Today's Agenda ISSUES TO ADDRESS... • What promotes bonding? • What types of bonds are there? • What properties are inferred from bonding? 1

GECKO WHY ?

Here is the answer

Chapter 2: Atomic Structure and Interatomic Bonding Electron Configuration Periodic Table Primary Bonding Ionic Covalent Metallic Secondary Bonding or van der Waals Bonding Three types of Dipole Bonding Molecules

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 Actually neutron is a little heavier than proton but no problem with that… Electrons are so much lighter so neglect in calculating the weight of an atom… Hydrogen, deuterium, tritium How is Carbon 14 technique used to measure the age of old artifacts ?

AMU and Mole Concept Nav = 1 gram/1 amu. The atomic mass unit (amu) : 1 amu is defined as the 1/12 of the atomic mass of the most common isotope of carbon, carbon 12 (12C6). Atomic mass of 12C6 is 12 amu : Carbon= 6 protons (Z=6) + 6 neutrons (N=6) Mproton ~ Mneutron = 1.67 x 10-24 g = 1 amu Atomic Mass = Z + N The atomic weight is often expressed in mass per mole 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 number, Nav = 6.023 × 1023. Nav = 1 gram/1 amu. Example: Atomic weight of iron = 55.85 amu/atom = 55.85 g/mol Inverse of mass of neutron + protons = Avagadro’s number

MOLE CONCEPT ANY IDEAS WHY ?

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

Atomic Models Towards the end of 19th century physicists realized Newtonian physics has serious difficulty in explaining many phenomena involving electrons => quantum mechanics Bohr atomic model Electrons assume very well defined orbits around the nucleus (protons + neutrons) Electrons in each shell orbit assumes the same energy level Severe issues when considering events involving electrons such as emission spectra and photoelectrons…) See figure 2.1 in Callister page 18 (page 13 for 6th ed.) Wave mechanical model Electrons in an atom or molecule are permitted to have only specific values of energy, energy is quantized… Electrons do not move in circular orbits but in “fuzzy” orbits. At any given time we can only talk about the probability of finding an electron at a radius from the orbit. Every electron is characterized by four quantum numbers. The size, shape, spatial orientation of an electrons probability density are specified by these numbers BOHR model represents an early attempt to describe electrons in atoms, interms of both position (electron orbital) and energy (quantized energy levels)

BOHR ATOM Nucleus: Z = # protons = 1 for hydrogen to 94 for plutonium Adapted from Fig. 2.1, Callister 6e. Nucleus: Z = # protons = 1 for hydrogen to 94 for plutonium N = # neutrons Atomic mass A ≈ Z + N 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

FUZZY ORBITS

ELECTRON ENERGY STATES Electrons... • have discrete energy states • tend to occupy lowest available energy state. Adapted from Fig. 2.5, Callister 6e. 3

A product of Schrodinger’s Equation Quantum Numbers (I) A product of Schrodinger’s Equation n, l , ml , ms n principal quantum number, distance of an electron from the nucleus l subshell, describes the shape of the subshell ml number of energy states in a subshell ms spin moment Pauli’s exclusion principle: only one electron can have a given set of four quantum numbers.

Quantum Numbers (II) l ml ms = ±½

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 !

SURVEY OF ELEMENTS • Most elements: Electron configuration not stable. Adapted from Table 2.2, Callister 7e. • Why? Valence (outer) shell usually not filled completely. 5

STABLE ELECTRON CONFIGURATIONS • have complete s and p subshells • tend to be unreactive. Adapted from Table 2.2, Callister 6e. 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

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. 6

Periodic Table Draft of the first periodic table, Mendeleev, 1869

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 7

Concept Checks Question: Why are the atomic weights of the elements generally not integers? Cite two reasons. Answer: The atomic weights of the elements ordinarily are not integers because: (1) The atomic masses of the atoms normally are not integers (except for 12C), and (2) the atomic weight is taken as the weighted average of the atomic masses of an atom's naturally occurring isotopes. Question: Give electron configurations for the Fe3+and S2- ions. Answer: The Fe3+ ion is an iron atom that has lost three electrons. Since the electron configuration of the Fe atom is 1s22s22p63s23p63d64s2 (Table 2.2), the configuration for Fe3+ is 1s22s22p63s23p63d5. The S2- ion a sulfur atom that has gained two electrons. Since the electron configuration of the S atom is 1s22s22p63s23p4 (Table 2.2), the configuration for S2- is 1s22s22p63s23p6.

Atomic Bonding in Solids 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 r

IONIC BONDING (I) • Occurs between + and – ions (anion and cation). • Requires electron transfer. • Large difference in electronegativity required. • Example: Na+ Cl- Easiest to describe…. 8

Ionic bond – metal + nonmetal 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]

IONIC BONDING (II) Easiest to describe…. 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 8

Ionic Bonding Energy – minimum energy most stable r A B EN = EA + ER = Energy balance of attractive and repulsive terms r A n B EN = EA + ER = - Attractive energy EA Net energy EN Repulsive energy ER Interatomic separation r 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 (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. In covalent bonding, electrons are shared between the molecules, to saturate the valency. The simplest example is the H2 molecule, where the electrons spend more time in between the nuclei than outside, thus producing bonding. Adapted from Fig. 2.10, Callister 6e. 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. In covalent bonding, electrons are shared between the molecules, to saturate the valency. The simplest example is the H2 molecule, where the electrons spend more time in between the nuclei than outside, thus producing bonding. 10

Primary Bonding Metallic Bond -- delocalized as electron cloud Ionic-Covalent Mixed Bonding % ionic character =   where XA & XB are Pauling electronegativities %) 100 ( x Ex: MgO XMg = 1.3 XO = 3.5

COVALENT BONDING (III) Very few materials have pure ionic or covalent bonding; electronegativity inpart defines how much time electrons spend between ion cores… In covalent bonding, electrons are shared between the molecules, to saturate the valency. The simplest example is the H2 molecule, where the electrons spend more time in between the nuclei than outside, thus producing bonding. 10

EXAMPLES: COVALENT BONDING 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) 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 12

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

Bonding Energies

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: E • Elastic modulus, E • E ~ curvature at ro E is larger if Eo is larger. 16

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

ANNOUNCEMENTS Read Chapter 3!!!