Week 2 PDF - Atomic Structure & Interatomic Bonding

Chapter 2: Atomic Structure & Interatomic Bonding ISSUES TO ADDRESS... What characteristics of atoms/molecules promote interatomic/intermolecular bonding? What types of interatomic/intermolecular bonds exist ? What properties of materials depend on the magnitude of in...

Chapter 2: Atomic Structure & Interatomic Bonding ISSUES TO ADDRESS... What characteristics of atoms/molecules promote interatomic/intermolecular bonding? What types of interatomic/intermolecular bonds exist ? What properties of materials depend on the magnitude of interatomic/intermolecular bonds ? Chapter 2 - 1 Atomic Structure (Freshman Chem.) atom – electrons – 9.11 x 10-31 kg protons neutrons } 1.67 x 10 -27 kg atomic number = # of protons in nucleus of atom = # of electrons in neutral species atomic mass unit = amu = 1/12 mass of 12C A = Atomic wt = wt of 6.022 x 1023 molecules or atoms 1 amu/atom = 1 g/mol C 12.011 H 1.008 etc. Chapter 2 - 2  Atomic Structure (cont.) Some of the following properties are determined by an atom's electronic structure: 1) Chemical 2) Electrical 3) Thermal 4) Optical Chapter 2 - 3  Electronic Structure Electrons have wave-like and particle-like characteristics. Two wave-like characteristics are – Electron position in terms of probability density – shape, size, orientation of probability density determined by quantum numbers – Quantum # Designation/Values n = principal (shell) K, L, M, N, O (1, 2, 3, 4, etc.) ℓ = azimuthal (subshell) s, p, d, f (0, 1, 2, 3,…, n-1) mℓ = magnetic (no. of orbitals) 1, 3, 5, 7 (-ℓ to +ℓ) ms = spin ½, -½ Chapter 2 - 4  Electron Energy States Electrons... have discrete energy values tend to occupy lowest available energy states 4d 4p N-shell n = 4 3d 4s Energy 3p M-shell n = 3 3s 2p L-shell n = 2 2s 1s K-shell n = 1 Chapter 2 - 5  SURVEY OF ELEMENTS Most elements: Electron configurations not stable. Element Atomic # Electron configuration Hydrogen 1 1s 1 Helium 2 1s 2 (stable) Lithium 3 1s 2 2s 1 Beryllium 4 1s 2 2s2 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) Why not stable? Valence (outer) shell usually not completely filled. Chapter 2 - 6  Electron Configurations Valence electrons – those in outer unfilled shells Filled shells are more stable – require more energy to gain or lose electrons Valence electrons available for bonding and tend to determine an atom’s chemical properties – example: C (atomic number = 6) 1s2 2s2 2p2 valence electrons Chapter 2 - 7  Electronic Configurations (cont.) ex: Fe (atomic # = 26) Electron configuration 1s2 2s2 2p6 3s2 3p6 3d 6 4s2 4d 4p N-shell n = 4 valence electrons 3d 4s Energy 3p M-shell n = 3 3s 2p L-shell n = 2 2s 1s K-shell n = 1 Chapter 2 - 8  The Periodic Table Elements in each column: Similar valence electron structure inert gases give up 1e- give up 2e- accept 2e- accept 1e- give up 3e- 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 Electropositive elements: Electronegative elements: Readily give up electrons Readily acquire electrons to become + ions. to become - ions. Chapter 2 - 9  Electronegativity Ranges from 0.7 to 4.0, Large values: tendency to acquire electrons. Smaller electronegativity Larger electronegativity Chapter 2 - 10  Ionization Process metal atom + nonmetal atom donates accepts electrons electrons Dissimilar electronegativities ex: MgO Mg 1s2 2s2 2p6 3s2 O 1s2 2s2 2p4 [Ne] 3s2 Mg2+ 1s2 2s2 2p6 O2- 1s2 2s2 2p6 [Ne] [Ne] Chapter 2 - 11  Ionic Bonding Occurs between + and - ions. Requires electron transfer. Large difference in electronegativity required. Example: NaCl Na (metal) Cl (nonmetal) unstable unstable electron Na (cation) + - Cl (anion) stable Coulombic stable Attraction Chapter 2 - 12  Ionic Bonding (cont.) Energy – minimum energy most stable – Net energy = sum of attractive and repulsive energies – Equilibrium separation when net energy is