Presentation on theme: "I. Waves & Particles (p. 91 - 94) Ch. 4 - Electrons in Atoms."— Presentation transcript:
I. Waves & Particles (p ) Ch. 4 - Electrons in Atoms
A. Waves zWavelength ( ) - length of one complete wave zFrequency ( ) - # of waves that pass a point during a certain time period yhertz (Hz) = 1/s zAmplitude (A) - distance from the origin to the trough or crest
A. Waves A greater amplitude (intensity) greater frequency (color) crest origin trough A
B. EM Spectrum LOWENERGYLOWENERGY HIGHENERGYHIGHENERGY
B. EM Spectrum zFrequency & wavelength are inversely proportional c = c:speed of light ( m/s) :wavelength (m, nm, etc.) :frequency (Hz)
B. EM Spectrum GIVEN: = ? = 434 nm = m c = m/s WORK : = c = m/s m = Hz zEX: Find the frequency of a photon with a wavelength of 434 nm.
C. Quantum Theory zPlanck (1900) yObserved - emission of light from hot objects yConcluded - energy is emitted in small, specific amounts (quanta) yQuantum - minimum amount of energy change
C. Quantum Theory zPlanck (1900) vs. Classical TheoryQuantum Theory
C. Quantum Theory zEinstein (1905) yObserved - photoelectric effect
C. Quantum Theory zEinstein (1905) yConcluded - light has properties of both waves and particles wave-particle duality yPhoton - particle of light that carries a quantum of energy
C. Quantum Theory E:energy (J, joules) h:Plancks constant ( J·s) :frequency (Hz) E = h zThe energy of a photon is proportional to its frequency.
C. Quantum Theory GIVEN: E = ? = Hz h = J·s WORK : E = h E = ( J·s ) ( Hz ) E = J zEX: Find the energy of a red photon with a frequency of Hz.
II. Bohr Model of the Atom (p ) Ch. 4 - Electrons in Atoms
A. Line-Emission Spectrum ground state excited state ENERGY IN PHOTON OUT
B. Bohr Model ze - exist only in orbits with specific amounts of energy called energy levels zTherefore… ye - can only gain or lose certain amounts of energy yonly certain photons are produced
B. Bohr Model zEnergy of photon depends on the difference in energy levels zBohrs calculated energies matched the IR, visible, and UV lines for the H atom
C. Other Elements zEach element has a unique bright-line emission spectrum. yAtomic Fingerprint Helium zBohrs calculations only worked for hydrogen!
C. Other Elements zExamples: yIron zNow, we can calculate for all elements and their electrons – next section
III. Quantum Model of the Atom (p ) Ch. 4 - Electrons in Atoms
A. Electrons as Waves zLouis de Broglie (1924) yApplied wave-particle theory to e - ye - exhibit wave properties EVIDENCE: DIFFRACTION PATTERNS ELECTRONS VISIBLE LIGHT
B. Quantum Mechanics zHeisenberg Uncertainty Principle yImpossible to know both the velocity and position of an electron at the same time
B. Quantum Mechanics zSchrödinger Wave Equation (1926) yfinite # of solutions quantized energy levels ydefines probability of finding an e -
B. Quantum Mechanics Radial Distribution Curve Orbital zOrbital (electron cloud) yRegion in space where there is 90% probability of finding an e -
C. Quantum Numbers UPPER LEVEL zFour Quantum Numbers: ySpecify the address of each electron in an atom
C. Quantum Numbers 1. Principal Quantum Number ( n ) yMain energy level occupied the e- ySize of the orbital yn 2 = # of orbitals in the energy level
C. Quantum Numbers s p d f 2. Angular Momentum Quantum # ( l ) yEnergy sublevel yShape of the orbital
C. Quantum Numbers zn=# of sublevels per level zn 2 =# of orbitals per level zSublevel sets: 1 s, 3 p, 5 d, 7 f
C. Quantum Numbers 3. Magnetic Quantum Number ( m l ) yOrientation of orbital around the nucleus Specifies the exact orbital within each sublevel
C. Quantum Numbers pxpx pypy pzpz
zOrbitals combine to form a spherical shape. 2s 2p z 2p y 2p x
C. Quantum Numbers 4. Spin Quantum Number ( m s ) yElectron spin +½ or -½ yAn orbital can hold 2 electrons that spin in opposite directions.
C. Quantum Numbers 1. Principal # 2. Ang. Mom. # 3. Magnetic # 4. Spin # energy level sublevel (s,p,d,f) orientation electron zPauli Exclusion Principle yNo two electrons in an atom can have the same 4 quantum numbers. yEach e - has a unique address: