What gives fireworks their different colors?

1. What gives fireworks their different colors?

2. The Atom and Unanswered Questions Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it. The model doesn’t explain how the electrons were arranged around the nucleus. The model doesn’t explain why negatively charged electrons are not pulled into the positively charged nucleus. Section 5-1

3. The Atom and Unanswered Questions (cont.) In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame. Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms. Section 5-1

4. The Wave Nature of Light Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space. All waves can be described by several characteristics. Section 5-1

5. The Wave Nature of Light (cont.) The wavelength(λ) is the shortest distance between equivalent points on a continuous wave. The frequency(ν) is the number of waves that pass a given point per second. The amplitude is the wave’s height from the origin to a crest. Section 5-1

6. The Wave Nature of Light (cont.) Section 5-1

7. The Wave Nature of Light (cont.) • Wavelengths represented by λ (lambda). • Expressed in meters • Frequency represented by v (nu) • Expressed in waves per second • Hertz (Hz) is SI unit = one wave/second • Speed of light represented by c • This is a constant value in a vacuum Section 5-1

8. The Wave Nature of Light (cont.) The speed of light (3.00  108 m/s) is the product of its wavelength and frequency c = λν. λ and v are inversely proportional Section 5-1

9. The Wave Nature of Light (cont.) Light from a green leaf is found to have a wavelength of 4.90 x 10-7m (or 490 nm or 4,900Å) Note: Å = ångström= 10-10 m. What is the frequency of the light? c = λνwhere c= 300 million meters/second c/λ= ν 3.00*108m/s / 4.90 x 10-7 m = ν ν= 6.12 * 1014 Hz Section 5-1

10. The Wave Nature of Light (cont.) Sunlight contains a continuous range of wavelengths and frequencies. A prism separates sunlight into a continuous spectrum of colors. The electromagnetic spectrum includes all forms of electromagnetic radiation. Section 5-1

11. The Wave Nature of Light (cont.) Section 5-1

12. Visible Spectrum

13. The Particle Nature of Light The wave model of light cannot explain all of light’s characteristics. German physicist Max Planck observed in 1900 that matter can gain or lose energy only in small, specific amounts called quanta. A quantum is the minimum amount of energy that can be gained or lost by an atom. Section 5-1

14. The Particle Nature of Light • Energy of a quantum: • Equantum = hv • Planck’s constant (h) has a value of 6.626  10–34 Joules ● second. • Found that energy can only be emitted or absorbed in whole number multiples of hv(e.g. 3hv or 8hv). Section 5-1

15. The Particle Nature of Light (cont.) Albert Einstein proposed in 1905 that light has a dual nature. A beam of light has wavelike and particlelike properties. Think of light as a stream of tiny packets of energy called photons. A photon is a particle of electromagnetic radiation which carries a quantum of energy, but has no mass. Thus: Ephoton = h Ephotonrepresents energy of lighthis Planck's constant. represents frequency. Section 5-1

16. The Particle Nature of Light (cont.) The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it. Thus, depends on Ephoton of light, not intensity (brightness). Section 5-1

17. The Particle Nature of Light (cont.) • Calculate the quantum of energy associated with purple light of 380 nm. • v = c/λ = 3*108 m/s / 3.80 x 10-7 m = 7.9 x 1014 Hz • Ephoton = h = 6.626  10–34 J*sec x (7.9 x 1014/sec) = 5.2 x 10-19 J

18. Atomic Emission Spectra Since the energy of a beam of light is related to its frequency, we can determine its energy by measuring its frequency or wavelength. Thus the color of light(i.e. wavelength) tells us about its energy level. For example, light in a neon sign is produced when electricity is passed through a tube filled with neon gas and excites the neon atoms. The excitedneon atoms emit light to release energy. Only colors related to the quantum of energy released will be generated. Section 5-1

19. Atomic Emission Spectra (Discrete) Section 5-1

20. Atomic Emission Spectra (cont.) The atomic emission spectrumof an element is the set of wavelengths of the electromagnetic waves emitted by the atoms of the element. Each element’s atomic emission spectrum is unique. Section 5-1

21. Section 5-1 Atomic Emission Spectra (cont.) The Emission Spectra for Calcium (top) and Sodium (bottom).

22. A B C D Assessment What is the smallest amount of energy that can be gained or lost by an atom? A.electromagnetic photon B.beta particle C.quantum D.wave-particle Section 5-1

23. A B C D Assessment What is a particle of electromagnetic radiation with no mass called? A.beta particle B.alpha particle C.quanta D.photon Section 5-1

