![Molecular Orbital Diagrams
An MO diagram, just like an atomic orbital diagram,
shows the relative energy and number of electrons in
each MO.
The MO diagram also shows the AOs from which each
MO is formed.
Bond order is calculated as follows:
½[(# of e- in bonding MO) – (# of e- in antibonding MO)]](https://image.slidesharecdn.com/atomicstructure-230923175928-26eb3452/85/Atomic-structure-pdf-47-320.jpg)
Atomic structure.pdf
- 2. Atomic Structure All matter is composed of atoms. Understanding the structure of atoms is critical to understanding the properties of matter
- 3. Subatomic Particles Particle Mass (g) Charge (Coulombs) Charge (units) Electron (e- ) 9.1 x 10-28 -1.6 x 10-19 -1 Proton (p) 1.67 x 10-24 +1.6 x 10-19 +1 Neutron (n) 1.67 x 10-24 0 0 mass p = mass n = 1840 x mass e-
- 7. Rutherford observed that most of the alpha-particles passed through the foil without any deflection from their path and struck the screen at its centre, causing illuminations. A few of them were deflected at some angles (90 0 or wider angles) after passing through the foil. Very few (not more than one in 10,000) turned back to their original path.
- 8. Postulates of Rutherford's Model 1. Atom contains a massive (heavy) and positively charged part at its centre. This central part of the atom is called nucleus. 2. The volume occupied by the nucleus is only a minute fraction of the total volume of the atom. 3. Atom is not all solid, as was earlier suggested by Dalton, but is extraordinarily hollow, since it consists of a lot of empty space round the nucleus. 4. Electrons are revolving round the nucleus in closed orbits with a fast speed and hence almost all the space round the nucleus is occupied by the revolving electrons.
- 10. Atomic Structure Atoms are composed of -protons – positively charged particles -neutrons – neutral particles -electrons – negatively charged particles Protons and neutrons are located in the nucleus. Electrons are found in orbitals surrounding the nucleus.
- 11. Atomic Structure Every different atom has a characteristic number of protons in the nucleus. atomic number = number of protons Atoms with the same atomic number have the same chemical properties and belong to the same element.
- 12. Atomic Structure Each proton and neutron has a mass of approximately 1 dalton. The sum of protons and neutrons is the atom’s atomic mass. Isotopes – atoms of the same element that have different atomic mass numbers due to different numbers of neutrons.
- 13. 1. e- can have only specific (quantized) energy values 2. light is emitted as e- moves from one energy level to a lower energy level Bohr’s Model of the Atom (1913) En = -RH ( ) 1 n2 n (principal quantum number) = 1,2,3,… RH (Rydberg constant) = 2.18 x 10-18J
- 14. The Bohr Model of the Atom
- 15. The Bohr Model of the Atom: Ground and Excited States • In the Bohr model of hydrogen, the lowest amount of energy hydrogen’s one electron can have corresponds to being in the n = 1 orbit. We call this its ground state. • When the atom gains energy, the electron leaps to a higher energy orbit. We call this an excited state. • The atom is less stable in an excited state and so it will release the extra energy to return to the ground state. – Either all at once or in several steps.
- 17. Barbara A. Gage PGCC CHM 1010 Why Do Atoms Bond? • Chemical bonds form because they lower the potential energy between the charged particles that compose atoms • A chemical bond forms when the potential energy of the bonded atoms is less than the potential energy of the separate atoms
- 18. Barbara A. Gage PGCC CHM 1010 Types of Bonds Types of Atoms Type of Bond Bond Characteristic metals to nonmetals Ionic electrons transferred nonmetals to nonmetals Covalent electrons shared metals to metals Metallic electrons pooled • We can classify bonds based on the kinds of atoms that are bonded together
- 19. Barbara A. Gage PGCC CHM 1010 Ionic Bonds • When a metal atom loses electrons it becomes a cation – metals have low ionization energy, making it relatively easy to remove electrons from them • When a nonmetal atom gains electrons it becomes an anion – nonmetals have high electron affinities, making it advantageous to add electrons to these atoms • The oppositely charged ions are then attracted to each other, resulting in an ionic bond
- 20. Barbara A. Gage PGCC CHM 1010 Covalent Bonds • Nonmetal atoms have relatively high ionization energies, so it is difficult to remove electrons from them • When nonmetals bond together, it is better in terms of potential energy for the atoms to share valence electrons – potential energy lowest when the electrons are between the nuclei • Shared electrons hold the atoms together by attracting nuclei of both atoms
- 22. Barbara A. Gage PGCC CHM 1010 Metallic Bonds • The relatively low ionization energy of metals allows them to lose electrons easily • The simplest theory of metallic bonding involves the metal atoms releasing their valence electrons to be shared as a pool by all the atoms/ions in the metal – an organization of metal cation islands in a sea of electrons – electrons delocalized throughout the metal structure • Bonding results from attraction of cation for the delocalized electrons
- 23. Hydrogen Bonding These three bonds all have; • A strong permanent dipole • A hydrogen atom • An atom with lone pair electrons The three types of bonds which give molecules significant hydrogen bonding are; (i) N – H (ii) O – H (iii) F – H
- 25. Hydrogen Bonding Hydrogen bonding in water results in some unusual properties; • Higher than expected boiling point • High specific heat capacity (absorbs a lot of heat energy with only a small change in temperature) • Ice is less dense than water
- 34. Valence Bond (VB) Theory The basic principle of VB theory: A covalent bond forms when the orbitals of two atoms overlap and a pair of electrons occupy the overlap region. The space formed by the overlapping orbitals can accommodate a maximum of two electrons and these electrons must have opposite (paired) spins. The greater the orbital overlap, the stronger the bond. Extent of orbital overlap depends on orbital shape and direction.
