Introduction to Bonding

Put simply, chemical bonding joins atoms together to form more complex structures (like molecules or crystals). Bonding can occur between atoms of the same element, or between atoms of different elements. There are several types of chemical bonding which have different properties and give rise to different structures.

Ionic bonding occurs between positive and negative ions (charged atoms). This type of bonding seldom occurs between just two atoms, but typically leads to the formation of an ionic solid, in which the ions arrange themselves into a rigid crystal lattice. NaCl (common salt) is an example of an ionic substance.

Covalent bonding occurs when atoms share electrons with each other. This gives rise to two types of structures: molecules and covalent network solids. Methane (CH₄) is a covalent molecule and glass is a covalent network solid. In molecules we can distinguish individual covalent bonds between pairs of atoms.

Whether the interaction between two atoms has a covalent or ionic character 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 bonding occur between metal atoms. In a metallically bonded substance, the atoms' outer electrons are able to freely move around - they are delocalised. In principle, all electrons can behave like delocalized waves, but in substances with covalent or ionic bonding the number of accessible energy states equals the number of electrons and that makes this wave character less apparent. In metals there are far more accessible states than electrons and that leaves the electrons free to move and conduct electrically. Aluminum, tin, lead, silver are all metallically bonded elements, but there are also compounds with metallic bonding, such as alloys and intermetallic compounds.

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 Bonding

4.1.1 : Ionic bonding - +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 are negatively charged. 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 Bonding

Ionic bonding occurs 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

Formation of Ions

Ions result when metals and non-metals chemically react. Due to its low ionization energy, a metal atom is destabilized only a little if it loses electrons to form a complete valence shell and becomes positively charged. Likewise, a non-metal is stabilized strongly by gaining 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. An ionic bond is the electrostatic attraction between cations and anions due to Coulombic forces. The process of attraction typically does not stop with two ions, but proceeds to involve a great many of them that stack in a solid lattice structure, in which individual 'bonds' or individual molecules cannot be distinguished.

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 more easily.

4.1.3 : Group 6 ions will form 2- ions, Group 7 ions will form 1- ions. Examples : O^2-, 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, nitride or phosphide...

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.

Diagram of a covalent bond between hydrogen atoms

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 positive 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.

Characteristics

$NaCl$ . $KNO_{3}$

Ionically bonded substances typically have the following characteristics.

High melting point (solid at room temp)
Hard
Brittle (can shatter)
Some dissolve in water
Conduct electricity when dissolved or melted
Typically stronger than covalent bonds.

HL Material

Topic 14 is the additional HL material for Topic 4.

IB Chemistry/Bonding

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