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Hollywood High School School for Advanced Studies AP Chemistry Mr. Brombach

Hollywood High School School for Advanced Studies AP Chemistry Mr. Brombach. Unit I Structure of an Atom. Unit I. Schedule ____________________________________________________. Lesson 1.1. Introduction Lesson 1.2 . Composition of an Atom. Isotopes

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Hollywood High School School for Advanced Studies AP Chemistry Mr. Brombach

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  1. Hollywood High School School for Advanced Studies AP Chemistry Mr. Brombach

  2. Unit I Structure of an Atom

  3. Unit I.Schedule____________________________________________________ Lesson 1.1. Introduction Lesson 1.2. Composition of an Atom. Isotopes Lesson 1.3. The Nature of Light. Electromagnetic Spectrum Lesson 1.4. Bohr’s Model of an Atom. Wave- Particle Nature of an Electron Lesson 1.5. Orbitals. Quantum Numbers. Lesson 1.6. PracticeQuantum Numbers Lesson 1.7. Unit Review Lesson 1.8. Test # 1

  4. Lesson 1.1. Introduction ____________________________________________________________________________________ • Go to Hollywood HS website • Open Mr. Brombach’s web log: • Go to AP-Chemistry • Find the following: • Syllabus • Unit schedule • HW assignments • Handouts • Lecture notes • Lab Assignments

  5. HW Format____________________________________________ Name_____________________Period____ Date_________ #_____ HW 1.1. # 2.42, p.80 Question……………………………………………………… Answer……………………………………………………….. _______________________________________________ # 2.46, p.80 Question……………………………………………………… Answer……………………………………………………….. _____________________________________________

  6. Lesson 1.2.Composition of an Atom. Isotopes_____________________________________________________________________________________________ MicroworldAtomsMolecules Elements Compounds MacroworldPure Substances Mixtures Matter

  7. Lesson 1.2.Composition of an Atom. Isotopes_____________________________________________________________________________________________ • Physical and chemical properties of a substance depend on its chemical structure • That includes the arrangement of the atoms in a molecule and types of bonding between them

  8. Lesson 1.2.Composition of an Atom _____________________________________________________________________________________________ • Each atom is represented by the notation mass numberA X symbolatomic numberZ • Atomic number (Z) equals the number of protons in the nucleus Z = # p+

  9. Lesson 1.2.Composition of an Atom_____________________________________________________________________________________________ • An atom is neutral (the number of protons equals the number of electrons) # p+ = # e- • Mass number (A) is the total number of protons and neutrons A =# p+ + # no • Number on neutrons can be found from the formula: #no = A - Z

  10. Lesson 1.2.Composition of an Atom_____________________________________________________________________________________________ • Since the mass of an atom is so small, to measure atomic mass we use a group called “Dalton (D)” (the old name amu) 1 D = 1.66 x 10-24 g • Atomic mass in the Periodic Table is done in Daltons • For example, the mass of carbon atom 12 Cis exactly mc = 12 D6 or (12)(1.66 x 10-24 g) = 1.99 x 10-23 g

  11. Lesson 1.2. Isotopes _________________________________________________________________________________ • Not all atoms of a particular element have the same mass • The difference in their mass number (A) is due to the presence of different number of neutrons (no) • For ex.: There are two types of Boron (B) atom: • 10B or Boron – 10 (5 p+ + 5 no) • 11B or Boron – 11 (5 p+ + 6 no)

  12. Lesson 1.2. Isotopes _________________________________________________________________________________ • Isotopes of an element are atoms that have different number of neutrons and, therefore, different mass numbers • An element occurs as a mixture of isotopes • The atomic mass of an element is the average of its isotopic masses according to their natural abundances

  13. Lesson 1.2. Isotopes _________________________________________________________________________________

  14. Lesson 1.2. Isotopes_________________________________________________________________________________ • Find average atomic mass of Mg • Atomic mass portion: 24Mg = 23.9850 x 0.7899 = 18.9458 25Mg = 24.9858 x 0.1000 = 2.4986 26Mg = 25.9826 x 0.1101 = 2.8607 24.3024 D

  15. Lesson 1.3. Nature of Light _________________________________________________________________________________ • How do we know about atoms, as we cannot see them? • To learn about atomic structure, scientists treat matter with different kind of energy (heat, electricity, ionization, magnetic field…) EnergyAn ElementEMR • As a result, the matter gives away electromagnetic radiation (EMR) • By studying EMR, the scientists are able to develop models of the atom