- 35. Figure 11.1 Orbital overlap and spin pairing in H2. A covalent bond results from the overlap of orbitals from two atoms. The shared space is occupied by two electrons, which have opposite spins.
- 36. Figure 11.2 Orbital orientation and maximum overlap. Hydrogen fluoride, HF. Fluorine, F2. The greater the extent of orbital overlap, the stronger the bond.
- 37. VB Theory and Orbital Hybridization The orbitals that form when bonding occurs are different from the atomic orbitals in the isolated atoms. If no change occurred, we could not account for the molecular shapes that are observed. Atomic orbitals “mix” or hybridize when bonding occurs to form hybrid orbitals. The spatial orientation of these hybrid orbitals correspond with observed molecular shapes.
- 38. Features of Hybrid Orbitals The number of hybrid orbitals formed equals the number of atomic orbitals mixed. The type of hybrid orbitals formed varies with the types of atomic orbitals mixed. The shape and orientation of a hybrid orbital maximizes overlap with the other atom in the bond.
- 39. Figure 11.3 Formation and orientation of sp hybrid orbitals and the bonding in BeCl2. orbital box diagrams atomic orbitals hybrid orbitals One 2s and one 2p atomic orbital mix to form two sp hybrid orbitals.
- 40. Figure 11.3 continued box diagram with orbital contours Overlap of Be and Cl orbitals to form BeCl2.
- 41. Figure 11.4 The sp2 hybrid orbitals in BF3. Mixing one s and two p orbitals gives three sp2 hybrid orbitals. The third 2p orbital remains unhybridized.
- 42. Figure 11.4 continued The three sp2 orbitals point to the corners of an equilateral triangle, their axes 120° apart. Each half-filled sp2 orbital overlaps with the half-filled 2p orbital of a F atom.
- 43. Figure 11.5 The sp3 hybrid orbitals in CH4. The four sp3 orbitals adopt a tetrahedral shape.
- 44. Figure 11.6 The sp3 hybrid orbitals in NH3. The N lone pair occupies an sp3 hybrid orbital, giving a trigonal pyramidal shape.
- 46. Molecular Orbital (MO) Theory The combination of orbitals to form bonds is viewed as the combination of wave functions. Atomic wave functions (AOs) combine to form molecular wave functions (MOs). Addition of AOs forms a bonding MO, which has a region of high electron density between the nuclei. Subtraction of AOs forms an antibonding MO, which has a node, or region of zero electron density, between the nuclei.
- 47. Molecular Orbital Diagrams An MO diagram, just like an atomic orbital diagram, shows the relative energy and number of electrons in each MO. The MO diagram also shows the AOs from which each MO is formed. Bond order is calculated as follows: ½[(# of e- in bonding MO) – (# of e- in antibonding MO)]
- 48. Figure 11.17 MO diagram for H2. H2 bond order = ½ (2 − 0) = 1
- 49. Electrons in Molecular Orbitals • MOs are filled in order of increasing energy. • An MO can hold a maximum of 2 e- with opposite spins. • Orbitals of equal energy are half-filled, with spins parallel, before pairing spins. Electrons are placed in MOs just as they are in AOs. A molecular electron configuration shows the type of MO and the number of e- each contains. For H2 the configuration is (σ1s)2.