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  • Revision:Chemistry unit 1 3 - structure and bonding

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A Level Chemistry Unit 1.3 - Structure and bonding

Contents

Ionic bonding

Ionic bonds are a type of chemical bond based on electrostatic forces between two oppositely-charged ions. In ionic bond formation, a metal donates an electron, due to a low electronegativity to form a positive ion or cation. Often ionic bonds form between metals and non-metals. The non-metal atom has an electron configuration just short of a noble gas structure. They have high electronegativity, and so readily gain electrons to form negative ions or anions. The two or more ions are then attracted to each other by electrostatic forces. Such bonds are stronger than hydrogen bonds, but similar in strength to covalent bonds.

Ionic bonding occurs only if the overall energy change for the reaction is favourable when the bonded atoms have a lower energy than the free ones. The larger the resulting energy change the stronger the bond.

Pure ionic bonding is not known to exist. All ionic bonds have a degree of covalent bonding or metallic bonding. The larger the difference in electronegativity between two atoms the more ionic the bond. Ionic compounds conduct electricity when molten or in solution. They generally have a high melting point and tend to be soluble in water.

Covalent Bonding

Covalent bonding is a description form of chemical bonding that is characterized by the sharing of one or more electrons between two atoms. In general bonds are defined by a mutual attraction that holds the resultant molecule together. Often bonding occurs in such a way that the the outer electron shells of the participating atoms becomes filled. Such bonds are always stronger than the intermolecular hydrogen bond and similar in strength to or stronger than the ionic bond.

In contrast to the ionic and metallic bond, the covalent bond is directional, i.e. the bond angles have a great impact on the strength of the interaction. This impact arises because covalent bonds are formed by the overlap of atomic orbitals, with greater overlap producing a greater strength of interaction. Atomic orbitals all have highly directional character, resulting in a highly directionally-dependent interactions in bonding.

Covalent bonding most frequently occurs between atoms with similar electronegativities. For this reason, non-metals tend to engage in covalent bonding more readily since metals have access to metallic bonding.

Dative bonding

A dative covalent bond (also known as coordinate covalent bond) is a special type of covalent bond in which the shared electrons "come from" one of the atoms only. Once the bond has been formed, its strength is no different from that of a covalent bond. The process of forming a dative bond is called coordination. The electron donor acquires a positive formal charge, while the electron acceptor acquires a negative formal charge.

A compound that contains a lone pair of electrons is capable of forming a dative bond. Dative bonds can be found in many different substances, such as in simple molecules like carbon monoxide \mathsf{(CO)} or the ammonium ion \mathsf{(NH_4)}. Dative bonds are also formed by electron deficient compounds, such as beryllium chloride \mathsf{(BeCl_2)}. The Beryllium atom in this compound tends to bind two additional chlorine atoms, in which every beryllium atom is bonded to four chlorine atoms, two with normal covalent bonding, and the other two with dative bonds.

Dative bonding can also be found in coordination complexes involving metal ions, especially if they are transition metal ions. In such complexes, substances in a solution donate their free pairs of electrons to the metal ion, which accepts the electrons. Dative bonds form and the resulting compound is called a coordination complex, while the electron donors are called ligands. A common ligand is water \mathsf{(H_2O)}, which will form coordination complexes with metal ions, like \mathsf{Cu^{2+}}, which will form \mathsf{[Cu(H_2O)_6]^{2+}} in aqueous solution.

A dative bond is sometimes represented by an arrow pointing from the donor of the electron pair to the acceptor of the electron pair.

Metallic bonding

Metallic bonding is the bonding within metals. It involves the delocalised sharing of free electrons among a lattice of metal atoms. Thus, metallic bonds may be compared to molten salts.

Metal atoms typically contain a high number of electrons in their valence shell compared to their period or energy level. These become delocalised and form a sea of electrons surrounding a giant lattice of positive ions. The surrounding electrons and the positive ions in the metal have a strong attractive force between them. This means that more energy is required to negate these forces. Therefore metals often have high melting or boiling points. The principle is similar to that of ionic bonds.

