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Chapter 2: Atomic Structure and Interatomic Bonding • • • • Atomic Structure Electron Configuration Periodic Table Primary Bonding – Ionic – Covalent – Metallic • Secondary Bonding or van der Waals Bonding – Three types of Dipole Bonding • Molecules Chapter 2- Atomic Models Chapter 2- ~ 400 BC - Democritus • Ancient Greek philosopher • Democritus coined the term átomos which means "uncuttable" or "the smallest indivisible particle of matter". Structure of Matter Physical world “VOID + BEING” Chapter 2- 1803 – John Dalton • English instructor and natural philosopher • “Each element consists of atoms of single unique type and can join to form chemical compounds.” • Originator of the modern atomic theory Chapter 2- 1869 - Mendeleev • Building upon earlier discoveries by scientists, Mendeleev published the first functional PERIODIC TABLE. • Certain chemical properties of elements repeat periodically when arranged by atomic number. Chapter 2- Periodic Table Draft of the first periodic table, Mendeleev, 1869 Chapter 2- 1869… Chapter 2- Today: Periodic Table of the Elements Chapter 2- The Structure of the Atom Status report end of the 19th century • Atom is electrically neutral • Negative charge carried by electrons • Electron has very small mass – bulk of the atom is positive, – most mass resides in positive charge Chapter 2- The Structure of the Atom particle symbol charge (C) mass (kg) electron e– –1.6×10–19 9.11×10–31 proton p+ +1.6×10–19 1.673×10–27 neutron no 0 1.675×10–27 Question: what is the distribution of charge inside an atom? Chapter 2- 1897 – Sir J. J. Thomson • Discovered the electron (1906 Nobel Prize in Physics). • Plum Pudding (1904): “The atom as being made up of electrons swarming in a sea of positive charge. Chapter 2- 1909 – E. Rutherford • Tested the Plum Pudding Model. • Results: – Majority of a particles transmitted (pass through) or deflected through small angles – Tiny fraction deflected through large angles Chapter 2- 1909 – E. Rutherford • Conclusion: – Disproved the Plum-Pudding Model – Large amount of the atom's charge and mass is concentrated into a small region – Atom was mostly empty space • Objections to Rutherford model – The laws of classical mechanics predict that the electron will release electromagnetic radiation while orbiting a nucleus. Because the electron would lose energy, it would gradually spiral inwards, collapsing into the nucleus. – This atom model is unsuccessful, because it predicts that all atoms are unstable. Chapter 2- 1912 – N. Bohr • Many phenomena involving electrons in solids could not be explained in terms of CLASSICAL MECHANICS. • We need QUANTUM MECHANICS… Chapter 2- Bohr Postulates for the Hydrogen Atom 1. 2. 3. 4. 5. Rutherford atom is correct Classical EM theory not applicable to orbiting eNewtonian mechanics applicable to orbiting eEelectron = Ekinetic + Epotential e- energy quantized through its angular momentum: L = mvr = nh/2π, n = 1, 2, 3,… 6. Planck-Einstein relation applies to e-transitions: ΔE = Ef-Ei= hν = hc/λ c = νλ Chapter 2- BOHR ATOM orbital electrons: n = principal quantum number 1 2 n=3 Adapted from Fig. 2.1, Callister 6e. Nucleus: Z = # protons N = # neutrons Atomic mass A ≈ Z + N Chapter 2- 2 1853 - A. Ångström Chapter 2- 1913 - Sommerfeld • German theoretical physicist • Modified the Bohr Model • “suppose we have plurality of orbits” – a shell containing multiple orbits: ORBITALS • How to capture these new ideas quantitatively? • We need new quantum numbers: n, l, m, s n principal quantum number, distance of an electron from the nucleus l subshell, describes the shape of the subshell m number of energy states in a subshell s spin moment Chapter 2- Wave mechanics to arrive at same place: E=E(n,l,m,s) • The Bohr model – significant limitations • Resolution: Wave-mechanical model (electron is considered to exhibit both wave-like and particlelike characteristics). – De Broglie: “If a photon which has no mass, can behave as a particle, does an electron which has mass can behave as a wave (1920)?” λ = h/p = h/mv – Heisenberg: Uncertainty Principle “I don’t know where any of one of electrons is, but I can tell you an average where any of one of them is likely to be” – Schrodinger Chapter 2- Beyond Bohr’s Model In 1924 de Broglie : dual character of electrons In 1927 Heisenberg : uncertainity, it is not possible to measure simultaneously both the momentum (or velocity) and the position of a microscopic particle with absolute accuracy. Schrodinger, math expression for the behavior of an electron around an atom Chapter 2- FUZZY ORBITS Chapter 2- What is the filling sequence of electrons in orbitals by n, l, m, s is not adequate? AUFBAU PRINCIPLE 3 principles: 1. Pauli Exclusion Principle:only one electron can have a given set of four quantum numbers. 