ENG2000 Chapter 2 Atoms and Bonding

1. ENG2000 Chapter 2Atoms and Bonding

2. Overview • Atomic structure • fundamentals • electrons and atoms • Atomic bonding • bonding mechanisms and forces • bonding types • molecule • Atomic bonding is determined by the electronic configurations of the atoms • Atomic bonding determines all the fundamental physical and electronic, magnetic, optical etc properties

3. Atoms • For our purposes, atoms are made from three fundamental particles • proton (charge = +q, m = 1.67 x 10-27kg) • neutron (charge - 0, m = 1.67 x 10-27kg) • electron (charge = -q, m = 9.1 X 10-31kg) • q = 1.6 x 10-19C • An element is defined by its atomic number, Z • Z = number of protons in the atomic nucleus • 1 (H) ≤ Z ≤ 92 (U) for naturally occurring elements • The atomic mass (A) is the sum of proton and neutron masses in the nucleus • # neutrons (N) can vary to give different isotopes of the same element

4. Masses • The atomic weight (really a mass) is typically given in units of grams per mole (g/mol.) • 1 mole of a substance contains 6.023 x 1023 particles – Avogadro’s Number • e.g. iron has an atomic weight of A = 55.85 g/mol. • Where several isotopes of a substance are present, the atomic weight is calculated from the appropriate fractions of the weights of the individual isotopes

5. nucleus orbitingelectrons Bohr Atom • In the early years of the 20th century, atomic spectroscopy indicated that electron energies are quantised • the Bohr planetary model of the atom is an early attempt to visualise a system that would yield quantised energies • it is incomplete because it does not explain why the orbiting electrons do not emit electromagnetic radiation http://www.marxists.org/reference/subject/philosophy/images/bohr.jpg http://csep10.phys.utk.edu/astr162/lect/light/bohrframe/bohr2.gif

6. 0 n=3 -1.5eV n=2 -3.4eV n=1 -13.6eV potentialenergy Energy levels • These are the first three energy levels for an isolated H atom • 1eV = 1.6 x 10-19J • the energy gained by an electron accelerated through a potential difference of 1V • To move between energy levels requires a ‘quantum jump’ • More refined measurements showed that each ‘n’ level was in fact composed of several discrete energies • Better models needed

7. 0 3d n=3 3p -1.5eV 3s 2p n=2 -3.4eV 2s n=1 -13.6eV 1s potentialenergy Other energy levels • Due to electrostatic (and other) interactions between electrons, each primary energy is in fact several closely spaced levels • These are named s, p, d, f • after the shapes of the spectroscopic lines in the early experiments • sharp, principal, diffuse, fine • Energy levels are identified by four quantum numbers

8. Wave mechanics • Numerous pieces of evidence suggest that all particles can be thought of as both particles and waves • interference effects – quintessentially wave-like phenomena – can be seen with electrons • quantum-mechanical tunnelling (see later) • called wave-particle duality • A particle’s wavelength is calculated from the de Broglie formula (1924) • where h is Planck’s constant (1901); h = 6.62 x 10-34 Js • m is the mass, v is the velocity

9. The spatial properties of the wave (x, y, t, intensity) are closely related to the probability of finding the particle at a particular location • the important part here is that the wave mechanical nature of an electron implies that we do not know the precise position • only a probability function giving the likelihood of an electron’s position Callister

10. Quantum numbers • Principal quantum numbers are n = 1, 2, 3, 4 … • they correspond to energy shells K, L, M, N, … • Second quantum number, l, is [s, p, d, f] • related to the spatial shape of the energy level • the number of sub-shells is limited to the ‘n’ for the level • A third number, ml (the magnetic quantum number), describes the number of available energy states per sub-shell • 1 for s, 3 for p, 5 for d, 7 for f • the energies of these states are identical in the absence of a magnetic field, but split when a field is applied • The last quantum number is the spin, ms • ms = ± 1/2

11. l ml ms n Planetary picture • Very, very roughly these for quantum numbers can be visualised in terms of a planetary orbit • n corresponds to the radius of the orbit • l corresponds to the shape of the orbit • ml corresponds to the tilt (or inclination) of the orbit • ms represents the two directions the ‘planet’ can spin

12. Maximum number of states * # states x 2, because two electrons (with ± spin) can exist in each state

13. Notation • The conventional notation is: n [s,p,d,f]# • where # is the number of available states that actually contain electrons • For example: • H = 1s1 • He = 1s2 • Li = 1s22s1 • Be = 1s22s2 • B = 1s22s22p1 • Ne = 1s22s22p6 • Na = 1s22s22p63s1 • Al = 1s22s22p63s23p1

