Chapter 2 Atomic Structure & Bonding in Solids

1. Chapter 2 Atomic Structure & Bonding inSolids

2. Issues to address • What promotes bonding? • What types of bonds are there? • What properties are inferred from bonding?

3. atomic structure Fundamental Concept Atom Basic Unit of an Element Diameter : 10 –10 m. Neutrally Charged Nucleus Diameter : 10 –14 m Positive Charge Accounts for almost all mass Electron Cloud Mass : 9.109 x 10 –28 g Charge : -1.602 x 10 –9 C Accounts for all volume Neutron Mass : 1.675 x 10 –24 g Neutral Charge Proton Mass : 1.673 x 10 –24 g Charge : 1.602 x 10 –19 C

4. atomic structure H He Li Be O F Ne Na Mg S Cl Ar K Ca Sc Se Br Kr Rb Sr Y Te I Xe Cs Ba Po At Rn Fr Ra Fundamental Concept Periodic table of the elements Atomic number, Z Atomic weight, A 29 Cu 63.54

5. atomic structure Electron Principle Atomic structure Principle electron nucleus quantum mechanics classical mechanics  electron structure electron position Bohr atomic model Wave mechanical atom model electron energy electron configuration

6. atomic structure Electron Principle electron structure & position Bohr atomic model Wave mechanical atom model ▣ structure: assume electrons as particle-like & revolve around the atomic nucleus in discrete paths. ▣ position: in terms of its orbital in discrete path. ▣ limitation: inability to explain several phenomena involving electron. ▣ further refined from Bohr model. ▣ structure: assume electrons as particle-like & wave-like. ▣ Position: in terms of probability distribution (various location around the atomic nucleus).

7. atomic structure Electron Principle electron energy Bohr atomic model Wave mechanical atom model Both model assume: ▣ Energy are quantized. ▣ Electrons are permitted to have only specific values of energy (have energy levels @ states). ▣ Each adjacent orbital/state are separated by finite energies. ▣ Electron may change energy by make a quantum jump. ▣ Energy is absorbed to move to higher energy level. ▣ Energy is emitted during transition to lower level.

8. atomic structure Electron Principle electron energy Bohr atomic model Wave mechanical atom model emit energy (photon) absorb energy (photon) energy levels

9. atomic structure n=1 n=2 n=3 Electron Principle electron energy Bohr atomic model Wave mechanical atom model ▣ Eachorbital at discreteenergy levels separate into electron subshells & quantum numbers dictate the number of state within each subshell. ▣ electron in its orbital (position) & separate by energy levels. ▣ Every electron characterized by 4 quantum numbers. 1. Principle quantum number, n. (energy level/shell). 2. Subsidiary quantum number, l. (orbital/subshell). 3. Magnetic quantum number, ml. (state). 4. Spin quantum number, ms. (spin)

10. atomic structure Electron Principle electron energy Bohr atomic model Wave mechanical atom model Orbital/subshell s orbital (l=0) Energy level s orbital (l=0) orbital n=1 n=2 p orbital (l=1) n=3 n=2 n=1 Energy level/shell Bohr atomic model

11. atomic structure Electron Principle electron energy Wave mechanical atom model Electron energy level 4d eV N-shell n = 4 4p Energy 3d 4s 3p M-shell n = 3 -1.5 3s 2p L-shell n = 2 -3.4 2s -13.6 K-shell n = 1 1s Energy level/shell Orbital/ subshell State

12. atomic structure Electron Principle electron configurations Bohr atomic model Wave mechanical atom model ◈ Electron Configuration: lists the arrangement of electrons in orbital. ◈ Maximum number of electrons in each atomic shell is given by 2n2. n# of electrons 1 2 2 8 3 18 4 32 ◈ Apply Pauli exclusion principle: Each electron state can hold no more than 2 electrons which must have opposite spin. ◈ Electrons have discrete energy state & tend to occupy lowest energy state. ◈ Atomic size (radius) increases with addition of shells.

