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Chapter 2 Atomic Structure and Interatomic Bonding

Chapter 2 Atomic Structure and Interatomic Bonding. ATOMIC STRUCTURE Fundamental Concepts Electrons in Atoms Atomic Models Quantum Numbers Electron Configurations The Periodic Table. ATOMIC BONDING IN SOLIDS Bonding Forces and Energies Primary Interatomic Bonds Ionic Bonding

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Chapter 2 Atomic Structure and Interatomic Bonding

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  1. Chapter 2Atomic Structure and Interatomic Bonding • ATOMIC STRUCTURE • Fundamental Concepts • Electrons in Atoms • Atomic Models • Quantum Numbers • Electron Configurations • The Periodic Table

  2. ATOMIC BONDING IN SOLIDS • Bonding Forces and Energies • Primary Interatomic Bonds • Ionic Bonding • Covalent Bonding • Metallic Bonding • Secondary Bonding or VAN DER WAALS BONDING • Fluctuating Induced Dipole Bonds • Polar Molecule-Induced Dipole Bonds • Permanent Dipole Bonds • Molecules

  3. ATOMIC STRUCTURE • Fundamental Concepts • atom , nucleus , protons , neutrons, ,electrons • atomic number (Z)=number of protons 1- 92 • atomic mass (A) = sum of the masses of protons and neutrons within the nucleus • number of protons is the same for all atoms of a given element, the number of neutrons (N) may be variable • Isotopes = atoms of some elements have two or more different atomic masses • atomic mass unit (amu) = 1.66 x 10-24 gm = 12.0 gm/mole (c12)/ 0.602x 1024 (Avogadro Number) atoms/mole • atomic weight = weighted average of the atomic masses of the atom’s naturally occurring isotopes

  4. Individual atoms and ions ATOMIC STRUCTURE • Atoms = nucleus (protons and neutrons) + electrons • Charges:Electrons and protons have negative and • positive charges of the same magnitude, 1.6 × 10-19 • Coulombs. Neutrons are electrically neutral. • Masses:Protons and Neutrons have the same mass, 1.67 • × 10-27 kg. Mass of an electron is much smaller, 9.11 × 10-31 • kg and can be neglected in calculation of atomic mass. • # protons gives chemical identification of the element • # protons = atomic number (Z) • # neutrons defines isotope number • The atomic mass (A) = mass of protons + mass of • neutrons

  5. ATOMIC STRUCTURE • The atomic mass unit (amu) is often used to express • atomic weight. 1 amu is defined as 1/12 of the atomic • mass of the most common isotope of carbon atom • that has 6 protons (Z=6) and six neutrons (N=6). Mproton ≈ Mneutron = 1.66 x 10-24 g = 1 amu. The atomic mass of the 12C atom is 12 amu. • The atomic mass of the 13C atom is 13 amu • The atomic weight of an element = weighted average • ofthe atomic masses of the atoms naturally occurring • isotopes. Atomic weight of carbon is 12.011 amu. • A= Z + N • The atomic weight is often specified in mass/mole.

  6. ATOMIC STRUCTURE • A mole is the amount of matter that has a mass in • grams equal to the atomic mass in amu of the atoms • (A mole of carbon has a mass of 12 grams). • The number of atoms in a mole is called the • Avogadro’s number, Nav = 6.023 × 1023. Nav = 1 gram/1 amu. • Example: Atomic weight of iron = 55.85 amu/atom = 55.85 g/mol

  7. ATOMIC STRUCTURE • The number of atoms per cm3, n, for a material of • density  (g/cm3) and atomic mass M (g/mol): n = Nav x  / M • Graphite (carbon):  = 2.3 g/cm3, M = 12 g/mol n = 6×1023 atoms/mol × 2.3 g/cm3/(12 g/mol) = 11.5 x 1022 atoms/cm3 • Diamond (carbon):  = 3.5 g/cm3, M = 12 g/mol n = 6×1023 atoms/mol x 3.5 g/cm3/12 g/mol = 17.5 x1022 atoms/cm3 • Water (H2O) d = 1 g/cm3, M = 18 g/mol n = 6×1023 molecules/mol x 1 g/cm3/18 g/mol = 3.3 x 1022 molecules/cm3

  8. ATOMIC STRUCTURE • For material with n = 6 × 1022 atoms/cm3 we can • calculate the mean distance between atoms L = (1/n)1/3 • = 0.25 nm. • the scale of atomic structures in solids – a fraction of 1 • nm or a few A.