a minimum EN = EA + ER = - A + Bn r r Repulsive energy ER Interatomic separation r Net energy EN Fig. 2.10(b), Callister & Rethwisch 10e. Attractive energy EA Chapter 2 - 13  Ionic Bonding (cont.) Predominant bonding in Ceramics Examples: NaCl MgO CaF 2 CsCl Give up electrons Acquire electrons Chapter 2 - 14  Covalent Bonding Similar electronegativities ∴ share electrons Bonds involve valence electrons – normally s and p orbitals are involved Example: H2 H2 Each H: has 1 valence e-, needs 1 more H H Electronegativities are the same. shared 1s electron shared 1s electron from 1st hydrogen from 2nd hydrogen atom atom Fig. 2.12, Calliser & Rethwisch 10e. Chapter 2 - 15 Covalent Bonding: Bond Hybrization Carbon can form sp3 hybrid orbitals sp3 109.5° 3 sp C sp3 sp3 Fig. 2.14, Callister & Rethwisch 10e. (Adapted from J.E. Brady and F. Senese, Chemistry: Matter and Its Changes, 4th edition. Reprinted with permission of John Wiley and Sons, Inc.) Fig. 2.13, Callister & Rethwisch 10e. Chapter 2 - 16  Covalent Bonding (cont.) Hybrid sp3 bonding involving carbon Example: CH4 H C: each has 4 valence electrons, 1s needs 4 more sp3 sp3 H: each has 1 valence electron, needs 1 more sp3 H C 1s H 1s H sp3 Electronegativities of C and H 1s are similar so electrons are Region of overlap shared in sp3 hybrid covalent Fig. 2.15, Callister & Rethwisch 10e. bonds. (Adapted from J.E. Brady and F. Senese, Chemistry: Matter and Its Changes, 4th edition. Reprinted with permission of John Wiley and Sons, Inc.) Chapter 2 - 17  Metallic Bonding Electrons delocalized to form an “electron cloud” Fig. 2.19b, Callister & Rethwisch 10e. Chapter 2 - 18  Mixed Bonding Most common mixed bonding type is Covalent-Ionic mixed bonding " − ( X A −XB )2 % % ionic character = $$1 − e 4 ' x (100%) ' # & where XA & XB are electronegativities of the two elements participating in the bond Ex: MgO XMg = 1.2 XO = 3.5 ⎛ − (3.5−1.2)2 ⎞ ⎜ % ionic character = 1− e 4 ⎟ x (100%) = 73.3% ⎜ ⎟ ⎝ ⎠ Chapter 2 - 19  Secondary Bonding Arises from attractive forces between dipoles Fluctuating dipoles asymmetric electron ex: liquid H 2 clouds H2 H2 + - + - H H H H secondary secondary bonding bonding Permanent dipoles secondary -general case: + - bonding + - secondary -ex: liquid HCl H Cl bonding H Cl secon -ex: polymer d a ry b ondin g linear polymer molecule Chapter 2 - 20  Properties Related to Bonding I: Melting Temperature (Tm) Bond length, r Melting Temperature, Tm Energy r Bond energy, Eo ro r Energy smaller Tm unstretched length ro larger Tm r Eo = The larger Eo, the higher Tm “bond energy” Chapter 2 - 21  Properties Related to Bonding II: Coefficient of Thermal Expansion (αl) Coefficient of thermal expansion, αl length, L o ΔL = αl (T2 -T1) unheated, T1 Lo ΔL heated, T2 The smaller Eo, the larger αl. unstretched length Increase in bond length is due to Energy ro asymmetry of the E vs. r curve. This r results in an increase in αl. larger αl As E0 increases this asymmetry Eo decreases. Eo smaller αl Chapter 2 - 22  Summary: Properties Related to Bonding Type and Bonding Energy Ceramics Large bond energy (Ionic & covalent bonding): high Tm large E small αl Metals Variable bond energy (Metallic bonding): moderate Tm moderate E moderate αl Polymers Weak bond energy (between chains) (Covalent & Secondary): Secondary bonding responsible for most physical properties secon low Tm dary b ondin g small E large αl Chapter 2 - 23  SUMMARY A material’s chemical, electrical, thermal, and optical properties are determined by electronic configuration. Valence electrons occupy the outermost unfilled electron shell. Primary bonding types include covalent, ionic, and metallic bonding. Secondary or van der Waals bonds are weaker than the primary bonding types. The percent ionic character of a covalent-ionic mixed bond between two elements depends on their electronegativities. Chapter 2 - 24