Peter's Physics Pages

An introduction to Electricity and Strength of Materials with Peter Eyland

Lecture 7 (Types of Chemical Bonds)

In this lecture the following are introduced:
•The Classification of Chemical Bonds
•The Periodic Table with Electron Configurations
•Ionic Bonds
•Covalent Bonds
•Metallic Bonds
•Electric Dipoles
•Hydrogen Bonds
•Van der Waals Bonds

The Classification of Chemical Bonds

There are two major bond classifications, each with identifiable sub-groups:

The Periodic Table gives a guide to the type of primary bond that will form.

The Periodic Table with Electron Configurations

 For a background on the periodic table, see Western Oregon University Atoms have their electrons confined in potential energy wells around them, but not at any energy, there are only a few discrete energy levels allowed. These levels are like steps on a ladder where your feet cannot be supported at any height, but only on the rungs. The electrons then cannot be at intermediate energies, they will fall down to the next "rung" on the energy ladder, grouping in three dimensional "shells". The periodic table arises because only certain numbers of electrons are allowed in each shell, as also illustrated in the diagram. Electrons are held at upper levels because they cannot all fall down to the bottom. The valence electrons are in the outer (highest) energy levels and are the main determiners of chemical activity and bonding. "Valence" comes from a Latin word meaning "power" and used for the "combining power" of an atom, i.e. how many neighbouring atoms it can bond with.

As the outer shells determine most of the chemical properties, the Periodic Table below shows only the electron numbers in the two outer shells, the colours indicate how shells are filled.

 S Filling 2 S levels (l=0) at the outer level P Filling 6 P levels (l=1) at the outer level D Filling 10 D levels (l=2) at a level deeper F Filling 14 F levels (l=3) two levels deeper

In the outermost shell the general pattern is an increase in electron numbers from 1 to 8 across the table.
These form "families" of elements with similar chemical properties. It is a little like music with "octaves". "Family members" are down the table in higher octaves.

The families are designated with Roman numerals and further divided into A and B.
In light blue, there is IA, IIA, then in green there is IIIA, IVA, VA, VIA, VIIA and 0. Family 0 has one light blue element at the top.
B families are in yellow and start with IIIB, IVB, VB, VIB, VIIB, VIIIB, then IB and IIB.
 IA IIA IIIB IVB VB VIB VIIB {- VIIIB -} IB IIB IIIA IVA VA VIA VIIA 0 1 H1 He2 2 Li2.1 Be2.2 B2.3 C2.4 N2.5 O2.6 F2.7 Ne2.8 3 Na.8.1 Mg.8.2 Al.8.3 Si.8.4 P.8.5 S.8.6 Cl.8.7 Ar.8.8 4 K.8.1 Ca.8.2 Sc.9.2 Ti.10.2 V.11.2 Cr.13.1 Mn.13.2 Fe.14.2 Co.15.2 Ni.16.2 Cu.18.1 Zn.18.2 Ga.18.3 Ge.18.4 As.18.5 Se.18.6 Br.18.7 Kr.18.8 5 Rb.8.1 Sr.8.2 Y.9.2 Zr.10.2 Nb.12.1 Mo.13.1 Tc.14.1 Ru.15.1 Rh.16.1 Pd.18 Ag.18.1 Cd.18.2 In.18.3 Sn.18.4 Sb.18.5 Te.18.6 I.18.7 Xe.18.8 6 Cs.8.1 Ba.8.2 La*.9.2 Hf.10.2 Ta.11.2 W.12.2 Re.13.2 Os.14.2 Ir.17 Pt.17.1 Au.18.1 Hg.18.2 Tl.18.3 Pb.18.4 Bi.18.5 Po.18.6 At.18.7 Rn.18.8 7 Fr.8.1 Ra.8.2 Ac**.9.2 Rf Ha Sg Ns Hs Mt Uun Uuu Uub Uut Uuq Uup Uuh Uus Uuo Ce*.8.2 Pr.8.2 Nd.8.2 Pm.8.2 Sm.8.2 Eu.8.2 Gd.9.2 Tb.8.2 Dy.8.2 Ho.8.2 Er.8.2 Tm.8.2 Yb.8.2 Lu.9.2 Th**.10.2 Pa.9.2 U.9.2 Np.9.2 Pu.9.2 Am.9.2 Cm.9.2 Bk Cf Es Fm Md No Lr

The inert gases of Group 0, namely Helium, Neon, Argon, Krypton, Xenon, and Radon, each have 8 valence electrons and form few chemical compounds. (Helium is an exception with only 2 outer electrons). This suggests that 8 valence electrons form a stable electronic structure where an atom tends not to share electrons with other atoms.

An element will be chemically reactive if it can get to the stable electronic configuration of an inert gas, either
• by losing one or two electrons to another atom, or
• by gaining one or two electrons from another atom (at most three), or
• by sharing three or more electrons.

Three Primary Bonds

The three types of primary bonding reflect these ways in which atoms can group together by gaining or losing or sharing electrons, so they can get inert gas electron configurations.

(P1) Ionic Bonds

Atoms near the left or right sides of the periodic table can loose or gain 1 (or 2) electrons to form charged "ions". For example, a Sodium atom (row 3, column IA) can loose one electron to have 8 valence electrons and become a positively charged "cation". A Chlorine atom (row 3, column VIIA) can gain one electron to have 8 valence electrons and become a negatively charged "anion". These two ions then will be attracted to each other by non-directional electrostatic force and form an ionic (or electrovalent) bond.

