ENS 205 Materials Science I Chapter 2: Atomic Bonding

1. ENS 205Materials Science IChapter 2: Atomic Bonding http://hyperphysics.phy-astr.gsu.edu/hbase/hframe.html

2. Objectives At the end of this chapter: • Know the quantum number of elements and apply them. • Know the periodic table of elements • Know the 4 methods by which atoms bond to each other • Understand the energy/force relationship between atoms making atomic bonds.

3. Material Infrastructure • What makes their materials behavior, mechanical for instance, different? • Microstructure- major properties result from mechanisms occurring at either atomic or the microscopic level • Chemical or Atomic Bonding • Strong bonding of ceramics: high strength and stiffness, and resistance to temperature and corrosion, but brittle • Weakly bonding of chain molecules in polymers: low strength and stiffness, creep deformation

4. Atom • Atoms = nucleus (protons and neutrons) + electrons • Protons and Neutrons have the same mass, and determines the weight of the atom • Mass of an electron is much smaller than mass of proton/neutron, and can be neglected in calculation of atomic mass.

5. Atom

6. Atom: Definitions • Consider the number of protons and neutrons in the nucleus as the basis of the chemical identification  periodic table (placed by the number of protons) • Atomic Mass Unit (amu) =mass of proton or neutron ~ 1.66x10-24 gr

7. Atom: Definitions • Atomic number = the number of protons in the nucleus • Avogadro’s number, Nav : 6.023x1023 # of protons or neutrons necessary to produce a mass of 1 gr. Avogadro’s number (Nav )of atoms of a given element termed as gram-atom amu=1/ Nav 1.66x10-24 = 1/6.023x1023

8. Atom: definitions • 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). Example: C12 carbon isotope 1 C12 atom 6 protons+6 neutrons 12 amu Nav many C12 atom1 mole C12 atom 12 gr • Mole of a compound contains Avogadro’s number of each constituent atom • E.g. 1 mole of NaCl, 6.023x1023 of Na atoms + 6.023x1023 Cl atoms

9. Atomic number Atomic mass (in amu)

10. Quantum Numbers Electronic energy levels in atoms are specified by using quantum numbers The principal quantum is “n”. • n indicates the primary electron shell in an atom where the shells are represented by K=1, L=2, M=3, etc.

11. Planetary atomic model The atomic structure of sodium, atomic number 11, showing the electrons in the K, L, and M quantum shells. the most inner K-shell can accommodate only two electrons, called s-electrons; the next L-shell two s-electrons and six p-electrons; the M-shell can host two s-electrons, six p-electrons, and ten d-electrons; and so on. The electronic configuration of the different energy levels fill in a relatively straight forward pattern in a shorthand notation. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 5s2 … . eg., for Carbon, which has an atomic number of 6, it has 6 protons and 6 electrons. It’s electronic configuration in shorthand notation is 1s2 2s2 2p2 .

13. Electron (Atomic) Orbitals

14. Electron (Atomic) Orbitals • The electron volt (eV) – energy unit convenient for description of atomic bonding • Electron volt - the energy lost / gained by an electron after it has moved through a potential difference of 1 volt . E = q × V • For q = 1.6 x 10-19 Coulombs V = 1 volt 1 eV = 1.6 x 10-19 J

15. Identification of the Elements We can identify the elements using their florescence energy when a material is irradiated by an x-ray, electron or gamma ray.

16. Identification of the Elements The energy of an x-ray emitted from a K, L or M shell electron can be used to identify the atomic number of the element present in a material.

17. Atomic Bonding • Classification of engineering materials may be based on the nature of atomic bonding. Understanding the atomic bonding requires the understanding of the structure of the individual atoms • Chemical bonds: hold atoms and molecules together in solids. • Most of the materials not composed of just a single specie of atoms. They are compounds, composed of molecules made up of atoms from two or more elements. • When two or more atoms combine to form molecules of a compound, they form atomic bonds between them through chemical bonding. • Chemical bonding is essentially the interaction of electrons from one atom with the electrons of another atom. The bonding of adjacent atoms is essentially an electronic process • Primary Bonding • Secondary Bonding

18. Atomic Bonding • When atoms are combined into solids, there are several bonding mechanisms that can occur, which result in properties that may differ substantially from those of the atom alone. Hence, it is necessary to understand the types of bonding that can occur In the Solid Sphere Model, there are three primary or strong bonds and one weaker or secondary (but important!) type of bond between atoms or ions. • 1) Ionic bonds • 2) Covalent bonds • 3) Metallic bonds • 4) Van der Waals bonds