Metallic bonding is non-polar, because there is no difference in the electronegativity among the atoms participating in the bonding interaction, and the electrons involved in that interaction are delocalised across the crystalline structure of the metal.

The metallic bond accounts for many physical characteristics of metals, such as strength, malleability, ductility, conduction of heat and electricity, and lustre. Due to the fact that the electrons move independently of the positive ions in a sea of negative charge, the metal gains some electrical conductivity. It allows the energy to pass quickly through the electrons generating a current.

Metallic bonding is the electrostatic attraction between the metal atoms or ions and the delocalised electrons. This is why atoms or layers are allowed to slide past each other, resulting in the characteristic properties of malleability and ductility.

Pi(π) and sigma (σ) bonds

Sigma bonds (σ bonds) are a type of covalent chemical bond. Sigma bonds are the strongest type of covalent bonds. Electrons in sigma bonds are sometimes referred to as sigma electrons. The symbol σ is the Greek letter for s.

Pi bonds (π bonds) are also chemical bonds of the covalent type. The Greek letter π in their name refers to p orbitals, since the orbital symmetry of the pi bond is the same as that of the p orbital when seen down the bond axis. P orbitals usually engage in this sort of bonding. However, d orbitals can engage in pi bonding also. Pi bonds are usually weaker than sigma bond. Although the pi bond by itself is weaker than a sigma bond, pi bonds are most often found in multiple bonds together with sigma bonds and the combination is stronger than either bond by itself.

Intermediate nature of bonds

Difference in electronegativities

The above examples only show the ideal types of bond between two atoms, in reality most covalent bonds are in fact not totally covalent bonds and the is no such thing as a perfect ionic bond as none exist, (although some bonds do get very close). In reality the type of bond is more of a sliding scale based on the electronegativity of the atoms involved in the bonding process.

Electronegativity is a measure of the ability of an atom or molecule to attract electrons in the context of a chemical bond. The type of bond formed is largely determined by the difference in electronegativity between the atoms involved. Atoms with similar electronegativities will share an electron with each other and form a covalent bond. However, if the difference is too great (usually >2), the electron will be permanently transferred to one atom and an ionic bond will form. Furthermore, in a covalent bond if one atom pulls slightly harder than the other, a polar covalent bond will form.

Polarisation of bonds

Positive ions can have the effect of polarising (electrically distorting) nearby negative ions. The polarising ability depends on the charge density in the positive ion. Polarising ability increases as the positive ion gets smaller and the number of charges gets larger. As a negative ion gets bigger, it becomes easier to polarise. For example, in an iodide ion, I-, the outer electrons are in the 5-level - relatively distant from the nucleus. A positive ion would be more effective in attracting a pair of electrons from an iodide ion than the corresponding electrons in, say, a fluoride ion where they are much closer to the nucleus. Aluminium iodide is covalent because the electron pair is easily dragged away from the iodide ion. On the other hand, aluminium fluoride is ionic because the aluminium ion can't polarise the small fluoride ion sufficiently to form a covalent bond.

Polar bonds and polar molecules

In a complex molecule even if the bonds in the molecule are all polar bonds, the molecule overall may not be polar because of the slight charges cancelling out. This is true in molecules such as \mathsf{CHCl_4}, but not in molecules such as \mathsf{CHCl_3} where the charges do not cancel out.

Intermolecular forces

Hydrogen bonding

Hydrogen bonding occurs when a hydrogen atom is covalently bonded to a small highly electronegative atom such as nitrogen, oxygen, or fluorine. The result is a dipolar molecule. The hydrogen atom has a partial positive charge δ+ and can interact with another highly electronegative atom in an adjacent molecule (again N, O, or F). This results in a stabilizing interaction that binds the two molecules together.

Dipole-Dipole interactions

Dipole-dipole interactions are the force that occur between two molecules with permanent dipoles (spatially oriented δ+ within a molecule). These work in a similar manner to ionic interactions, but are weaker because only partial charges are involved.

Van der Waals forces

These involve the attraction between temporarily induced dipoles in non-polar molecules (often disappear within an instant). This polarization can be induced either by a polar molecule or by the repulsion of negatively charged electron clouds in non-polar molecules.

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