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Ch. 5 - Electrons in Atoms Wavelength () - length of one complete wave measured in m, cm, or nm Frequency () - # of waves that pass a point during a certain time period, In light it tells us which color it is hertz (Hz) = 1/s Amplitude (A) - distance from the origin to the trough or crest how much energy the wave is carrying. It is the height of the wave. It is measured in meters. In SOUND it tells us how LOUD it is. In LIGHT it tells how BRIGHT it is. crest A greater amplitude origin A trough greater frequency • To understand the electronic structure of atoms we must understand light and how it is emitted or absorbed by substances. • We will examine visible light a type of Electromagnetic Radiation (EM) which carries (radiant) energy through space (speed of light) and exhibits wavelike behavior. • Also need to think of light as particle, to help understand how EM radiation and atoms interact H I G H E N E R G Y L O W E N E R G Y Move through a vacuum at the ‘speed of light’ 3.00 x 108 m/s Behaves like waves that move through water, which are the result of a transfer of energy to the water (from a stone), expressed as up and down movement of water Both electric and magnetic properties Wave Speed = (distance between wave peaks) x (frequency) = (wavelength) x (frequency) EM radiation moves through a vacuum at the “speed of light” 3.00 x 108 m/s also called c. A lower energy wave (infrared and red) has a longer wavelength() and lower frequency(f) A higher energy wave (blue - violet) has a shorter wavelength() and higher frequency(f). Frequency & wavelength are inversely proportional c = c: speed of light (3.00 108 m/s) : wavelength (m, nm, etc.) : frequency (Hz) EX: Find the frequency of a photon with a wavelength of 434 nm. GIVEN: WORK: =c =? = 434 nm = 4.34 10-7 m = 3.00 108 m/s -7 m 8 4.34 10 c = 3.00 10 m/s = 6.91 1014 Hz Planck (1900) Observed - emission of light from hot objects Concluded - energy is emitted (absorbed or released) in small, specific amounts (quanta) Quantum - smallest energy packet that can be emitted or absorbed as EM radiation by an atom. Planck proposed that the energy, E, of a single quantum energy packet equals a constant (h) times its frequency 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) EX: Find the energy of a red photon with a frequency of 4.57 1014 Hz. GIVEN: WORK: E=? E = h = 4.57 1014 Hz E = (6.6262 10-34 J·s) h = 6.6262 10-34 J·s (4.57 1014 Hz) E = 3.03 10-19 J Planck (1900) vs. Classical Theory Quantum Theory Energy is always emitted or absorbed in whole number multiples of hv, such as hv, 2 hv, 3 hv, 4hv, …. The allowed energies are quantized, that is their values are restricted to certain quantities. The notion of quantized rather than continuous energies is strange. Consider a ramp and a staircase, on a ramp you can vary the length your steps and energy used on the walk up. When walking up steps you must exert exactly the specific amount of energy needed to reach the next step. Your steps on steps are quantized, you cannot step between them. Einstein (1905) Observed – photoelectric effect Dispersed light falls on metal samples, the different frequencies produce different energetic photoelectrons Einstein (1905) Concluded - light has properties of both waves and particles (photons) “wave-particle duality” Photon - particle of light that carries a quantum of energy Used planck’s quantum theory to deduced that: Ephoton = hv Ch. 5 - Electrons in Atoms Set of frequencies of EM waves emitted by atoms an element when they absorb electrical energy, eˉ get excited, become somewhat unstable and release energy in the form of light excited state ENERGY IN PHOTON OUT ground state 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 Ground state: lowest allowable atomic electron energy state Excited state: any higher energy state 65 4 Energy 3 2 1 of photon depends on the difference in energy levels Bohr’s calculated energies matched the IR, visible, and