13. atomic structure Electron Principle electron configurations Number of electrons Wave mechanical atom model shell designation shell/energy level ▣ The # of available electron states in some of the electron shells & subshells Number of states subshells/orbital per subshell per shell 2 1 s 1 K 2 2 s 1 2 8 L p 6 3 s 2 1 p 6 3 3 18 M 10 d 5 s 2 1 p 6 3 4 N 32 d 10 5 f 14 7 Principle quantum number, n Subsidiary quantum #, l Magnetic quantum #, ml

14. atomic structure Electron Principle electron configurations Wave mechanical atom model example: Fe (Z = 26) electron configuration is 4d n = 4 Highest energy state 4p 1s2 2s2 2p6 3s2 3p6 3d 6 4s2 3d 4s Principal Quantum Numbers Orbital letters Energy 3p n = 3 3s # of electrons 2p Lowest energy state (ground state) n = 2 2s 1s n = 1

15. atomic structure Electron Principle electron configurations Wave mechanical atom model Method of arrangement: 5d 5f 5s 5p 4d 4f 4s 4p 3d 3s 3p 2s 2p 1s

16. atomic structure Electron configuration of some elements Element Atomic # Electron configuration 1 Hydrogen 1 1s 2 Helium 2 1s (stable) Lithium 3 1s 2 2s 1 Beryllium 4 1s 2 2s 2 Boron 5 1s 2 2s 2 2p 1 Carbon 6 1s 2 2s 2 2p 2 ... ... Neon 10 1s 2 2s 2 2p 6 (stable) Sodium 11 1s 2 2s 2 2p 6 3s 1 Magnesium 12 1s 2 2s 2 2p 6 3s 2 Aluminum 13 1s 2 2s 2 2p 6 3s 2 3p 1 ... ... Argon 18 1s 2 2s 2 2p 6 3s 2 3p 6 (stable) ... ... ... Krypton 36 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 (stable) Electron Principle electron configurations Wave mechanical atom model

17. atomic structure valence electrons Electron Principle electron configurations Wave mechanical atom model ▣ Most elements: electron configuration not stable. ▣ Why?  Valence (outer) shell usually not filled completely. • example: Carbon C • 1s22s2 2p2 • atomic number, Z = 6

18. atomic structure Electron Principle Valence electron • ▣ Electron that occupy the outermost (valence) shell. • ▣Valence electrons • ▷ those in unfilled shells (most elements) not stable & filled shells (inert gases) more stable. • ▣ Participate in the bonding (unfilled shell) between atoms to form atomic & molecular aggregates. • ▷ determine physical (optical, thermal & electrical) & chemical properties.

19. atomic structure H He Li Be O F Ne Na Mg S Cl Ar K Ca Sc Se Br Kr Rb Sr Y Te I Xe Cs Ba Po At Rn Fr Ra Periodic Table Electropositive & electronegative elements Inert gases Electropositive elements Electronegative elements

20. atomic structure Periodic Table electropositive elements electronegative elements Inert gases  readily give up electrons to become cations (+ions).  metallic elements.  smaller electro-negativity.  readily accept electrons to become anions (-ions).  non-metallic elements.  higher electro-negativity.  unfilled valence shell.  not stable electron configuration.  assume stable by losing @ gaining valence electrons to form charge ions.  filled valence shell.  stable electron configuration.  unreactive chemically.

21. atomic structure Periodic Table electronegativity ▣ Ranges from 0.7 to 4.0. ▣ Large values: tendency to acquire electrons. Smaller electronegativity Larger electronegativity

22. atomic bonding Type of bonding Atomic bonding type Primary bonding Secondary bonding Fluctuating dipoles bond Permanent dipoles bond Ionic bonding Covalent bonding Metallic bonding

23. atomic bonding Primary bonding Ionic bonding ▣ Strong atomic bonds due to transfer of electrons. ▣ It can form between metallic & nonmetallic elements. ▣ Electrons are transferred from electropositive to electronegative atoms. ▣ Large difference in electronegativity required. ▣ Occurs between + & - ions. ▣Non Directional bonding.

24. atomic bonding Primary bonding Ionic bonding Metal (electropositive element) - unstable Non metal (electronegative atom) - unstable accepts electrons electron transfer donates electrons Cation (+ve charge) - stable Anion (-ve charge) - stable electrostatic attraction Ionic bond

25. atomic bonding Primary bonding Ionic bonding Example: Magnesium oxide (MgO) Mg: X = 1.2, Z = 12 O: X = 3.5, Z = 8 electron transfer Non metal (electronegative atom) - unstable Metal (electropositive element) - unstable 1s2 2s2 2p63s2MgO 1s2 2s2 2p4 [Ne] 3s2 electrostatic attraction Anion (-ve charge) - stable Cation (+ve charge) - stable 1s2 2s2 2p6Mg2+O2- 1s2 2s2 2p6 [Ne] [Ne] MgO Ionic bond

26. atomic bonding Primary bonding Ionic bonding Bonding Force ▣ due to electrostatic attraction.

