Chapter 2: Atomic Structure and Interactive Bonding

1. Chapter 2: Atomic Structure and Interactive Bonding • Why Study Atomic Structure and Interactive Bonding? • An important reason to have an understanding of inter-atomic bonding in solids is that, in some instances, the type of bond allows us to explain a material’s properties. • For example, consider carbon, which may exist as both graphite and diamond. Whereas graphite is relatively soft and has a greasy feel to it, diamond is the hardest known material. This dramatic disparity in properties is directly attributable to a type of interatomic bonding found in graphite that does not exist in diamond.

2. 2.4 Periodic Table • All the elements have been classified according to electron configuration. • Elements are situated with increasing atomic number. • Seven horizontal rows are called periods. • All elements that are arrayed in a given column or group have similar valence electron structures, as well as chemical and physical properties.

3. 2.4 Periodic Table (Contd.)

4. 2.4 Periodic Table (Contd.) • Electropositive elements  capable of giving up their few valence electrons to become positively charged ions. • Electronegative elements  readily accept electrons to form negatively charged ions, or sometimes they share electrons with other atoms. • Figure 2.7 displays electronegativity values assigned to various elements. • As a general rule, electronegativity increases in moving from left to right and from bottom to top.

5. 2.4 Periodic Table (Contd.)

6. Ch. 2: ATOMIC BONDING IN SOLIDS • While there are only 118 or so elements listed on the periodic table, there are obviously more substances in nature than 118 pure elements.  This is because atoms can react with one another to form new substances called compounds.  When two or more atoms chemically bond together, the resulting compound is unique both chemically and physically from its parent atoms.  • Let's look at an example.  • The element sodium is a silver-colored metal that reacts so violently with water that flames are produced when sodium gets wet.  The element chlorine is a greenish-colored gas that is so poisonous that it was used as a weapon in World War I.  When chemically bonded together, these two dangerous substances form the compound sodium chloride, a compound so safe that we eat it every day - common table salt!

7. NaCl + 

8. Bonding Forces and Energies • Consider two isolated atoms: • When the atoms are at large inter-atomic separation distance, the atoms do not exert any force on each other. • When the distance is decreased, an attractive force FA starts to act pulling atoms closer. • FA increases as the atoms gets closer. • But as the atoms get closer a repulsive force FR begin to act. • The net force FN between the two atoms is given by: FN = FA + FR • At some inter-atomic distance ro, FR exactly equals FA and FN becomes Zero FN = 0 = FA + FR • ro is called the equilibrium inter-atomic separation distance at which atoms enter into bonding ro ≈ 0.3 nm

9. Force vs. Separation Distance Energy vs. Separation Distance

10. 2.5 Bonding Forces and Energies (Contd.) • Sometimes it is more convenient to work with the potential energies between two atoms instead of forces.Mathematically, energy (E) and force (F) are related as E =  F dr For atomic systems, EN, EA, and ER : the net, attractive, and repulsive energies for two isolated and adjacent atoms.

11. 2.5 Bonding Forces and Energies (Contd.) • Figure 2.8b plots potential energies. • The net curve, which is sum of attractive and repulsive energies, has a potential energy well around its minimum. • E0 Bonding energy at minimum point, energy required to separate these two atoms to an infinite separation. • A similar yet complex condition exists for solid materials because force and energy interactions among many atoms must be considered. • Nevertheless, a bonding energy, analogous to E0 above, may be associated with each atom.

12. 2.5 Bonding Forces and Energies (Contd.) • A number of material properties depend on E0, the curve shape, and bonding type. For example, • Materials having large bonding energies typically also have high melting temperature. • At room temperature, solids formed for large bonding energies, whereas for small energies the gaseous state is favored, liquids prevail when energies are of intermediate magnitude. • The mechanical stiffness (modulus of elasticity) is dependent on the shape of the force-versus-interatomic separation curve. • The coefficient of thermal expansion (how much expands or contracts per degree change in temperature) is related to the shape of its E0-versus-r0 curve.

13. Types of Atomic & Molecular Bonds • Primary Atomic Bonds • Ionic Bonds • Covalent Bonds • Metallic Bonds • Secondary Atomic & Molecular Bonds • Permanent Dipole Bonds • Fluctuating Dipole Bonds

14. Ionic Bonding • Large inter-atomic forces are created by the “coulombic” effect produced by positively and negatively charged ions. • Ionic bonds are “non-directional”. • The “cation” has a + charge & the “anion” has the - charge. • The cation is much smaller than the anion.

