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Interatomic Bonding

Interatomic Bonding. Bonding Forces and Energies Equilibrium atomic spacing Minimization of bonding energy Embedded Atom Method (EAM) Types of Bonding Ionic Covalent Secondary Metallic. Bonding Forces and Energy. Interatomic Forces attractive forces (F a ) repulsive forces (F r )

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Interatomic Bonding

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  1. Interatomic Bonding • Bonding Forces and Energies • Equilibrium atomic spacing • Minimization of bonding energy • Embedded Atom Method (EAM) • Types of Bonding • Ionic • Covalent • Secondary • Metallic

  2. Bonding Forces and Energy • Interatomic Forces • attractive forces (Fa) • repulsive forces (Fr) • When the atoms reach a critical distance (r0), the attractive and repulsive forces cancel each other and the atoms are at their equilibrium distance.

  3. Bonding Forces and Energy

  4. Bonding Forces and Energy • Sometimes it is easier to deal with potential energies (E) rather than forces. The relation of Energy to Force is as follows: • Equilibrium is reached by minimizing EN

  5. Bonding Forces and Energy

  6. Embedded Atom Method • Potentials also calculated through the embedded atom method (EAM) • potentials are calculated as a sum of pairwise (interactions between a pair of atoms) contributions and a many body term.

  7. Embedded Atom Method • If a ternary system is being studied, EAM potentials may be defined by considering the three individual binary systems that make up the ternary system. • As long as the interatomic interaction used for each of the pure components is the same in the description of the two binaries. • The volume term is calculated as the embedding energy of a local electron density.

  8. Embedded Atom Method • Effective pairs • equivalent potentials where the various contributions (pair and volume) are not the same but add up to the same total energy for all possible simulations. • Called the effective pair scheme, it is defined as when the first derivative of the embedding function is taken as zero.

  9. Embedded Atom Method • Potentials converted to Effective pair scheme: • Transformation where mixed potentials are originally derived:

  10. EAM Potentials • Some examples of EAM functions for various metals • Ag:

  11. EAM Potentials • Al: Au:

  12. EAM Potentials • Veff for various pure elements:

  13. Ionic Bonding • Most common bonding in metal-nonmetal compounds. • Atoms give up/receive electrons from other atoms in the compound to form stable electron configurations • Because of net electrical charge in each ion, they attract each other and bond via coulombic forces.

  14. Ionic Bonding • Attractive and repulsive energies are functions of interatomic distance and may be represented as follows: • A and B are constants depending upon the system. The value of n is usually taken as 12.

  15. Ionic Bonding • Properties of ionic bonding • nondirectional: magnitude of bond is equal in all directions around the ion. • High bonding energies (~600 - 1500 kJ/mol) • reflected in high melting temperatures • generally hard and brittle materials • most common bonding for ceramic materials • electrically and thermally insulative materials

  16. Covalent Bonding • Stable configurations are obtained by the sharing of valence electrons by 2 or more atoms. • Typical in nonmetallic compounds (CH4, H20) • Number of possible bonds per atom is determined by the number of valence electrons in the following formula: • number of bonds = 8 - (valence electrons) • Bonds also are angle dependent

  17. Covalent Bonding • Properties of covalent bonding • can be either very strong or very weak bonds, depending upon the atoms involved in the bond. This is also reflected in the melting temperature of the compound • ex: diamond (strong bond) -- Tm> 3350°C bismuth (weak bond) -- Tm ~ 270°C • most common form of bonding in polymers

  18. Secondary Bonding • Van der Waals bonding • weak bonds in comparison with other forms of bonding (~10 kJ/mol) • evident between all atoms, including inert gases and especially between covalently bonded molecules. • Bonds are created through both atomic and molecular dipoles

  19. Secondary Bonding • Hydrogen bonding • special type of secondary bond between molecules with permenant dipoles and hydrogen in the compound. • Ex: HF, H2O, NH3 • these secondary bonds can have strengths as high as ~50 kJ/mol and will cause increases in melting temperature above those normally expected.

  20. Metallic Bonding • Most common in bonding of metals and their alloys. • Proposed model of metallic bonding • metals usually have, at most, 3 valence electrons, all of which form an “electron sea”, which drift through the entire metal. • Base electrons form net-positive ion cores, which attract the free electrons from the “sea” as needed to maintain neutrality.

  21. Metallic Bonding • Bonding may be weak or strong, depending upon atoms involved. • Ex: Hg bonding energy = 68 kJ/mol W bonding energy = 850 kJ/mol

  22. Metallic Bonding • Potentials for metallic bonding are most commonly calculated via the EAM, especially in alloys and intermetallics • Link to Paper by Dr. Farkas

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