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Chapter 7

Chapter 7. Atomic Structure And Periodicity. How Often Does The Topic Appear On AP Exam? MC 10% of Questions FR Almost Every Year. Electromagnetic Radiation. Radiant energy that exhibits wavelength-like behavior and travels through space at the speed of light in a vacuum. Waves.

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Chapter 7

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  1. Chapter 7 • Atomic Structure • And • Periodicity How Often Does The Topic Appear On AP Exam? MC 10% of Questions FR Almost Every Year

  2. Electromagnetic Radiation • Radiant energy that exhibits wavelength-like behavior and travels through space at the speed of light in a vacuum.

  3. Waves • Waves have 3 primary characteristics: • 1. Wavelength: distance between two peaks in a wave. • 2. Frequency: number of waves per second that pass a given point in space. • 3. Speed: speed of light 3.00 X 108 m/s.

  4. Wavelength and frequency have an inverse relationship. • =c • λ= wavelength (in meters) lambda • ν = frequency (in cycles per second) nu • c = speed of light (m s-1) • Mnemonic Device: lambdanucharlie • The speed of light is 3.00 x 108

  5. In the flame test lab Sr(NO3)2 produced a brilliant magenta color light with a wavelength of around 6.50 x 102 nm. Calculate the frequency of the light. • Convert= c to= c/Use 3.00 x 108 m/s for the speed of light • Don’t forget to change units to meters • 1 nm = 109 • Plug ‘n Chug Do #32 p. 321

  6. Nature of Matter Planck’s Constant Planck’s Assumption: Transfer of energy is quantized, and can only occur in discrete units, called quanta. • E = change in energy, in J (kg x m2/ s2) • h = Planck’s constant, 6.63 X 10-34 J x s • = frequency, in cycle s-1 • = wavelength, in m

  7. The Cu(NO3)2 used in our flame test lab emitted a blue-green color having a wavelength of 500 nm. What is the quantum of energy emitted by the compound? Steps to solving: We know wavelength so use to calculate frequency Use To calculate the energy quantum emitted by Cu(NO3)2 Time to Plug ‘n Chug

  8. Tanning Beds? • Ultraviolet B radiation is in the wavelength range 280 to 320 nm. UV-B: • triggers direct DNA damage which in turn induces an increased melanin production • is more likely to cause a sunburn than UVA as a result of overexposure • is thought to cause the formation of moles and some types of skin cancer and causes skin aging (wrinkles before your time) • A photon of UV light possess enough energy to mutate a strand of human DNA. What is the energy of a single UVB photon with a wavelength of 300 nm?

  9. Einstein & Photoelectric Effect • Electrons are emitted from metals when light of a frequency (ν) above a specific threshold frequency (ν0 )strikes the surface. • Observed Characteristics: • No e- emitted ν < ν0 • No e- emitted ν < ν0 , even if ↑ intensity • If ν> ν0 , e- ↑ w/ ↑ intensity • If ν> ν0 , KEe- ↑ linearly ↑ intensity

  10. Einstein & Photoelectric Effect • KEelectron = ½ mv2 = hν – hν0 • m = mass of e- • v = velocity of e- • hν = energy of photon (Planck’s constant x frequency) • hν0 = threshold energy required to remove e- • Don’t confuse v (vee) for ν (nu) • Principle behind photocells in a camera light meter, security lights, and auto-open doors

  11. Energy and Mass • Energy has mass (Einstein said it 1st) • E = mc2 • E = energy • m = mass • c = speed of light • Solve for Mass • m = E / c2

  12. Energy and Mass (Evidence for the dual nature of light.)

  13. Wavelength and Mass de Broglie’s Equation • = wavelength, in m • h = Planck’s constant, 6.63 X10-34 J x s = kg m2 s-1 • m = mass, in kg • v = velocity • mv = momentum of electron • ALL particles exhibit a wave behavior. The equations allows us to calculate the wavelength of a particle. See Sample Exercise 7.3

  14. Atomic Spectrum of Hydrogen p284 • Continuous spectrum: Contains all the wavelengths of light when white light is passed through a prism. • Line (discrete) spectrum: When oneelements emission spectrum is passed through a prism only someof the wavelengths of light show up.