15. IONIC BONDING IN NaCl

16. NaCl

17. IONIC BONDING • The attractive energy EA is a function of inter-atomic separation distance. EA = -A/r Where A = (z1e)(z2e)/40 z1 and z2 are the valance of the two ions. e is the electron charge (1.6 x 10-19 C) o is the permittivity of vacuum (8.85 x 10-12 F/m) Repulsive energy for two isolated ions: ER = B / (rn) A, B, and n are constants and depends on ionic system, n is approximately 8.

18. Energy and Force relationship Energy is related to force as: E = ∫F dr and thus F = dE/dr

19. Useful Equations EA = -A/r FA = dE/dr = -A [d/dr(1/r)] FA = A/r2 Since A = (z1e)(z2e)/4o FA = (z1e)(z2e)/4or2 or r = √ (z1e)(z2e)/4oFA

22. Covalent Bonding • Large inter-atomic forces are created by the sharing of electrons to form directional bonds. • The atoms have small differences in electronegativity & close to each other in the periodic table. • The atoms share their outer s and p electrons so that each atom attains the noble-gas electron configuration.

23. H2 Molecule

24. Primary Interatomic Bonding (Contd.) • Examples of covalent bonding: Nonmetallic elemental molecules (H2, Cl2, F2, etc.) Dissimilar atoms ( CH4, H2O, HNO3, and HF ) Elemental solids ( Diamond (Carbon), silicon, germanium ) Number of covalent bond = { 8 – (No. of valence electrons) } Chlorine :8-7=1 can bond to only one other atom. Carbon: 8-4=4 can bond to four other atoms. • Covalent bonds may be very strong (diamond) or very weak (bismuth). See Table 2.3.

25. Primary Interatomic Bonding (Contd.) • It is possible to have interatomic bonds that are partially ionic and partially covalent, and, in fact, very few compounds exhibit pure ionic or covalent bonding. • For a compound, the degree of either bond type depends on the relative positions of the constituent atoms in the periodic table (Figure 2.6) or the difference in their electronegativities (Figure 2.7). • The percent ionic character of a bond between elements A and B (A being the most electronegative) may be approximated by the expression: % ionic character = {1 – exp[-(0.25)(XA – XB)2]} x 100 Where XA and XB are the electronegativities for the respective elements.

27. Metallic Bonding • Large inter-atomic forces are created by the sharing of electrons in a delocalized manner to form strong non-directional bonding.

28. Cloud of electrons surrounding the positively charged cores

30. Secondary Atomic & Molecular Bonds [Van der Waals Bonds] • Physical Bonds (no electron involvement). • Weak electrostatic bonds. • Bonding occurs between atomic or molecular electric dipoles. • Bonding Energies are low. The bondsare relatively weak.

31. Secondary Atomic & Molecular Bonds [Van der Waals Bonds] • Fluctuating Induced Dipole Bonds • Weak electric dipole bonding can take place among atoms due to an instantaneous asymmetrical distribution of electron densities around their nuclei. • This type of bonding is termed fluctuation since the electron density is continuously changing. • Example: Bonding between Argon atoms • Polar Molecule-Induced Dipole Bonds • Weak intermolecular bonds are formed between atoms or molecules which possess permanent dipoles. • A dipole exists in a molecule if there is asymmetry in its electron density distribution.

32. Fluctuating Dipole Bonds Bonding Between Argon electric dipoles - + Argon atom Argon Electric Diploe ≈ Van der waals bond

33. Polar Molecule-Induced Dipole Bonds • Permanent dipole exist in some molecules. • Such molecules are called Polar Molecules. • Polar Molecules can induce dipoles in adjacent non-polar molecules and bonding can take place. • Example: HCl molecule Cl H - +

34. Permanent Dipole Bonds • Van der waal bonds also occur between permanent polar molecules. • The bonding energies are higher than the fluctuating induced or polar molecule induced bonds. • The strongest Secondary Bonding is Hydrogen Bond. • Examples of Hydrogen Bonding: • HF, H2O, NH3