  16. Lesson 1.3. Nature of Light _________________________________________________________________________________ • EMR (light) travels as a wave • It is described by two independent variables: wavelength and frequency • Wavelength (λ – lambda) is the distance (nm) the wave travels during one cycle • Frequency (√ - nu) is the number of cycles the wave undergoes per second (1/s or Hz) • Speed of light in vacuum is constant and equals 3.00 x 108 m/s

  17. Lesson 1.3. Nature of Light _______________________________________________________________________ 750 nm 400 nm

  18. Lesson 1.3. Nature of Light _______________________________________________________________________ • The wavelength is inversely proportional to the frequency Cλ= ----- (1) √ C – speed of light, m/s λ – wavelength, nm √ - frequency, 1/s or Hz

  19. Lesson 1.3. Nature of Light _______________________________________________________________________ • At the beginning of 20th century, the three phenomena involving matter and light could not be explained based on the wave nature of light: • The pattern of intensity and wavelength of light emitted from hot, dense objects (blackbody radiation) • The electric current generated when light shines on a metal plate (photoelectric effect) • The individual colors emitted from electrically (or thermally) excited gases (atomic spectra)

  20. Lesson 1.3. Nature of Light _______________________________________________________________________ • Explaining these phenomena required a radically new view of energy (light): • Plank’s quantum hypothesis (1900) • A beam of light is not a continuous stream of energy; instead the beam consists of zillions of small, discrete packets of energy, each called quantum • Einstein’s particulate nature of light (1905) • The quanta of light behave much like tiny particles of matter, each quantum of light was called a photon

  21. Lesson 1.3. Nature of Light _______________________________________________________________________ • Thus, the light has properties of both, a wave and a particle • To represent this duality, the photon is illustrated as a burst of light with a wave drawn inside the burst • The scientists are free to choose which of these two modes fits their needs the best

  22. Lesson 1.3. Nature of Light _______________________________________________________________________ • The energy carried by the wave is directly proportional to the frequency E = h√ (2) E – energy, J h – Plank’s constant (6.626 x 10-34 J•s) √ - frequency, 1/s or Hz • The most powerful type of EMR are gamma rays that have the highest frequency 2.756 x 102 2.756 x 102

  23. Lesson 1.4. Bohr’s Model of an Atom_______________________________________________________________________ Energy is absorbedEnergy is released• ground stateexcited state • Ground state or stationary state is the most stable (the lowest level of energy) • To move to the higher level an object absorbs energy and turns to excited state (less stable) • To go back to stable state, the object gives away (emits) energy •

  24. Lesson 1.4. Bohr’s Model of an Atom _______________________________________________________________________ • Accepting Plank’s and Einstein’s idea about quantized energy, Bohr proposed that the hydrogen atom had only certain energy levels • If gaseous hydrogen is turned from ground state to excited state by electric discharge, it goes back to ground state by emitting EMR

  25. Lesson 1.4. Bohr’s Model of an Atom _______________________________________________________________________ See p.264 fig. 7.9

  26. Lesson 1.4. Bohr’s Model of an Atom _____________________________________________________

  27. Lesson 1.4. Atomic Spectra _______________________________________________________________________ • As emitted EMR passes through a slit and a prism, the EMR will be divided into individual wavelength • The EMR does not create a continuous spectrum, or rainbow, as sunlight does • Rather, it produces a line spectrum – a series of fine lines of individual colors separated by colorless spaces

  28. Lesson 1.4.Atomic Spectra ___________________________________________________________ • The pattern of wavelength (frequencies) formed by • a given element is referred to as element’s atomic • spectrum • The wavelength at which the colored lines occur • is individual characteristic of the element, its • “fingerprint” that allows to identify an element

  29. Lesson 1.4. Atomic Spectra _______________________________________________________________________ • To find the position and wavelength of any line in a given series, use the Rydberg equation 1 1 1 ------ = R (------ - ------) (3) λ n12 n22 λ – wavelength of a particular spectral line n1, n2 – integers representing energy levels (n2 >n1) R – Rydberg constant = 1.097 x 107 1/m

  30. Lesson 1.3. Bohr’s Model of an Atom_______________________________________________________________________ Lesson 1.4. Bohr’s Model of an Atom______________________________________ ∆ In an atom, an electron can move from one energy level to another only by absorbing or emitting a photon of energy

  31. Lesson 1.4. Bohr’s Model of an Atom _______________________________________________________________________ • The amount of energy an atom emits is the difference between energy of final and initial state ∆ Ephoton = Efin – Ein = = Eexcited state – Egroundstate(4) • The greater the energy level, the greater the energy n E • The energy of any excited state equals: E = -2.18 x 10-18(1/n2), J(5) • The energy emitted of absorbed by H atom ∆E = -2.18 x 10-18(1/nfin2 – 1/nin2) (6)