  9. ELECTRONS IN ATOMS • Bohr atomic model • electrons are assumed to revolve around the atomic nucleus in discrete orbitals, and the position of any particular electron is more or less well defined in terms of its orbital

  10. wave-mechanical model • electron is considered to exhibit both wave-like and particle-like characteristics • position is considered to be the probability of an electron’s being at various locations around the nucleus

  11. Quantum Numbers • principal quantum number n 1, 2, 3, 4, 5, . . . , K, L, M, N, O, (shell) • quantum number is related to the distance of an electron from the nucleus, or its position • second quantum number, l = s, p, d, or f (subshell) • third quantum number ml number of energy states per subshell • fourth quantum number msspin moment (1/2 or -1/2)

  12. Quantum Numbers

  13. Quantum Numbers • Energy states • 1s state is less than that of a 2s state • 3d state is greater than a 3p, which is larger than 3s

  14. Electron Configurations • Pauli exclusion principle : each electron state can hold no more than two electrons, which must have opposite spins.

  15. Quantum Numbers • ground state: all the electrons occupy the lowest possible energies • electron configuration or structure: represents the manner in which these states are occupied • the electron configurations for hydrogen, helium, and sodium are, respectively, 1s1, 1s2, and 1s22s22p63s1

  16. Sodium Atom 1s12s22p63s1 Schematic representation of the filled and lowest unfilled energy states for a sodium atom.

  17. valence electrons: occupy the outermost shell • stable electron configurations: • states within the outermost or valence electron shell are completely filled • occupation of just the s and p states for the outermost shell by a total of eight electrons • Example: neon, argon, and krypton (inert, or noble, gases) Neon Ne 1s22s22p6

  18. Answer the following ( 1 bounus for first 2 right answers) • Give electron configurations for the Fe3+ and S2_ ions

  19. Elements in the same column (Elemental Group) share • similar properties. Group number indicates the number • of electrons available for bonding. • 0: Inert gases (He, Ne, Ar...) have filled subshells: chem. • Inactive • IA: Alkali metals (Li, Na, K…) have one electron in • outermost occupied s subshell - eager to give up • electron – chem. Active • VIIA: Halogens (F, Br, Cl...) missing one electron in • outermost occupied p shell - want to gain electron - • chem. active

  20. Electronegativity - a measure of how willing atoms • are to accept electrons. • Subshells with one electron – low electronegativity. • Subshells with one missing electron–high • electronegativity. • Electronegativity increases from left to right • Metals are electropositive – they can give up their few • valence electrons to become positively charged ions

  21. NigelsTableFull_edited.pdf

  22. Atomic Bonding in SolidsBONDING FORCES AND ENERGIES • FN = FA + FR • FA = Attractive Force • FR = repulsive Force • FA + FR = 0 Balance = dEN/dr • ro = equilibrium spacing=0.3 nm for most atoms

  23. Atomic Bonding in SolidsBONDING FORCES AND ENERGIES • material properties depend on E0, the curve shape, and bonding type. • materials having large bonding energies typically also have high melting temperatures • mechanical stiffness (or modulus of elasticity) of a material is dependent on the shape of its force- versus interatomic separation curve • slope for a relatively stiff material at the r= r0 position on the curve will be quite steep; slopes are shallower for more flexible materials • deep and narrow “trough low coefficient of thermal expansion

  24. PRIMARY INTERATOMIC BONDSIonic Bonding • value of n is approximately 8 • A = • Where is the permittivity of a vacuum (8.85 x 10-12 F/m) , Z1 and Z2 are the valences of the two ion types, and e is the electronic charge

