Atomic Structure and Bonding Chapter 2 Atomic Structure and Bonding What holds atoms together anyway?

Atomic Bonding Bonds are in equilibrium when forces are equal o zero. His occurs at some equilirbiurm distance, ro, or as r infinity.

Atomic Structure Valence electrons determine all of the following properties Chemical Electrical Thermal Optical

Bohr Atomic Model (Niels Bohr, 1915 - 1962) Nucleus Bohr Atomic Model (Niels Bohr, 1915 - 1962) Electrons have fixed energy and are resticed to specific “shells” Bohr Atomic Model: very early Electrons aranged in orbitals around nucleas. Orbitals define positions of electrons General QM: each electron has an energy level/state assoc. with orbital. “What do we know about electrons?” Heisenberg uncertainty - you can never know where an electron is at a given instant. dual wave/particle nature Must use notion of Electron clouds. Bohr model insufficient Electron shells

Electronic Structure Electrons have wavelike and particulate properties. This means that electrons are in orbitals defined by a probability. Each orbital at discrete energy level determined by quantum numbers.   Quantum # Designation n = principal (energy level-shell) K, L, M, N, O (1, 2, 3, etc.) l = subsidiary (orbitals) s, p, d, f (0, 1, 2, 3,…, n -1) ml = magnetic 1, 3, 5, 7 (-l to +l) ms = spin ½, -½

Electron Clouds (Orbitals) Heisenberg – the Orbital is just a state of energy, a level of energy but even with the energy state known, the position of the electron is only known as a probability of its positions. s,p,d and f orbitals give probability densities of positions in increasingly higher order energy states

Wave-mechanical Model and Quantum Numbers Principal Quantum Number, n - 1,2,3,4,... => K.L.M.N, ... => Bohr Shells Second Quantum Number, l - Subshells => s,p,d, and f Third and Fourth - spin moment and orientation Points of interest: Farther from the nucleus, the energy level goes up. Change in shell or subshell position requires energy loss/absorption Move out - must have absorbed energy Move in - gives off energy Electrons on outer shell are called- valence electrons v.e. take part in forming bonds with other atoms Atoms like to have their outer shells filled. Shell may loose, gain, or share electrons to complete shell. Types of Chemical Bonds in Between Atoms

Electron Energy States Electrons... • have discrete energy states • tend to occupy lowest available energy state. 1s 2s 2p K-shell n = 1 L-shell n = 2 3s 3p M-shell n = 3 3d 4s 4p 4d Energy N-shell n = 4 Adapted from Fig. 2.4, Callister 7e.

SURVEY OF ELEMENTS • Most elements: Electron configuration not stable. ... 1s 2 2s 2p 6 3s 3p 3d 10 4s 4p Atomic # 18 36 Element 1 Hydrogen Helium 3 Lithium 4 Beryllium 5 Boron Carbon Neon 11 Sodium 12 Magnesium 13 Aluminum Argon Krypton Adapted from Table 2.2, Callister 7e. • Why? Valence (outer) shell usually not filled completely.

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

Electronic Configurations 26 1s2 2s2 2p6 3s2 3p6 3d 6 4s2 valence electrons ex: Fe - atomic # = 1s 2s 2p K-shell n = 1 L-shell n = 2 3s 3p M-shell n = 3 3d 4s 4p 4d Energy N-shell n = 4 Adapted from Fig. 2.4, Callister 7e.

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The Periodic Table • Columns: Similar Valence Structure give up 1e inert gases accept 1e accept 2e O Se Te Po At I Br He Ne Ar Kr Xe Rn F Cl S Li Be H Na Mg Ba Cs Ra Fr Ca K Sc Sr Rb Y Adapted from Fig. 2.6, Callister 7e. Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions.

The Periodic Table • Columns: Similar Valence Structure give up 1e inert gases accept 1e accept 2e O Se Te Po At I Br He Ne Ar Kr Xe Rn F Cl S Li Be H Na Mg Ba Cs Ra Fr Ca K Sc Sr Rb Y Adapted from Fig. 2.6, Callister 7e. Electropositive elements: Readily give up electrons to become + ions. Electronegative elements: Readily acquire electrons to become - ions.

Types of Chemical Bonds in Between Atoms Primary Ionic Covalent Metallic Secondary Van der Waals Forces

Ionic Bonding - + • Occurs between + and - ions. • Requires electron transfer. • Large difference in electronegativity required. • Example: NaCl Na (metal) unstable Cl (nonmetal) electron + - Coulombic Attraction Na (cation) stable Cl (anion) Example: NaCl Na's outer shell has one extra electron, which it gives up. Cl's outer shell has a one empty location, that accepts from Na. Na  Na+ and Cl  Cl- The attractive forces between the two form the ionic bond.

Atomic Bonding Bonds are in equilibrium when forces are equal o zero. His occurs at some equilirbiurm distance, ro, or as r infinity.

Ionic Bonding r A B EN = EA + ER = Attractive and repulsive terms exist. The _______ energy state is most stable r A n B EN = EA + ER = - Attractive energy EA Net energy EN Repulsive energy ER Interatomic separation r Force versus seperation distance curve Interatomic Seperation Str depends w/ type of bond Str. Vary with distance Net = attract + repel FN=FA+FR Can also define bond energy. A is found from the valences of atom and from physical constants. Adapted from Fig. 2.8(b), Callister 7e.

Examples: Ionic Bonding • Predominant bonding in Ceramics NaCl MgO Give up electrons Acquire electrons CaF 2 CsCl 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.

Covalent Bonding similar electronegativity  share electrons bonds determined by valence – s & p orbitals dominate bonding Example: CH4 shared electrons from carbon atom from hydrogen atoms H C CH 4 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 7e.

Metallic Bonds “Sea of Electrons”

SECONDARY BONDING + - Arises from interaction between dipoles • Fluctuating dipoles asymmetric electron clouds + - secondary bonding H 2 ex: liquid H Adapted from Fig. 2.13, Callister 7e. • Permanent dipoles-molecule induced + - -general case: secondary bonding Adapted from Fig. 2.14, Callister 7e. Cl Cl -ex: liquid HCl secondary H H bonding secondary bonding -ex: polymer secondary bonding

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

Properties From Bonding: Tm • Bond length, r • Melting Temperature, Tm r o Energy r • Bond energy, Eo Eo = “bond energy” Energy r o unstretched length smaller Tm larger Tm Tm is larger if Eo is larger.

Properties From Bonding : a • Coefficient of thermal expansion, a coeff. thermal expansion D L length, o unheated, T 1 heated, T 2 D L = a ( T - T ) 2 1 L o • a ~ symmetry at ro r o larger a smaller a Energy unstretched length Eo a is larger if Eo is smaller.

Summary: Primary Bonds Ceramics Large bond energy large Tm large E small a (Ionic & covalent bonding): Metals Variable bond energy moderate Tm moderate E moderate a (Metallic bonding): Polymers Directional Properties Secondary bonding dominates small Tm small E large a (Covalent & Secondary): secondary bonding