19. Atomic or ionic radius • An atomic or ionic radius refers to the radius corresponding to the average electron density

20. Valence Electrons • Valence electrons are those electrons in the outer shells that are easily removed or added to form either a positive or negative charge for the purpose of combinations with other atoms. • These then form ions, which we shall see, are important for ceramics and semiconductors. • Valence electrons are the single most important structure of an atom or ion as they determine the physical (mechanical), electrical, photonic and magnetic properties of materials.

21. Valence Electrons What is the valence of an atom? • The valence is the ability of the atom to enter into chemical combination with other elements and is often determined by the number of outermost combined s, p, and /or d levels. • Examples are: • Mg: 1s2 2s2 2p63s2 valence = 2 • Al: 1s2 2s2 2p63s2 3p1valence = 3 • Ge: 1s2 2s2 2p6 3s2 3p1 3d104s2 4p2 valence = 4 Valence electrons determine all of the properties of the material! http://hyperphysics.phy-astr.gsu.edu/hbase/hframe.html

22. Valence Electrons(cont’d) There are exceptions to the filling order of the electronic shells • e.g., Iron, Fe – atomic no. = 26; 1s2 2s2 2p6 3s2 3p63d8 [3d64s2 ]; instead of completely filling the 3d orbital with 8 electrons, Fe first fills the 4s orbital. Electron Configuration of Nickel http://www.webelements.com/webelements/elements/text/Fe/econ.html

23. Exceptions in 3d, 4d, 5d • A d subshell that is half-filled or full (ie 5 or 10 electrons) is more stable than the s subshell of the next shell. This is the case because it takes less energy to maintain an electron in a half-filled d subshell than a filled s subshell. • For instance, copper (atomic number 29) has a configuration of [Ar]4s1 3d10, not [Ar]4s2 3d9 • Likewise, chromium (atomic number 24) has a configuration of [Ar]4s1 3d5, not [Ar]4s2 3d4 where [Ar] represents the configuration for argon.

24. Valence Electrons

25. Atomic Structure • Filled outermost shells are the most stable (non-reactive) configurations. The atoms with unfilled valence shells strive to reach the stable configuration by gaining or loosing electrons or sharing electrons with other atoms. This transference/sharing of electrons result in a strong bonding among atoms,

26. Electronegativity Some properties of elements include: • Electronegativity is the tendency of an atom to gain an electron. High electronegativity atoms tend to be on the right side of the Periodic Table and low electronegativity atoms are on the left side. What is the most electronegative element? • Electropositivity is the tendency of an atom to loss electrons. • High electronegative atoms tend to react with high electropositive atoms to form ionic molecules and ceramic materials. • The sharing of electrons tends to make very strong atomic bonds. In the case of ceramics these bonds may break abruptly making the ceramic brittle.

27. The electronegativities of selected elements relative to the position of the elements in the periodic table.

28. http://hyperphysics.phy-astr.gsu.edu/hbase/hframe.html

29. Periodic Table The atomic number, atomic mass, density and crystal structure are given.

30. Atomic Bonding • Primary Bonds are formed when outer orbital electrons are transferred or shared between atoms. strong and stiff, hard to melt, metals and ceramics, • Ionic • Covalent • Metalic http://hyperphysics.phy-astr.gsu.edu/hbase/Chemical/eleorb.html

31. Secondary bonds • Secondary bonds: relatively weak, behavior of liquids, bonds between carbon-chain molecules in polymers, due to subtle attraction between positive and negative charges (no transfer or sharing) • Van der Waals • Hydrogen

32. Primary Chemical Bonds: Ionic Bonding • An ionic bond is created between two unlike atoms with different electronegativities. When sodium donates its valence electron to chlorine, each becomes an ion; attraction occurs due to their opposite electrostatic charges, and the ionic bond is formed. • The size of the Cl ion is big compared to its elemental size whereas the size of Na ion is small compared to its elemental size. • eg. Na and Cl form NaCl where the properties of the resultant material (salt) is very different from either of the atoms. Cl and Na are both highly corrosive where Cl is associated with acids and Na is associated with bases.