  32. Lesson 1.5. The Wave-particle Nature of an Electron _______________________________________________________________________ • Does photon have a mass? • The famous Einstein’s equation states the relationship between energy and mass E = mc2(7) E – energy, J m – mass, g c – speed of light, m/s

  33. Lesson 1.5. The Wave-particle Nature of an Electron _______________________________________________________________________ • As light exists as a wave and as a particle, each model has the equation of energy: E = mc2(mass represents a particle) E = h√(frequency represents a wave) mc2 = h√√ = c/λmc2 = hc/λ h m = ----- (8) λc m – mass of photon (EMR or particle)

  34. Lesson 1.5. The Wave-particle Nature of an Electron _______________________________________________________________________ • De Broglie proposed the equation, which connects wave and particle properties of any object such as planet, baseball, or electron h λ = ----- (9) m v v – velocity (speed), m/s • Since electron moves with a speed close to the speed of light, it also exists as a wave and as a particle (duality)

  35. Lesson 1.5. The Atomic Orbital _______________________________________________________________________ • If an electron has the properties of both a particle and a wave, what can we determine about its position in the atom? • The Heisenberg’s Uncertainty Principle states that it is impossible to know simultaneously the exact position and velocity of a particle • That means that we cannot prescribe exact paths for electrons, such as the circular orbits of Bohr’s model

  36. Lesson 1.5. The Atomic Orbital _______________________________________________________________________ • The wave motion of objects on the atomic scale is examined in the field of quantum mechanics • In 1926, Schrodinger formulated an equation from which the probability of finding the electron in hydrogen atom could be determined • If we could plot the positions of an electron of a given energy over time as a series of tiny dots, the resulting pattern would resemble what is called a probability cloud

  37. Lesson 1.5. The Atomic Orbital _______________________________________________________________________ • The electron density diagram represents the probability of finding the electron at a particular point at a given distance r along a line from the nucleus outward • The probability of the electron being far from the nucleus is very small, but not zero • An atomic orbital, like a probability cloud, specifies a volume of space where the electron is most likely to be found

  38. Lesson 1.5. The Atomic Orbital _______________________________________________________________________

  39. Lesson 1.5. The Atomic Orbital _______________________________________________________________ S-orbital (1)

  40. Lesson 1.5. The Atomic Orbital _______________________________________________________________ p-orbital (3)

  41. Lesson 1.5. The Atomic Orbital _______________________________________________________________ d-orbital (5)

  42. Lesson 1.5. The Atomic Orbital _______________________________________________________________ f-orbital (7)

  43. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________ • Each orbital can be described by a set of characteristics called quantum numbers (QN): • n – principal QN (characterizes energy level and size of the orbital) • l – azimuthal QN (energy sublevel and shape) • ml – magnetic QN (orientation in space)

  44. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________ Values Principal QN n = 1, 2, 3, 4, 5 … The greater the “n” value, the higher energy level and the bigger the orbital 5 4 3 2 1

  45. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________ Values Azimuthal QN: “l” = 0, 1, 2, 3, 4…. n-1 • “l” represents: • Energy sublevels: l = 0(s); l = 1(p); l = 2(d); l = 3(f) • Shape of the orbital: s – sphere; p – double-lobe d and f – shape varies f d p s sublevels Energy level

  46. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________ Values Magnetic QN: “ml” = -l…0…+l “ml”represents the orientation of the orbital in space: • l = 0 ml = 0 (only 1 orientation) • l = 1 ml = -1, 0, +1 (3 orientations x, y, z) • l = 2 ml = -2, -1, 0, +1, +2 (5 orientations) • l = 3 ml = -3, -2, -1, 0, +1, +2, +3 (7 orientations)

  47. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________ -On a particular energy level, there are: • 1 s-orbital • 3 p-orbitals • 5 d-orbitals • 7 f-orbitals

  48. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________

  49. Lesson 1.6.Quantum Numbers _________________________________________________________________________________________ Lesson 1.5.a. The Atomic Orbital _______________________________________________________________ l 2px l n ml

  50. Lesson 1.6.Quantum Numbers_________________________________________________________________________________________ • The total number of orbitals on a particular energy level equals: # orbitals = n2 4f 4d 4p 4s 42 = 16 orbitals n = 4 3d 3p 3s n = 3 32 = 9 orbitals 2p 2s n = 2 22 = 4 orbitals 1s n = 1 12 = 1 orbital

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