2. Electrons -have discrete energy states -fill orbitals from lowest en. to highest en. 3. Hund’s rule Chapter 2- Niels Bohr, Werner Heisenberg, and Wolfgang Pauli talking in the Niels Bohr Institute lunchroom, possibly 1934 or 1936 Chapter 2- Quantum Numbers (II) l ml ms = ±½ Chapter 2- Quantum Numbers (III) Electrons fill quantum levels in order of increasing energy ( only n and l make significant differences in energy configurations). 1s, 2s, 2p, 3s,3p,4s,3d,4p,5s,4d,5p,6s,4f,5d,…. When all electrons are at the lowest possible energy levels => ground state Excited states do exist such as in glow discharges etc… Valence electrons occupy the outermost filled shell. Valence electrons are responsible for all bonding ! Chapter 2- SURVEY OF ELEMENTS • Most elements: Electron configuration not stable. Electron configuration 1s1 1s2 (stable) 1s22s1 1s22s2 Adapted from Table 2.2, 1s22s22p1 Callister 7e. 1s22s22p2 ... 1s22s22p6 (stable) 1s22s22p63s1 1s22s22p63s2 1s22s22p63s23p1 ... 1s22s22p63s23p6 (stable) ... 1s22s22p63s23p63d10 4s246 (stable) • Why? Valence (outer) shell usually not filled completely. Chapter 2- 5 STABLE ELECTRON CONFIGURATIONS Stable electron configurations... • have complete s and p subshells • tend to be unreactive. Adapted from Table 2.2, Callister 6e. Chapter 2- 4 Electron Configurations • Valence electrons – those in unfilled shells • Filled shells more stable • Valence electrons are most available for bonding and tend to control the chemical properties – example: C (atomic number = 6) 1s2 2s2 2p2 valence electrons Chapter 2- THE PERIODIC TABLE • Columns: Similar Valence Structure, Similar Properties Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions. Chapter 2- 6 ELECTRONEGATIVITY • Ranges from 0.7 to 4.0, • Large values: tendency to acquire electrons; reactivity Metals like to give up, halogens like to acquire electrons ! Smaller electronegativity Larger electronegativity Chapter 2- 7 REVIEW OF ATOMIC STRUCTURE (FRESHMAN CHEMISTRY) ATOMS = (PROTONS+NEUTRONS) + ELECTRONS NUCLEUS BONDING • Mass of an atom: – Proton and Neutron: ~ 1.67 x 10-27 kg – Electron: 9.11 x 10-31 kg • Charge: – Electrons and protons: (±) 1.60 x 10-19 C – Neutrons are neutral The atomic mass (A): total mass of protons + total mass of neutrons Atomic weight ~ Atomic mass # of protons are used to identify elements (Z) # of neutron are used to identify isotopes ( e.g. 14C6 and 12C6 ) Isotopes are written as follows: AXZ , i.e. 1H1, 2H1, 3H1 Chapter 2- Atomic Structure Valence electrons determine all of the following properties: 1) 2) 3) 4) Chemical Electrical Thermal Optical Chapter 2- Atomic bonding in solids Things are made of atoms—little particles that move around, attracting each other when they are a little distance apart, but repelling upon being squeezed into one another. In that one sentence ... there is an enormous amount of information about the world. — Richard P. Feynman Chapter 2- Atomic Bonding in Solids r • Start with two atoms infinitely separated • Attractive component is due to nature of the bonding (minimize energy thru electronic configuration) • Repulsive component is due to Pauli exclusion principle; electron clouds tend to overlap • Essentially atoms either want to give up (transfer) or acquire (share) electrons to complete electron configurations; minimize their energy – Transfer of electrons => ionic bond – Sharing of electrons => covalent – Metallic bond => sea of electons Chapter 2- IONIC BONDING (I) • • • • Occurs between + and – ions (anion and cation). Requires electron transfer. Large difference in electronegativity required. Example: Na+ Cl- Chapter 2- 8 Ionic bond – metal + donates electrons nonmetal accepts electrons Dissimilar electronegativities ex: MgO Mg 1s2 2s2 2p6 3s2 [Ne] 3s2 Mg2+ 1s2 2s2 2p6 [Ne] O 