Note: cations are + like the t, anions are - like the i.
 When large numbers of such ion pairs come together an ionic solid is formed. Common salt (NaCl) is an ionic solid which has the cubic structure shown on the right.

In ionic solids:
• there is a charge requirement for stacking atoms. Each ion must have nearest neighbours of opposite charge. (For the cubic array above, each ion has six nearest neighbours of opposite charge.)
• there are no directional requirements, so stacking depends on meeting charge and size requirements and the bonding can be at any angle.
• there are long range requirements because they attract or repel other ions beyond the nearest and next-nearest neighbours.

(P2) Covalent bonds

Atoms at the centre of the periodic table (group IVA) have 4 valence electrons. It is difficult to completely lose or gain this many electrons so by compromise they end up sharing electrons.

 In the diagram on the right, the central atom is missing the 4 electrons to form an inert gas configuration, but can borrow them for a while from like-minded neighbours to form an electron "cloud" between the two atoms. The atomic cores are then attracted to the negative electron cloud between them, forming a covalent bond. The repulsion between electron "clouds" maximises the angles between the bonds. In 3-dimensions, four equally spaced bonds form a tetrahedral structure. In the diagram, the blue atom is at the centre of the tetrahedron forming four bonds with green atoms at the tetrahedral vertices. A sloping light green triangular face is highlighted between two base atoms and the top atom. When large numbers of such atoms come together, all sharing some of their electrons to fill outer shells, a covalent solid is formed. A notable example of a covalent solid is diamond, which is a three-dimensional carbon network built up from tetrahedral units.

In covalent bonding:
• there are no charge requirements. Each atom does not have to have nearest neighbours of opposite charge.
• there are strong directional requirements which determine structural geometries.
• there are no long range requirements. Bonds are only between nearest-neighbour atoms sharing electrons.

(P3) Metallic Bonds

 A lot of metals fall in the yellow area of the periodic table shown. They share electrons in a different way to covalent bonding. The diagram (top right) shows a single atom potential well with all its electrons bound to it. Under it (middle right), is the example of five metal atoms brought closely together so that their potential wells overlap. At the bottom right, a combined potential well is formed as the overlapping potentials interact and are added together. The outermost electron levels now combine to form levels that extend throughout the whole solid.

In thes outer combined levels a few of the electrons are shared by all the atoms of the solid. Most of the electrons will still be confined to their own atoms as they are further down the wells. The outer electrons are "nearly free" in that they are free from individual atoms but not free to leave the solid as they still have negative potential energy.

Metallic bonding occurs between the positive atom cores and the "nearly free" electrons.
In metallic bonding:
• there are no charge requirements,
• there are no directional requirements, and
• there are long range effects.

 This means that in metallic bonding the atoms pack together as closely as possible. Metallic solids occur when large numbers of atoms bond together in close-packed structures. They can be modelled as ping-pong balls glued together as shown in the diagram.

Secondary Bonds

Secondary or weak bonds are formed when there is effectively a partial and/or momentary charge. They are secondary in terms of strength but not necessarily in terms of importance, as life is only made possible because of them.

Electric Dipoles

 To understand these forces the idea of dipoles is needed. An electric dipole is basically a pair of equal positive and negative charges separated by a small distance. These dipoles will arise, for example, in a molecule, where atoms share an electron, but the electron spends more time with one atom, because it is bigger, and less time with the smaller atom. The effective charge separation may be only a fraction (perhaps a tenth) of the electron's charge. Bonding between dipoles happens when the positive end of one dipole is attracted to the negative end of another. Since the effective charges are small compared with those involved in primary bonding, secondary bonding is about one tenth of primary bond strengths.

Hydrogen Bonding

Hydrogen bonding is the most common type of bonding between permanent dipoles.

The situation that leads to Hydrogen bonding arises with a normal bond between a Hydrogen atom and a neighbour. Since any other atom will bind the electron from the Hydrogen atom more tightly, the electron will spend more time with the other atom. This creates a permament dipole (a partially exposed proton) that can interact with other dipoles nearby.

 The water molecule (H2O) is the classic situation, where the Oxygen molecule binds electrons from both Hydrogen atoms more tightly than Hydrogen can. Hydrogen bonding between water molecules makes ice less dense than water so, in winter, rivers freeze down from the top and not up from the bottom, enabling life to survive. Hydrogen bonding is also one of the major forces responsible for the attraction between chains in polymers.

Van der Waals Bonds

The dipoles involved in Van der Waals bonding come from fluctuations in the symmetry of the electron distribution surrounding the nucleus of an atom. Momentary electric dipoles are set up and give rise to weak, very short-range, non-directional attractive forces between molecules or atoms. Permanent dipoles can also be involved, e.g. by inducing other temporary dipoles.

Summarising:

The Periodic Table gives an indication of how electrons will be shared.
Ionic bonds tend to form between elements at the edges.
Covalent bonds tend to form between elements in the middle.
Metallic bonds have overlapping potentials releasing some electrons to form a "glue".
Secondary bonds are the result of electric dipole interaction.
Hydrogen bonds form with permament dipoles.
Van der Waals bonds generally form from fluctuating dipoles.

 Ionic Covalent Metallic Van der Waals Hydrogen strong strong strong weak weak n=1, m=9 n=6, m=12 n=6, m=12 n=2 non-directional directional non-directional non-directional non-directional long range short range long range short range short range charge stacking requirements no charge stacking requirements no charge stacking requirements dipole orientation dipole orientation

Ionic crystals tend to be brittle because of the charge restriction.
Covalent crystals tend to be brittle because of the directional restriction.
Metals tend to be plastic because there are no directional or charge restrictions.

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