Bonding forces and energies Primary interatomic bonds Secondary bonding Molecules

Transcription

1 Chapter 2. Atomic structure and interatomic bonding 2.1. Atomic structure Fundamental concepts Electrons in atoms The periodic table 2.2. Atomic bonding in solids Bonding forces and energies Primary interatomic bonds Secondary bonding Molecules

2 2.1. Atomic structure Fundamental concepts Atom consists of a nucleus (protons and neutrons) and electrons that move in their orbits. e : 1.6 X C M p = m n = 1.67 X kg M e = 9.11 X kg A chemical element has Atomic number (Z) : the number of proton in the nucleus, Atomic mass (A) : the sum of the masses of protons and neutrons in the nucleus, Neutron number (N) : the number of neutron may vary for a given elements. Isotopes : elements with two or more different atomic masses Atomic weight (amu) A Z + N

3 Electrons in atoms Bohr atomic model An electron is an particle Electrons revolve around an atomic nucleus in discrete orbitals

4 Wave-mechanical model Electrons are considered as both wave-like and particle-like Every electron in an atom is characterized by four parameters called quantum numbers. Principal quantum number, n n = 1 K; n = 2 L; and so on Second quantum number, l s, p, d, f Third quantum number, m l (number of energy states) s m l = 1; p m l = 3 ; d m l = 5 ; f m l = 7 Fourth quantum number, m s (spin orientation when a magnetic field is applied) m s = + ½ and ½

5 The schematic of the relative energies of the electrons for various shells and subshells

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7 Comparison between the two models a. Bohr atomic model Electrons are in fixed positions and energy (quantized energy levels). b. Wave-mechanical model Electrons position is considered to be the probability of an electron s being at various locations around the nucleus.

8 Electron configurations Pauli exclusion principle Each electron state can hold no more than two electrons, which must have opposite spins Ground state is a level where electrons occupy the lowest possible energy according to Pauli exclusion principle Electron configurations represent the manner in which these states are occupied. Examples: Hydrogen (H) 1s 1 Helium (He) 1s 2 Sodium (Na) 1s 2 2s 2 2p 6 3s 1 Calcium (Ca) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2

9 The periodic table The periodic table consists of all elements that have been classified according to electron configuration. Groups IA and IIA (alkali and alkaline earth metals) Elements with one and two electrons in excess of stable structures. Groups IIIA, IVA, and VA Elements between metals and nonmetals by virtue of their valence electron structures. Groups VIA and VIIA Elements with two and one electron in deficient of stable structures. (Group VIIA is called the halogens) Group VIIIA (Inert gases) Elements which have filled electron shells and stable configurations.

10 2.2. Atomic bonding in solids Bonding forces and energies Interatomic forces determine the physical properties of materials. There are two types of forces: the attractive and the repulsive forces Most of time, it is more convenient to work with potential energies, which are attractive and the repulsive energies. Both terms depend on the distance between the centre of two atoms. The equilibrium distance, r o For many atoms, r o = 0.3 nm (3Å )

11 Bonding force The net force, F N F N = F A + F R In equilibrium: F A + F R = 0

12 Bonding energy The potential energy between two atoms E = Fdr The net energy, E N E N = E A + E R

13 Types of bonding: A. Primary bonding or chemical bonding This bonding is found in solids and involves the valence electrons. This type of bonding is strong (» 100 kj/mol) Examples: ionic, covalent, and metallic bonds B. Secondary bonding or physical bonding or van der Waals This bonding is found in most solids and arises from atomic or molecular dipoles. This type of bonding is weak ( 10kJ/mol) Examples: fluctuating induced dipole bonds, polar molecule- Induced dipole bonds, and pemanent dipole bonds

14 A. Primary bonding or chemical bonding Ionic bonding It is always found in compounds that are composed of both metallic and nonmetallic elements. Atoms of a metallic element easily give up their valence electrons to the nonmetallic atoms. This bonding is a nondirectional bonding, the magnitude of the bond is equal in all directions around an ion. Coulombic bonding force Attractive energy: Repulsive energy: EA E B = = A, B, n = constants, n ~ 8 B r A r n

15 A. Primary bonding or chemical bonding Covalent bonding It is usually found in many nonmetallic elemental molecules (H 2, Cl 2, F 2 ) and molecules containing dissimilar atoms (CH 4, H 2 0, HNO 3, HF) This bonding is formed on stable electron configurations by sharing of electrons between adjacent atoms. A very strong covalent bond Diamond with a very high melting temperature (713 kj/mol; 3550 ºC) A very weak covalent bond Bismuth with a very low melting temperature (270 ºC)

16 A. Primary bonding or chemical bonding Metallic bonding It is found in many metals and their alloys (group IA and IIA). Metallic materials have 1, 2 or at most 3 valence electrons. These valence electrons are not bound to any particular atom to any Particular atom in the solid and are free to drift throughout the entire metal. sea of electrons or electron cloud Net negative charge Ion cores Net positive charge Weak metallic bond Hg (68 kj/mol; -39 ºC) Strong metallic bond W (850 kj/mol; 3410 ºC)

17 B. Secondary bonding or physical bonding or van der Waals Fluctuating induced dipole bonds All atoms have constant vibrational motion and it causes electrical symmetry and creates small electric dipoles

18 B. Secondary bonding or physical bonding or van der Waals Polar molecule-induced dipole bonds It causes by virtue of an asymmetrical arrangement of positively and negatively charged regions

19 B. Secondary bonding or physical bonding or van der Waals Permanent dipole bonds It exist between adjacent polar molecules. The hydrogen bond is the strongest secondary bonding type.