14. Filling the energy levels • Electrons occupy the lowest energy state available • note that e.g. 4s < 3d, so fills first http://www.webelements.com/webelements/elements/media/e-config/H.gif http://www.chemtutor.com/scheme.gif

15. Valence electrons • The number of electrons occupying the outermost shell of an atom – the valence electrons – is important for determining the chemical properties of the atom • because these electrons will be involved with the bonding of atoms • Atoms with one electron too many (e.g. Na) or one too few (e.g. F) are highly reactive • Atoms with full shells (e.g. Ne, Ar) tend to be inert

16. Periodic table • The periodic table of the elements was originally drawn up according to the chemical properties of the elements • as we have seen these properties are closely related to the atomic electron configurations • the seven horizontal rows are called ‘periods’ • chemical properties vary from one end of the period to the other • each column – a ‘group’ – displays similar chemical properties and similar valence structures

17. http://helios.augustana.edu/physics/301/periodic-table-fix.jpghttp://helios.augustana.edu/physics/301/periodic-table-fix.jpg

18. Groups • Group 0 contains the inert (noble) gases • Group IV includes Si • important materials in Si chip manufacture are B (III) and P (V) – as we will see later • together, these materials are between metals and non-metals • Group VII are the ‘halogens’ and are one electron deficient in the valence shell • Groups IA and IIA are the alkali and alkaline earth metals • Groups IIIB – IIB are the transition metals, which have partially filled lower (d) energy states • includes ‘real’ metals and magnetic materials

19. Bonding • Atomic bonding determines many of the physical properties of a material • If two isolated atoms are brought closer together the net force varies with distance • there is a mechanism-specific attractive force (FA) • and a repulsive force (FB), which increases when the atoms are sufficiently close for the outer shells to overlap • equilibrium is reached when FA + FB = 0 • this is at r0 on the following page • The potential energy at r0 is the bonding energy, E0 • and represents the energy required to separate the atoms to an infinite distance • e.g. thermal energy to melt the material

20. Callister

21. Ionic bonding • Ionic bonding occurs in materials composed of a metallic and a non-metallic element • the metallic element easily donates its electron to the non-metallic element • the metal becomes a positive ion, while the non-metal is negatively ionised http://www.agen.ufl.edu/~chyn/age4660/lect/lect_02/2_11a.gif http://www.astro.lsa.umich.edu/users/cowley/lecture11/images/NaCl.jpg

22. Here, the attractive forces are coulombic, arising from the attraction of oppositely charged ions • E0 ≈ 600 – 1500 kJ/mol., or 3 – 8 eV/atom • this relatively large bonding energy is reflected in typically high melting temperatures for ionically bonded materials • including ceramics

23. Covalent bonding • As the name suggests, covalent bonds are formed by sharing valence electrons between the constituent atoms • thereby causing all atoms to achieve a full – and stable – outer shell • the classic example is methane, CH4 http://www.mse.cornell.edu/courses/engri111/images/covalent.gif

24. Covalent bonds are also common in elements from the right-hand side of the periodic table • notably the semiconductors silicon and germanium, as well as carbon • also compound semiconductors, e.g. GaAs and InP • The number of atoms participating in the bond is determined by the number of valence electrons • Si is in group IV, so has 4 valence electrons, and therefore bonds with 4 neighbouring atoms

25. +ve ion cores -ve electron sea Metallic bonding • Metallic elements have one or two (possibly three) ‘loose’ valence electrons • which are relatively freely donated by all atoms • The result is a structure in which ionised atoms (because they have donated their electron) are ‘suspended’ in a ‘sea’ of electrons • the ions are fixed in place because the negatively charged electron sea exerts an equal attraction in all directions

26. Metallic bonding • Because the donated electrons are freely mobile, the electrical conductivity of metals is high • heat can also be transmitted by electrons, so metals are good thermal conductors • Ionically and covalently materials are typically good electrical insulators • there is another mechanism for thermal transport which means that e.g. ceramics can be good thermal conductors http://207.10.97.102/chemzone/lessons/03bonding/mleebonding/metallicblue.gif

27. + – + – Other bonding types • Ionic, covalent and metallic are the primary bonding types • Secondary bonds are those that exist between all atoms, but are relatively weak and may be obscured by the primary bonds • van der Waals bonds are typically only 0.1eV/atom (c.f. 8eV/atom for ionic) • and results from atomic or molecular dipoles • Dipoles can result from molecular bonds – especially those involving H atoms – atomic vibrations, or external electric fields

28. Melting temperatures

29. Summary • Atomic structure is determined by quantum mechanics • four quantum numbers determine energy states • states may or may not be occupied by electrons • Atomic structure determines chemical and physical properties of the elements • periodic table • Structure also determines how atoms bond • primary – ionic, covalent, metallic • secondary – van der Waals, hydrogen