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. Practice Quantum 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 _____________________________________________________________________________________________ Microworld AtomsMolecules Elements Compounds Macroworld Pure 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 number A X symbol atomic number Z 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 + + # n o Number on neutrons can be found from the formula: # n o = 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 C is exactly m c = 12 D 6 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 (n o ) For ex.: There are two types of Boron (B) atom: – 10 B or Boron – 10 (5 p + + 5 n o ) – 11 B or Boron – 11 (5 p + + 6 n o )

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 _________________________________________________________________________________ Isotopic form Mass (D)Abundance, % Fraction 24 Mg23.985078.990.7899 25 Mg24.985810.000.1000 26 Mg25.982611.010.1101

14. Lesson 1.2. Isotopes _________________________________________________________________________________ Find average atomic mass of Mg Atomic mass portion: 24 Mg = 23.9850 x 0.7899 = 18.9458 25 Mg = 24.9858 x 0.1000 = 2.4986 26 Mg = 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…) Energy An Element EMR 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 10 8 m/s

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

18. Lesson 1.3. Nature of Light _______________________________________________________________________ The wavelength is inversely proportional to the frequency C λ = ----- (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 20 th 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 Js) √ - frequency, 1/s or Hz The most powerful type of EMR are gamma rays that have the highest frequency 2.756 x 10 2

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 n 1 2 n 2 2 λ – wavelength of a particular spectral line n 1, n 2 – integers representing energy levels (n 2 >n 1 ) R – Rydberg constant = 1.097 x 10 7 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 ∆ E photon = E fin – E in = = E excited state – E ground state (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/n 2 ), J (5) The energy emitted of absorbed by H atom ∆ E = -2.18 x 10 -18 (1/n fin 2 – 1/n in 2 ) (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 = mc 2 (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 = mc 2 (mass represents a particle) E = h√ (frequency represents a wave) mc 2 = h√ √ = c/λ mc 2 = 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)  m l – 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 1 2 3 4 5

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 Energy level s p d f sublevels

46. Lesson 1.6. Quantum Numbers _________________________________________________________________________________________ Values Magnetic QN: “ m l ” = -l…0…+l “m l ” represents the orientation of the orbital in space:  l = 0 m l = 0 (only 1 orientation)  l = 1 m l = -1, 0, +1 (3 orientations x, y, z)  l = 2 m l = -2, -1, 0, +1, +2 (5 orientations)  l = 3 m l = -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 1.5.a. The Atomic Orbital _______________________________________________________________ Lesson 1.6. Quantum Numbers _________________________________________________________________________________________ 2p x n l l mlml

50. Lesson 1.6. Quantum Numbers _________________________________________________________________________________________ The total number of orbitals on a particular energy level equals: # orbitals = n 2 n = 3 n = 2 n = 1 3 2 = 9 orbitals 2 2 = 4 orbitals 1 2 = 1 orbital n = 4 1s 2s 2p 3s 4s 3p 4p 3d 4d 4f 4 2 = 16 orbitals