33. Primary Chemical Bonds: Ionic Bonding • A collection of such charged ions, form and electrically neutral solid by arranging themselves into regular crystalline array • Makes material hard and brittle • Non-directional: A cation (Na+) will attract any adjacent anion (Cl-) equally in all directions

34. When a voltage is applied to an ionic material, entire ions must move to cause a current to flow. Ion movement is slow and the electrical conductivity is poor. Thus ionic materials like SiO2 and Al2O3 make good insulators of electricity.

35. Primary Chemical Bonds: Ionic Bonding Nature of the bonding force for the ionic bond  coulombic attractions force Fc With small a, Fcgets large, then a ideallybe equal to zero?

36. Primary Chemical Bonds Ionic Bonding • With small a, FCgets large, then a ideallybe equal to zero? • Oppositely charged ions gets closer, leads to increase in FC, but it is counteracted by an opposing repulsive force FR due to • overlapping of the similarly charged electric fields from each ions • the attempt to bring the two positively charged nuclei closer together • where λ and ρ are experimentally determined constants for a given ion pair

37. Primary Chemical Bonds: Ionic Bonding Interatomic spacing The equilibrium distance between atoms is caused by a balance between repulsive and attractive forces. Equilibrium separation occurs where the total-atomic energy of the pair of atoms is at a minimum, or when no net force is acting to either attract or repel the atoms.The interatomic spacing is approximately equal to the atomic diameter or, for ionic materials, the sum of the two different ionic radii. Bonding Force, Net force F=FR+FC Equilibrium bond length where F=0

38. Primary Chemical Bonds: Ionic Bonding • Bonding energy, E is related to bonding force through the differential expression Equilibrium bond length a0 corresponds to • F = 0 and • A minimum in the energy curve  stable ions positions

39. Primary Chemical Bonds: Ionic Bonding A material that has a high binding energy will also have a high strength and high melting temperature.

40. Bonding Energy • How does bonding energy relate to melting point? • Modulus of Elasticity? • Coefficient of Thermal Expansion? • Hint: The higher the bonding energy the more tightly the atoms are held together.

41. Primary Chemical Bonds: Ionic BondingCoordination number • Coordination number is the number of adjacent ions (or atoms) surrounding a reference ion (or atom) • Depends directly on the relative sizes of the oppositely charged ions • Radius ratio r/R (smaller ion to the larger ion)

42. Coordination number Larger ions overlap: instability because of high repulsive forces

43. Coordination number MORE TO COME IN Ch 3…

44. Coordination Number As r/R→1, a coordination number as high as 12 is possible

45. Questions to think on ? • Why don’t we have a coordination number greater than unity. • Why coordination numbers of 5, 7, 10, 11 are absent?

46. Covalent Bonds Materials with covalent bonds tend to occur among atoms with small differences in electronegativity and therefore the elements are close to one another in the periodic table. • Two or more atomsshare two or more electrons. • The atoms most commonly share their outer s and p electrons so that each atom can tend to approach an inert gas structure. • Example, Si; Z = 14; 1s2 2s2 2p63s2 3p2 or 1s2 2s2 2p63s1 3p3 are possible with the second configuration being more stable.

47. Electron orbitals are represented as particles orbiting at a fixed radius. In reality, electrons charge is found in a range of radii. Representation of the actual electron density Highly directional due to sharing of electrons with specific neighboring atoms

48. Primary Chemical Bonds: Covalent Bonding • While ionic bonds are non-directional, covalent bonds are very directional so atoms can best share their electrons. • Covalent bonds are very strong. • They tend to be brittle with poor electrical conductivity. Why then is Silicon and other like materials used in the electronics industry? • Many hydrocarbons, eg., C2 H4 , are covalently bonded. Many polymeric materials such as polyvinyl chloride (PVC), used as molded plastic on cars, have primarily covalent bonds.

49. Primary Chemical Bonds: Covalent Bonding • A continuous covalent bond arrangement to form a 3D network of a solid • Diamond is a cubic crystal structure of carbon (formed at a temperature of 1325°C, a pressure of 50000 kg/cm2 is required to grow diamond) • Highest melting temperature • Highest hardness • Highest elastic modulus

50. Carbons’s electronic configuration in shorthand notation is 1s22s2 2p2 Double bond  covalent sharing of two pairs of valence electrons When energy provided Bonding of adjacent molecules, double bondsingle bond between each adjacent molecule pair