UV lines for the H atom Each element has a unique bright-line emission spectrum. Helium Examples: “Atomic Fingerprint” Iron Now, we can calculate for all elements and their electrons Ch. 5 - Electrons in Atoms Louis de Broglie (1924) Proposed eˉ in their orbits behave like a wave Wavelength of an eˉ depends on its mass(m) and its velocity (v): λ = _h _ mv EVIDENCE: DIFFRACTION PATTERNS VISIBLE LIGHT ELECTRONS Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time Attempting to observe an electron’s position changes its momentum and attempting to observe an electron’s momentum changes its position. Therefore electrons cannot be locked into well-defined circular orbits around the nucleus. Schrödinger Wave Equation (1926) proposed a wave equation incorporating both the wave and particle nature of the electron. The result of the equation, wave functions, shows the probability that an electron will be in a certain region of space at a given instant. This electron density is represented by a distribution of dots which represents where electrons are located about 90% of the time finite # of solutions quantized energy levels defines probability of finding an e- Orbital (“electron cloud”) a specific distribution of electron density in space. Each orbital has a characteristic energy and shape. Orbital Specify the “address” of each electron in an atom UPPER LEVEL 1. Principal Quantum Number (n = 1, 2, 3, …) (see periodic table left column) Indicates the relative size and energy of atomic orbitals As (n) increases, the orbital becomes larger, the electron spends more time farther from the nucleus Each major energy level is called a principle energy level Ex: lowest level = 1 ground state, highest level = 7 excited state 2. Energy Sublevel Defines the shape of the orbital (s, p, d, f) # of orbital related to each sublevel is always an odd # s = 1, p = 3, d = 5, f = 7 Each orbital can contain at most 2 electrons s p d f Subscripts x, y, z designates orientation Specifies the exact orbital within each sublevel px py pz 4. Spin Quantum Number ( ms ) Electron spin +½ or -½ An orbital can hold 2 electrons that spin in opposite directions. Pauli Exclusion Principle A maximum of 2 electrons can occupy a single atomic orbital Only if they have opposite spins 1. Principal # 2. Energy sublevel 3. Orientation 4. Spin # energy level (s,p,d,f) x, y, z exact electron Ch. 5 - Electrons in Atoms IV. Electron Configuration A. General Rules Aufbau Principle Electrons fill the lowest energy orbitals first. “Lazy Tenant Rule” A. General Rules Hund’s Rule Within a sublevel, place one e- per orbital before pairing them. “Empty Bus Seat Rule” WRONG RIGHT Notation s p 1 2 3 4 5 6 7 f(n-2) d (n-1) 6 7 © 1998 by Harcourt Brace & Company B. Notation Orbital Diagram O 8e- 1s 2s Electron Configuration 2 2 4 1s 2s 2p 2p B. Notation Longhand Configuration S 16e- 1s2 2s2 2p6 3s2 3p4 Core Electrons Valence Electrons Valence electrons: determine chemical properties of that element & are the electrons in the atoms outermost orbital Shorthand Configuration S 16e- [Ne] 3s2 3p4 Shorthand Notation Shorthand Configuration Core e-: Go up one row and over to the Noble Gas. Valence e-: On the next row, fill in the # of e- in each sublevel. 1 2 3 4 5 6 7 C. Periodic Patterns Example - Germanium 1 2 3 4 5 6 7 [Ar] 2 4s 10 3d 2 4p D. Stability Full energy level Full sublevel (s, p, d, f) Half-full sublevel 1 2 3 4 5 6 7 D. Stability Electron Configuration Exceptions Copper EXPECT: [Ar] 4s2 3d9 ACTUALLY: [Ar] 4s1 3d10 Copper gains stability with a full d-sublevel. D. Stability Electron Configuration Exceptions Chromium EXPECT: [Ar] 4s2 3d4 ACTUALLY: [Ar] 4s1 3d5 Chromium gains stability with a half-full d-sublevel. D. Stability Ion Formation Atoms gain or lose electrons to become more stable. Isoelectronic with the Noble Gases. 1 2 3 4 5 6 7 D. Stability Ion Electron Configuration Write the e- config for the closest Noble Gas EX: Oxygen ion O2- Ne 2O 10e [He] 2 2s 6 2p Read Section 5-3!