The delocalized nature of the electrons explains a number of unique characteristics of metals: they are good conductors of electricity, they are ductile, meaning they can be made into wires, and they are malleable, meaning they can easily be hammered into thin sheets.
4.5 Physical Properties
Melting and Boiling point: High with Ionic, Metallic Bonding and Network Covalent. Low with Covalent Molecular Bonding.
Volatility: Covalent Molecular Substances are volatile, others are not.
Conductivity: Metallic substances conduct. Polar molecular substances conduct, non-polar ones don't. Ionic substances do conduct when molten or dissolved in water but never when solid.
Solubility: Ionic substances --> generally dissolve in polar solvents (like water). Metallic substances --> soluble in liquid metal. Non-polar molecules are generally soluble in non-polar solvents, and polar in polar. Organic molecules with a polar head --> Short chain molecules are solubility in polar solvents but long chains can eventually outweigh the polar 'head' and will dissolve in non-polar solvents.
Ionically bonded substances typically have the following characteristics.
- High melting point (solid at room temp)
- Brittle (can shatter)
- Some dissolve in water
- Conduct electricity when dissolved or melted
- Typically stronger than covalent bonds.
Topic 14 is the additional HL material for Topic 4.