27. atomic bonding Repulsive energy ER A B EN = EA+ ER = _ _ Interatomic separation r r r n Net energy EN Attractive energy EA Primary bonding Ionic bonding Bonding Energy • ▣ Energy – minimum energy most stable • ▷ Energy balance of attractive & repulsive terms

28. atomic bonding Primary bonding Ionic bonding Bonding Energy Energy r o r Eo = “bond energy”

29. atomic bonding NaCl MgO CaF 2 CsCl Primary bonding Ionic bonding ▣ Predominant bonding in Ceramics Give up electrons Acquire electrons

30. atomic bonding overlapping electron clouds Primary bonding Covalent bonding ▣ share valence electrons. ▣ similar @ small differences in electronegativity. ▣ bonds determined by valence – s & p orbitals dominate bonding. ▣ Directional bonding.

31. atomic bonding H H H C H Primary bonding Covalent bonding Example: Methane (CH4) 4 valence electrons C: X = 2.5, Z = 6 H: X = 2.1, Z = 1 C: 1s22s22p2 shared electrons H: 1s1 from carbon atom 1 valence electron shared electrons from hydrogen atoms

32. atomic bonding x ( 100 %) Primary bonding Ionic & Covalent bonding Percent ionic character ▣Ionic-Covalent Mixed Bonding ▣% ionic character = where XA = electronegativity value for element A XB = electronegativity value for element B

33. atomic bonding x ( 100 %) Primary bonding Ionic & Covalent bonding Percent ionic character Example: Magnesium oxide (MgO) XMg = 1.3 XO = 3.5  % ionic character = = 70.2 % ionic

34. atomic bonding Primary bonding Metallic bonding ▣ Atoms in metals are closely packed in crystal structure. ▣ Loosely bounded valence electrons are attracted towards nucleus of other atoms. ▣ Electrons spread out among atoms forming electron clouds. ▷ these free electrons are reason for a good electric conductivity & ductility. ▣Non-directionalbonding ▷ outer electrons are shared by many atoms.

35. atomic bonding Primary bonding Metallic bonding positive ion valence electron charge cloud

36. atomic bonding H H H H Secondary bonding Fluctuating dipoles bond ▣ Arises from interaction between dipoles. ▣ Very weak electric dipole bonds due to asymmetric distribution of electron densities. ▣ general case: example: liquid H2 asymmetric electron clouds H2 H2 + - + - secondary secondary bonding bonding

37. atomic bonding + - + - Cl Cl H H secondary bonding Secondary bonding Permanent dipoles bond ▣ Also, arises from interaction between dipoles. ▣ Weak electric dipole bonds due to molecule induced. ▣ general case: secondary bonding Example 1: liquid HCl acid secondary bonding Example 2: polymer

38. atomic bonding Summary bonding Primary bonding Secondary bonding Type Bond Energy Comments Ionic Large Nondirectional (ceramics) Directional (semiconductors, ceramics polymer chains) Covalent Variable large-Diamond small-Bismuth Metallic Variable large-Tungsten Nondirectional (metals) small-Mercury Secondary smallest Directional inter-chain (polymer) inter-molecular

39. atomic bonding Summary bonding Primary bonding Secondary bonding

40. atomic bonding Energy r o r smaller Tm larger Tm Properties from bonding Melting temperature  Tm is larger if bonding energy, Eo is larger.

41. atomic bonding length, L o unheated, T 1  L heated, T 2 r o Properties from bonding Coefficient of thermal expansion coeff. thermal expansion  L α = ( T - T ) 2 1 L o Energy unstretched length a is smaller if Eo is larger. r larger α Eo Eo smaller α

42. atomic bonding secondary bonding Properties from bonding Summary Ceramics Large bond energy large Tm large E small α (Ionic & covalent bonding): Metals Variable bond energy moderate Tm moderate E moderate α (Metallic bonding): Polymers Directional Properties Secondary bonding dominates small Tm small E large α (Covalent & Secondary):

43. End of Chapter 2