Atomic Theory

I. Waves & Particles (p. 91 - 94) Ch. 4 - Electrons in Atoms C. Johannesson

B. EM Spectrum HIGH ENERGY LOW ENERGY C. Johannesson

B. EM Spectrum HIGH LOW ENERGY ENERGY R O Y G. B I V red orange yellow green blue indigo violet C. Johannesson

C. Quantum Theory Planck (1900) Observed - emission of light from hot objects Concluded - energy is emitted in small, specific amounts (quanta) Quantum - minimum amount of energy change C. Johannesson

C. Quantum Theory Einstein (1905) Observed - photoelectric effect C. Johannesson

“wave-particle duality” C. Quantum Theory Einstein (1905) Concluded - light has properties of both waves and particles “wave-particle duality” Photon - particle of light that carries a quantum of energy C. Johannesson

C. Quantum Theory The energy of a photon is proportional to its frequency. E = h E: energy (J, joules) h: Planck’s constant (6.6262  10-34 J·s) : frequency (Hz) C. Johannesson

II. Bohr Model of the Atom (p. 94 - 97) Ch. 4 - Electrons in Atoms C. Johannesson

A. Line-Emission Spectrum excited state ENERGY IN PHOTON OUT ground state C. Johannesson

B. Bohr Model e- exist only in orbits with specific amounts of energy called energy levels Therefore… e- can only gain or lose certain amounts of energy only certain photons are produced C. Johannesson

B. Bohr Model 6 Energy of photon depends on the difference in energy levels Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom 5 4 3 2 1 C. Johannesson

C. Other Elements Helium Each element has a unique bright-line emission spectrum. “Atomic Fingerprint” Helium Bohr’s calculations only worked for hydrogen!  C. Johannesson

III. Quantum Model of the Atom (p. 98 - 104) Ch. 4 - Electrons in Atoms C. Johannesson

A. Electrons as Waves QUANTIZED WAVELENGTHS Louis de Broglie (1924) Applied wave-particle theory to e- e- exhibit wave properties QUANTIZED WAVELENGTHS C. Johannesson

A. Electrons as Waves QUANTIZED WAVELENGTHS C. Johannesson

A. Electrons as Waves EVIDENCE: DIFFRACTION PATTERNS VISIBLE LIGHT C. Johannesson

B. Quantum Mechanics Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time C. Johannesson

B. Quantum Mechanics Schrödinger Wave Equation (1926) finite # of solutions  quantized energy levels defines probability of finding an e- C. Johannesson

Radial Distribution Curve B. Quantum Mechanics Orbital (“electron cloud”) Region in space where there is 90% probability of finding an e- Orbital Radial Distribution Curve C. Johannesson

C. Quantum Numbers Four Quantum Numbers: Specify the “address” of each electron in an atom UPPER LEVEL C. Johannesson

C. Quantum Numbers 1. Principal Quantum Number ( n ) Energy level Size of the orbital n2 = # of orbitals in the energy level C. Johannesson

C. Quantum Numbers 2. Angular Momentum Quantum # ( l ) Energy sublevel Shape of the orbital f d s p C. Johannesson

C. Quantum Numbers n = # of sublevels per level n2 = # of orbitals per level Sublevel sets: 1 s, 3 p, 5 d, 7 f C. Johannesson

C. Quantum Numbers 3. Magnetic Quantum Number ( ml ) Orientation of orbital Specifies the exact orbital within each sublevel C. Johannesson

C. Quantum Numbers px py pz C. Johannesson

C. Quantum Numbers Orbitals combine to form a spherical shape. 2s 2px 2pz 2py 2px C. Johannesson

C. Quantum Numbers 4. Spin Quantum Number ( ms ) Electron spin  +½ or -½ An orbital can hold 2 electrons that spin in opposite directions. C. Johannesson

C. Quantum Numbers Pauli Exclusion Principle 1. Principal #  No two electrons in an atom can have the same 4 quantum numbers. Each e- has a unique “address”: 1. Principal #  2. Ang. Mom. #  3. Magnetic #  4. Spin #  energy level sublevel (s,p,d,f) orbital electron C. Johannesson

Feeling overwhelmed? Read Section 4-2! C. Johannesson