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Chapter 6 Arrangement of Electrons in Atoms 0r…. Matter waves and waves that don’t matter The nature of light Dual nature of light Wave characteristics Particle characteristics Wave nature of light Electromagnetic radiation Electromagnetic (EM) radiation Form of E w/ wavelength (l) behavior Speed = 3.0 x 1010 cm/s (speed of light) Wavelength (l) distance between pts. on a wave Frequency (n) # of waves that pass a given pt. in a specific time frequency C = ln Therefore, as l decreases, n increases C = speed of light, (186,000 miles/s, or 299,792,458 m/ s) Continuous spectrum All ls in a given range included Electromagnetic (EM) spectrum All EM radiation Particle nature of light Photoelectric effect Emission of e- by certain metals when light shines on them Max Planck (1900) When a hot object loses E, it is lost in sm. Specific amts. Called quanta Quantum Finite quantity of E that can be gained or lost by an atom Photon Individual quantum of light Albert Einstein (1905) Higher n = higher E Absorb. Of photons of a specific E explains photoelectric effect Dual (wave-particle) nature of light Important formulas E = hn h (Plank’s constant) = 6.626 x 10 -34 J . S n (frequency) c = ln c (speed of light) = 186,000 miles/s, or 299,792,458 m/ s Hydrogen atom spectrum Pass high voltage through H2 gas gas glows pass light through prism bright line spectrum Bright line spectrum Due to e-s boosted to high E state (excited state), then dropping to the ground state Lines represent E given off when e-s drop to ground state Hydrogen spectrum E of photon = difference between ground and excited state spectroscope Flame test Bohr Model of the atom (1913) The Hydrogen e- can circle the H nucleus only in certain orbits (like rungs of a ladder) Definite orbits occupied by electron particles Worked w/ H atom only According to this theory an electron moving between orbits would disappear from one and reappear instantaneously in another without visiting the space between “Quantum leap” “An electron doesn’t exist until it is observed” “Until it is observed an electron must be regarded as being at once everywhere and nowhere” Dennis Overbye Schrödinger Model (1926) Wave properties of atoms Worked w/ all atoms e- in orbitals e- clouds Can not pinpoint location of e- and path at a given instant immutable property of the universe Quantum numbers “Electron address” Location of e-s in the atom Quantum number 1 “Pennsylvania” Principle quantum number (main energy level) n= 1,2,3……7 Quantum number 2 “Hollidaysburg” Orbital quantum number (shape of orbital) s,p,d,f Quantum number 3 “N. Montgomery St.” Magnetic quantum number (orientation of orbital about the nucleus) Quantum number 4 “1510” Spin Quantum number (two possible states of electron) +1/2 or -1/2 Arrangement of electrons Arrangement of electrons Main energy levels: 1,2,3….. Sublevels: s,p,d,f Orbitals Each Each Each Each s has 1 p has 3 d has 5 f has 7 Each orbital can hold a max of 2 e- Orbital notation Unoccupied orbital ___ Orbital with 1 e- Orbital with 2 e- e.g. H # 1s # or $ #$ He #$ 1s Orbital notation Electron configuration notation Uses superscripts instead of lines e.g. H 1s1 or He 1s2 Electron dot notation Uses only e- in highest (outermost) main energy levels e.g. .Na . .He Aufbau (building up) principle Electrons occupy lowest energy orbital that will receive them, e.g. hydrogen’s electron goes into the 1s orbital Hund’s rule Orbitals of equal E are each occupied by one e- before 2nd e- is added, all e- in singly occupied orbitals must have same spin Pauli exclusion principle No two e- in same atom have the same set of four quantum numbers Electron fill chart Shorthand notation exceptions e.g. copper [Ar] 4s1 3d10