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1 Chapter 2 Atomic Structure & Bonding in Solids.

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Presentation on theme: "1 Chapter 2 Atomic Structure & Bonding in Solids."— Presentation transcript:

1 1 Chapter 2 Atomic Structure & Bonding in Solids

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

3 3 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 : x 10 –28 g Charge : x 10 –9 C Accounts for all volume Proton Mass : x 10 –24 g Charge : x 10 –19 C Neutron Mass : x 10 –24 g Neutral Charge atomic structure

4 4 Fundamental Concept atomic structure Periodic table of the elements O Se Te PoAt I Br He Ne Ar Kr Xe Rn F ClS LiBe H NaMg BaCs RaFr CaKSc SrRbY Atomic number, Z Atomic weight, A Cu

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

6 6 Electron Principle atomic structure Bohr atomic modelWave mechanical atom model electron structure & position ▣ 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 7 Electron Principle atomic structure Bohr atomic modelWave mechanical atom model electron energy ▣ Energy are quantized. ▣ Electrons are permitted to have only specific values of energy (have energy 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. Both model assume:

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

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

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

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

12 12 Electron Principle atomic structure Bohr atomic model electron configurations Wave mechanical atom model ◈ Electron Configuration: lists the arrangement of electrons in orbital. ◈ Maximum number of electrons in each atomic shell is given by 2n 2. n# of electrons ◈ 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 13 Electron Principle atomic structure electron configurations Wave mechanical atom model ▣ The # of available electron states in some of the electron shells & subshells 1 K 2 L 3 M Principle quantum number, n shell/energylevel shell designation Subsidiary quantum #, l Magnetic quantum #, m l subshells/ orbital s s p s p d s p d f N Number ofstates Number of electrons per subshellper shell

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

15 15 Electron Principle atomic structure electron configurations Wave mechanical atom model 1s 2s2p 3s3p 3d 4s4p 4d4f 5s5p 5d5f Method of arrangement:

16 16 Electron Principle atomic structure electron configurations Wave mechanical atom model

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

18 18 Electron Principle atomic structure 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 19 Periodic Table atomic structure Electropositive & electronegative elements O Se Te PoAt I Br He Ne Ar Kr Xe Rn F ClS LiBe H NaMg BaCs RaFr CaKSc SrRbY Electropositive elements Electronegative elements Inert gases

20 20 Periodic Table atomic structure electronegative elements  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 gaining valence electrons to form charge ions.  filled valence shell.  stable electron configuration.  unreactive chemically. electropositive elements Inert gases

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

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

23 23 Primary bonding atomic 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 24 atomic bonding Metal (electropositive element) - unstable Non metal (electronegative atom) - unstable electron transfer Cation (+ve charge) - stable Anion (-ve charge) - stable Ionic bond electrostatic attraction donates electrons accepts electrons Primary bonding Ionic bonding

25 25 atomic bonding Primary bonding Ionic bonding 1s 2 2s 2 2p 6 3s 2 Mg O 1s 2 2s 2 2p 4 [Ne] 3s 2 Mg: X = 1.2, Z = 12 O: X = 3.5, Z = 8 1s 2 2s 2 2p 6 Mg 2+ O 2- 1s 2 2s 2 2p 6 [Ne] [Ne] Example: Magnesium oxide (MgO) Metal (electropositive element) - unstable Non metal (electronegative atom) - unstable electron transfer Cation (+ve charge) - stable Anion (-ve charge) - stable Ionic bond MgO electrostatic attraction

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

27 27 atomic bonding Primary bonding Ionic bonding Bonding Energy ▣ Energy – minimum energy most stable ▷ Energy balance of attractive & repulsive terms Attractive energy E A Net energy E N Repulsive energy E R Interatomic separation r r A n r B E N = E A + E R = __

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

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

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

31 31 Primary bonding atomic bonding Covalent bonding Example: Methane (CH 4 ) C: X = 2.5, Z = 6H: X = 2.1, Z = 1 shared electrons from carbon atom shared electrons from hydrogen atoms H H H H C C: 1s 2 2s 2 2p 2 H: 1s 1 4 valence electrons 1 valence electron

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

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

34 34 Primary bonding atomic 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-directional bonding ▷ outer electrons are shared by many atoms.

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

36 36 Secondary bonding atomic bonding Fluctuating dipoles bond ▣ Arises from interaction between dipoles. ▣ Very weak electric dipole bonds due to asymmetric distribution of electron densities. asymmetric electron clouds +-+- secondary bonding HHHH H2H2 H2H2 secondary bonding example: liquid H 2 ▣ general case:

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

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

39 39 Summary bonding atomic bonding Primary bondingSecondary bonding

40 40 Properties from bonding atomic bonding Melting temperature  T m is larger if bonding energy, E o is larger. r o r Energy larger T m smaller T m

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

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

43 43 End of Chapter 2


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