  15. The Bohr Model Bohr proposed: electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits. • E= energy of the levels in the H-atom • Z= nuclear charge (for H, z = 1) • n = an integer indicating the orbital of electron

  16. Transitions of Electrons As electrons move (transition) between orbits, they give off photons of light. Moving from n=3 to n=2 produces a red photon.

  17. The Bohr Model • Ground State: The lowest possible energy state for an atom (n = 1).

  18. Energy Changes in the Hydrogen Atom • E = Efinal state- Einitial state

  19. Quantum Mechanics • Hydrogen electron is visualized as a standing wave around the nucleus. • Destructive interference will occur because the wave is not a standing wave, ends are mismatched.

  20. Quantum MechanicsSchrodinger’s Equation • Based on the wave propertiesof the atom •  = wave function of coordinate x, y, & z (psi) • = mathematical operator • E = total energy of the atom • A specific wave function is often called an orbital. • Wave function is based on KE + PE = Total Energy

  21. Heisenberg Uncertainty Principle • ∆x= uncertainty in position of electron • ∆mv = uncertainty in momentum (p) of electron • h = Planck’s constant • The more accurately we know a particle’s position, the less accurately we can know its momentum. • The limitation are very small if you are using large objects like baseball but for small objects the limitations are greatly magnified.

  22. Probability Distribution p292 • square of the wave function indicates the probabilityof finding an electron at a given position, the model does not tell us when it will be at that position (fig 7.11a) Radial probability distribution is the probability distribution in each spherical shell (fig7.11b)

  23. Quantum Numbers (QN) • Principal QN(n = 1, 2, 3, . . .) - related to size and energy of the orbital. Electron requires more energy to exist further from nucleus. • Angular Momentum QN(l = 0 to n-1) - relates to shape of the orbital. l=0 (spherical), l=1 (polar), l=2 (cloverleaf). • Magnetic QN(ml = l to - l) - relates to orientation of the orbital in space relative to other orbitals. • Electron Spin QN(ms= +1/2, -1/2) - relates to the spin statesof the electrons.

  24. Quantum Numbers

  25. Orbitals correspond to blocks on Periodic Table

  26. Orbital Shapes

  27. Pauli Exclusion Principle • In a given atom, no two electrons can have the same set of four quantum numbers (n, l, ml, ms). • Therefore, an orbital can hold only two electrons, and they must have opposite spins. • If 2 electrons in outer most orbital – substance is diamagnetic • If 1 electron in outer most orbital – substance is paramagnetic

  28. History of Periodic Table • Johann Dobereiner – 1829 - found several groups of 3 (triads) • John Newland – 1864 – elements arranged in octaves • Lothar Meyer & Dimitri Mendeleev -1872 – independently produced periodic table based on mass of elements. Mendeleev given most credit because of its predictive ability. • Henry Moseley – 1913 – arranged table in increasing atomic number – today’s version

  29. Aufbau Principle • Electrons are added one by one to elements from the inner most orbital to the outer most orbital. • Aufbau is German for “building up”

  30. Hund’s Rule • The lowest energy configuration for an atom is the one having the maximum number of unpaired electrons in a particular set of orbitals. • Ex. - The rule says you will put an electron in each p orbital BEFORE you pair them up.

  31. Valence Electrons The electrons in the highest principlequantum level of an atom. Inner electrons are called core electrons.

  32. Periodic Table Classifications p307 • Representative Elements(main group): filling s and p orbitals (Na, Al, Ne, O) • Transition Elements: filling dorbitals (Fe, Co, Ni) • Lanthanide and Actinide Series(inner transition elements): filling 4fand 5forbitals (Eu, Am, Es)

  33. Ionization Energy • The quantity of energy required to remove an electron from the gaseous atom or ion.

  34. Periodic Trends First Ionization Energy • increasesfrom left to right across a period( A# means nuclear attract.) • decreases going down a group(shielding effect)

  35. Electron Affinity • The energy change associated with the addition of an electron to a gaseous atom.

  36. Periodic TrendsAtomic Radii • Measure the distance between atoms in chemical compounds. • decrease going from left to right across a period(nuclear attraction) • increasegoing down a group(shielding effect)

  37. Information Contained in the Periodic Table • 1. Each group/family member has the same valence electron configuration (these electrons primarily determine an atom’s chemistry). • 2. Should know the electron configuration of any representative element. • 3. Certain groups have special names (alkali metals, halogens, etc). • 4. Metals and nonmetals are characterized by their chemical and physical properties.

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