  25. PRIMARY INTERATOMIC BONDSIonic Bonding

  26. PRIMARY INTERATOMIC BONDSCovalent Bonding • Two atoms that are covalently bonded will each contribute at least one electron to the bond, and the shared electrons may be considered to belong to both atoms. • Many nonmetallic elemental molecules(H2, Cl2, F2, etc.) as well as molecules containing dissimilar atomsCH4, H2O, HNO3 are covalently bonded Schematic representation of covalent bonding in a molecule of methane (CH4)

  27. PRIMARY INTERATOMIC BONDSCovalent Bonding • The number of covalent bonds that is possible for a particular atom is determined by the number of valence electrons. For N’ valence electrons, an atom can covalently bond with at most 8-N’ other atoms. • N’=7 for chlorine, and 8-N’=1 which means that one Cl atom can bond to only one other atom, as in Cl2 . • Similarly, for carbon,N’=4, and each carbon atom has 8-4 or four, electrons to share. Diamond is simply the three-dimensional interconnecting structure wherein each carbon atom covalently bonds with four other carbon atoms

  28. PRIMARY INTERATOMIC BONDSCovalent Bonding • It is possible to have interatomic bonds that are partially ionic and partially covalent, • For a compound, the degree of either bond type depends on the relative positions of the constituent atoms in the periodic table or the difference in their electronegativities. • The wider the separation (both horizontally—relative to Group IVA—and vertically) from the lower left to the upper-right-hand corner (i.e., the greater the difference in electronegativity), the more ionic the bond. • Conversely, the closer the atoms are together (i.e., the smaller the difference in electronegativity), the greater the degree of covalency. • The percentage ionic character of a bond between elements A and B (A being the most electronegative) may be approximated by the expression Where XAand XBarethe electronegativities for the respective elements

  29. PRIMARY INTERATOMIC BONDSMetallic bonding • This bond comes into existence if valence electrons are shared between number of atoms, i.e. arranged positive nucleuses are surrounded by electron pool. • Shared electrons are not specific to a pair of atoms, in contrast to Covalent bond, i.e. electrons are delocalized. • As shared electrons are delocalized, metallic bonds are non- directional. • Very characteristic properties of metals like high thermal and electrical conductivities are result of presence of delocalized electron pool.

  30. Solve • The mass of a small diamond is 3.1 mg. (a) How many C13 atoms are present if carbon contains 1.1 % of that isotope? (b) what is the weight percent of that isotope? • NAV= 6.022x1023 Atoms/mole

  31. P 2.7

  32. P2.11

  33. P 2.14

  34. P 2.18

  35. P 2.21

  36. The mass of a small diamond is 3.1 mg. (a) How many C13 atoms are present if carbon contains 1.1 % of that isotope? (b) what is the weight percent of that isotope? Procedure: (a) Determine the total number of carbon atoms present from the atomic masses; then 1.1 percent (on the atomic basis), (b) From the numbers of atoms in part (a) and the atomic masses, calculate the mass ( or weight) fraction. (a)# atoms = .0031[g]* 0.6022x1024 [atoms/mol]/12.011[g/mol] = 1.55x1020 C atoms C13 atoms = 1.55x1020 x .011 = 1.7x 1018 (b) Basis: 3.1 mg Since the mass of the C13 is greater than that of the average atom, the w/o will be greater than the a/o

  37. . 10 g of nickel are to be electroplated on a steel surface with an area of 0.8953 m2. The electrolyte contains Ni 2+ ions. (a) How thick will the nickel plate be? (b) what amperage is required of the plating is to be accomplished in 50 minutes? Procedure: (a) Set up your own equation based on ρ = m/v = m/At. (b) Determine the number of nickel atoms. Two electrons are required per Ni2+,each with 0.16 × 10-18 A.s ( or 0.16 × 10-18 C). The (amp)(sec) produced is obtained from these.

  38. (a) (b)

  39. Thanks

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