Transcript Chapter 5
Chapter 5 Electrons in Atoms Rutherford’s Model Not Complete • Couldn’t explain what electrons were doing • Why are they not pulled toward the nucleus? • Couldn’t explain the chemical properties of atoms • When burned, some elements produce visible light Electromagnetic • Form of energyRadiation that exhibits wave-like behavior as it travels • Visible light, microwaves, x-rays, radiowaves • Electromagn etic Wave Characteristics • • Wavelength (λ) = shortest distance between equivalent points on a wave Frequency (ʋ) = # waves that pass a given point per second • • • Amplitude = wave height from the origin to the crest Travel at 3 x 108 m/s in a vacuum (c) • • • Measured in Hz (1/sec) c = ʋλ Wavelength and frequency are inversely related Energy increases with smaller λ and higher ʋ Quanta • Waves don’t explain all of lights’ characteristics • Quantum concept = matter can gain or lose energy only in small and specific amounts called quanta • Quanta = the minimum amount of energy to be gained or lost by an atom • E = hʋ where h = Planck’s constant = 6.626 x 10-34 Js Photons • Einstein proposed that light has both wave and particle properties • Photon = particle of electromagnetic radiation with no mass that carries a quantum of energy Electrons Are Important • • • All of an atom’s properties are related to its electrons Ground state = lowest allowable energy state of an atom Excited state = state when atom gains energy • • • Can only gain energy in specific quanta Atom will be excited for a moment then fall back to the ground state and emit the same amount of energy Energy released (photon) can usually be seen as visible light Atomic Orbitals • • 3-D region around the nucleus that describes the probable location of an electron (Schrodinger) Made up of principal energy levels (n = 1-7) • • • n = 1 is low energy, close to nucleus n = 7 is high energy, far from nucleus Each principal energy level has sublevels Sublevels • 4 different sublevels with different numbers of orbitals that electrons can occupy • s = 1 orbital • p = 3 orbitals • d = 5 orbitals • f = 7 orbitals • Each orbital contains 2 electrons Sublevels • 4 different sublevels with different numbers of orbitals that electrons can occupy • s = 1 orbital • p = 3 orbitals • d = 5 orbitals • f = 7 orbitals • Each orbital contains 2 electrons Sublevels • 4 different sublevels with different numbers of orbitals that electrons can occupy • s = 1 orbital • p = 3 orbitals • d = 5 orbitals • f = 7 orbitals • Each orbital contains 2 electrons Sublevels • 4 different sublevels with different numbers of orbitals that electrons can occupy • s = 1 orbital • p = 3 orbitals • d = 5 orbitals • f = 7 orbitals • Each orbital contains 2 electrons Organization • Principal Level • n=1 • n=2 • n=3 • n=4 Sublevel Orbitals s 1 s, p 4 s, p, d 9 s, p, d, f 16 Electron Configuration • Arrangement of electrons in the atom • Usually in the most stable ground state Aufbau Principle • Each electron occupies the lowest energy orbital available • Energy sublevels have different energies • All orbitals in each sublevel are equal in energy • s<p<d<f • Orbitals in one principal level can overlap with another principal level • 4s < 3d Energy Level Diagonals Pauli Exclusion Principle • Maximum of two electrons in an orbital, but only if they have opposite spins • Example: Argon (atomic # = 18) • 1s__2s__2p__ __ __3s__3p__ __ __ Hund’s Rule • Single electrons with the same spin need to occupy each equal-energy orbital before additional electrons with opposite spin can occupy the same orbital • Ex: Phosphorus (atomic # = 15) • 1s__2s__2p__ __ __3s__3p__ __ __ Example • Sodium (atomic # = 11) • 1s __2s__2p__ __ __3s__ 2 6 1 • OR 1s 2s 2p 3s 2 • Calcium (atomic # = 20) • Silver (atomic # = 47) Shorthand • Find the first noble gas that has an atomic number lower than your given element • Put that symbol in brackets [He] • Complete the rest of the configuration for all of the electrons after the noble gas • Example: Magnesium (atomic # = 12) • [Ne]3s 2 Noble Gas • Helium 1s Configurations 2 • Neon • Argon • Krypton 4p6 2 2 1s 2s 6 2p 1s2 2s2 2p6 3s2 3p6 2 2 1s 2s 6 2 2p 3s 6 2 3p 4s 10 3d Valence Electrons • Valence electrons = electrons in the outermost primary energy level • These are the only electrons that determine chemical properties • They participate in bonding • Octet Rule = 8 electrons in the outermost energy level is the most stable • Noble gases Lewis (Dot) • Represent the electrons in the Structures outermost primary energy level • Only good for non-transition metal elements • Example: Carbon (atomic # 6) • 1s 2s 2p • n = 2 is the outermost energy 2 level 2 2