Unit 6B: Bonding Theory and Atomic Structure
Electromagnetic Radiation and Spectrum
gamma rays, x rays, UV light
high frequency, short wavelength
Visible light (VIBGYOR) (400-750 nm)
Infrared, microwaves, radio
low frequency, long wavelength
 frequency and wavelength inversely related, as demonstrated by equation c=λν
 c-speed of light constant-2.9979x108m/s
λ-wavelength(m)
ν-frequency(Hz or s-1)
Quantum Theory
 Einstein—light behaves as if its consists of quantized energy packets, meaning that energy can
have only certain allowed values, given by the equation Ephoton=hν
 Ephoton (J)
h=Planck’s constant—6.626x10-34 J-sec
ν-frequency (Hz or s-1)
 Ephoton=Eremove electron + Ekinetic
Quantum numbers
 n-principal quantum number, shell
i.e. 3s, n=3
 l-subshell (s=0, p=1, d=2, f=3)
i.e. n=3, can have l=0,1,2
 ml-magnetic (each box has number from –l to l)
i.e. 3p4, 1st box, ml= -1
 ms-spin (±1/2, if up +, if down -)
i.e. 3p4, points down, ms= -1/2
DeBroglie Wavelength
 Matter has characteristic wavelength that depends on its momentum, mv, λ=h/mv
 λ-wavelength(m) h-Planck’s constant-6.626x10-34 J-s m-mass (kg) (e-=9.11x10-31kg) v-velocity
(m/s)
Bohr’s Model
 Proposed a model of the hydrogen atom that explains its line spectrum
 Light emitted when the electron drops from a higher energy state to a lower energy state; light
absorbed to excite the electron from a lower energy state to a higher one
Electron Configuration of Atoms and Ions
 Use periodic table to write electron configuration
 D block-5 orbitals that can hold 2 electrons each for a total of 10 e-. D block one behind s/p block
 F block-7 orbitals @ 2e-each=14e-. 2 behind s/p block, one behind d block
 Core electron configuration—use largest noble gas that is smaller than atom/ion, then write
additional electrons
 When determining configurations for cations, remove electron first from orbital with largest
quantum number, n. For example, if you have 4s23d104p2 and you have element3+, first remove
electrons from 4p, then from 4s so you have 4s13d10
 Electron configurations are most stable when the orbitals are full or half-full
Valence Bond Theory—Covalent bonds formed when atomic orbitals on neighboring atoms overlap
 The greater the overlap between two orbitals, the stronger the bond that is formed
Hybridization—Mixing of s, p, and d orbitals to form hybrid orbitals.
 linear=sp, trigonal planar=sp2, tetrahedral=sp3, trigonal bipyramidal=sp3d,
octahedral=sp3d2
Sigma and Pi Bonds
 Sigma bonds are covalent bonds formed from end to end overlap
 Pi bonds are formed from the sideways overlap of p orbitals
 Each single bond consists of one sigma bond, 2x bond-1 sigma, 1 pi, 3x-1 sigma, 2 pi
Molecular Orbital Theory
 Electrons exist in allowed energy states called molecular orbitals (MOs)
 Like an atomic orbital, an MO can hold two electrons of opposite spin
 Occupation of bonding MOs favors bond formation, whereas occupation of antibonding MOs
(denoted with an *) is unfavorable