1s2 2s2 2p4 O2- 1s2 2s2 2p6 [Ne] Chapter 2- IONIC BONDING (II) Oppositely charged ions attract, attractive force is coulombic. Ionic bond is non-directional, ions get attracted to one another in any direction. Bonding energies are high => 2 to 5 eV/atom,molecule,ion Hard materials, brittle, high melting temperature, electrically and thermally insulating Chapter 2- 8 Ionic Bonding • Energy – minimum energy most stable – Energy balance of attractive and repulsive terms EN = EA + ER = - A r - B rn Repulsive energy ER Interatomic separation r Net energy EN Adapted from Fig. 2.8(b), Callister 7e. Attractive energy EA Chapter 2- Examples: Ionic Bonding • Predominant bonding in Ceramics NaCl MgO CaF 2 CsCl Give up electrons Acquire electrons Adapted from Fig. 2.7, Callister 7e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University. Chapter 2- COVALENT BONDING (I) • Requires shared electrons • Example: CH4 C: has 4 valence e, needs 4 more H: has 1 valence e, needs 1 more Electronegativities are comparable. Adapted from Fig. 2.10, Callister 6e. Chapter 2- 10 COVALENT BONDING (II) Diamond, sp3 Covalent bonds are formed by sharing of the valence electrons Covalent bonds are very directional Covalent bond model: an atom can have at most 8-N’ covalent bonds, where N’ = number of valence electrons Covalent bonds can be very strong, eg diamond, SiC, Si, etc, also can be very weak, eg Bismuth Polymeric materials do exhibit covalent type bonding. Chapter 2- 10 Primary Bonding • Metallic Bond -- delocalized as electron cloud • Ionic-Covalent Mixed Bonding % ionic character = (X A -X B )2 4 1e x (100%) where XA & XB are Pauling electronegativities Ex: MgO XMg = 1.3 XO = 3.5 (3.5 -1.3)2 4 % ionic character 1 - e x (100%) 70.2% ionic Chapter 2- EXAMPLES: COVALENT BONDING H2 H 2.1 Li 1.0 Na 0.9 K 0.8 Be 1.5 Mg 1.2 Ca 1.0 Rb 0.8 Cs 0.7 Sr 1.0 Fr 0.7 Ra 0.9 • • • • Ba 0.9 column IVA H2O C(diamond) SiC Ti 1.5 Cr 1.6 Fe 1.8 Ni 1.8 Zn 1.8 Ga 1.6 C 2.5 Si 1.8 Ge 1.8 F2 He O 2.0 As 2.0 Sn 1.8 Pb 1.8 F 4.0 Cl 3.0 Ne - Br 2.8 Ar Kr - I 2.5 Xe - At 2.2 Rn - Cl2 GaAs Adapted from Fig. 2.7, Callister 6e. (Fig. 2.7 is adapted from Linus Pauling, The Nature of the Chemical Bond, 3rd edition, Copyright 1939 and 1940, 3rd edition. Copyright 1960 by Cornell University. Molecules with nonmetals Molecules with metals and nonmetals Elemental solids (RHS of Periodic Table) Compound solids (about column IVA) Chapter 2- 11 METALLIC BONDING • Arises from a sea of donated valence electrons (1, 2, or 3 from each atom). Ion cores in the “sea of electrons”. Valance electrons belong no one particular atom but drift throughout the entire metal. “Free electrons” shield +’ly charged ions from repelling each other… Adapted from Fig. 2.11, Callister 6e. • Primary bond for metals and their alloys Chapter 2- 12 SECONDARY BONDING Arises from interaction between dipoles • Fluctuating dipoles asymmetric electron clouds + - + secondary bonding - ex: liquid H 2 H2 H2 H H H H secondary bonding Adapted from Fig. 2.13, Callister 7e. • Permanent dipoles-molecule induced -general case: -ex: liquid HCl -ex: polymer + - H Cl secondary bonding + secondary bonding H Cl Adapted from Fig. 2.14, Callister 7e. secondary bonding Chapter 2- Bonding Energies Chapter 2- Summary: Bonding Comments Type Bond Energy Ionic Large! Nondirectional (ceramics) Covalent Variable large-Diamond small-Bismuth Directional (semiconductors, ceramics polymer chains) Metallic Variable large-Tungsten small-Mercury Nondirectional (metals) Secondary smallest Directional inter-chain (polymer) inter-molecular Chapter 2- Properties From Bonding: Tm • Bond length, r • Melting Temperature, Tm Energy r • Bond energy, Eo ro Energy r smaller Tm unstretched length ro r Eo = “bond energy” larger Tm Tm is larger if Eo is larger. Chapter 2- PROPERTIES FROM BONDING: E • Elastic modulus, E Elastic modulus F L =E Ao Lo • E ~ curvature at ro Energy unstretched length ro r E is larger if Eo is larger. smaller Elastic Modulus larger Elastic Modulus Chapter 2- 16 Summary: Primary Bonds Ceramics (Ionic & covalent bonding): Metals (Metallic bonding): Polymers (Covalent & Secondary): Large bond energy large Tm large E small a Variable bond energy moderate Tm moderate E moderate a Directional Properties Secondary bonding dominates small Tm small E large a Chapter 2-