24. Bohr's Model of the Atom

25. Bohr's Model of the Atom Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum. His new model built a relationship between location of electrons and their energy levels. The lowest allowable energy state of an atom and its electrons is called its ground state. When an atom gains energy, it is said to be in an excited state. Section 5-2

26. Bohr's Model of the Atom (cont.) Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits. Section 5-2

27. Bohr's Model of the Atom (cont.) Each orbit was given a number, called the quantum number. Section 5-2

28. Bohr's Model of the Atom (cont.) Hydrogen’s single electron is in the n = 1 orbit in the ground state. When energy is added, the electron moves to the n = 2 orbit or higher (excited state). When the electron is no longer excited by an outside force, it drops back down to a lower energy level. When it “falls” it emits its extra energy in the form of electromagnetic radiation. Section 5-2

29. Bohr's Model of the Atom (cont.) Section 5-2

30. Bohr's Model of the Atom (cont.) Section 5-2

31. Electron Energy Transformations • Bohr’s atomic model attributes hydrogen’s emission spectrum to electrons dropping from higher-energy to lower-energy orbits closer to the nucleus. ∆E = E higher-energy orbit - E lower-energy orbit = E photon = hν= hc/λ • Bohr determined that the energy of each level was: E = -2.178 x 10-18 J x (Z2 / n2) • Where Z = the nuclear charge (= +1 for Hydrogen) and • n = orbit level 2.178

32. Electron Energy Transformations • If an electron moves from n=6 to n =1 (ground state), energy released would equal: • n=6: E6 = -2.178 x 10-18 J x (12/62) = -6.050 x 10-20 J • n=1: E1 = -2.178 x 10-18 J x (12/12) = -2.178 x 10-18 J • ∆E = energy of final state – energy of initial state • ∆E = (-2.178 x 10-18 J) – (-6.050 x 10-20 J) = = -2.117 x 10-18 J (atom released energy) λ = hc/ ∆E = (6.626x10-34J*s x 3.0x108m/s) / -2.117 x 10-18 J = 9.383 x10-8 m or 93.83nm (ultraviolet light emitted)

33. Bohr's Model of the Atom (cont.) Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines and their chemical reactions. The behavior of electrons is still not fully understood, but substantial evidence shows they do not move around the nucleus in circular orbits like planets. Section 5-2

34. The Quantum Wave Mechanical Model While the Bohr model was not perfect, its idea of quantum energy levels was on the right track. Louis de Broglie (1892–1987) noted that Bohr’s quantized energy orbits had similarities to “standing” waves. He proposed that electron particles also behaved like waves as they moved around the nucleus. Section 5-2

35. The Quantum Wave Mechanical Model (cont.) The next figure illustrates that waves with fixed endpoints, like the vibrations of a guitar string, can only vibrate in certain wavelengths (or else they don’t “fit”). Waves going around a circle can only vibrate in whole-number wavelengths. And only ODD numbers of whole wavelengths can fit around a circle or orbit. Section 5-2

36.  represents wavelengthsh is Planck's constant.m represents mass of the particle.represents velocity. The Quantum Wave Mechanical Model (cont.) The de Broglie equationpredicts that all moving particles, including baseballs and electrons, have wave characteristics. Since electrons can only move in specific wavelengths around the nucleus, and wavelength/frequency determines energy, we now have a mathematical reason for Bohr’s quantum energy levels. Section 5-2

37. The Quantum Wave Mechanical Model (cont.) Heisenberg showed it is impossible to take any measurement of an object without disturbing it. The Heisenberg uncertainty principlestates that it is fundamentally impossible to know precisely both the velocity and position of an electron particle at the same time. The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus. These regions are called “Orbitals”. Section 5-2

38. The Quantum Wave Mechanical Model (cont.) Building on de Broglie’s work, Erwin Schrödinger derived an equation in 1926 that treated electrons as waves. His new model was called the quantum wave mechanical model of the atom. Schrödinger’s wave equation applied equally well to elements other than hydrogen, thus improving on the Bohr Model. Section 5-2

39. The Quantum Wave Mechanical Model (cont.) The Schrodinger wave function defines a three-dimensional region of electron location probability called the atomic orbital. 90% Electron Probability Section 5-2

40. Every Electron needs an Orbital it can call Home • All the electrons for each element must occupy a specific orbital – there can be no overlap. • The Quantum Model describes the different types and energies of orbitals using various: • Sizes • Shapes • Orientations • Spins

41. Atomic Orbitals Principal quantum number (n) indicates the relative size and energy of atomicorbitals. • “n”specifies the atom’s major energy levels, called the principal energy levels. • These are similar to Bohr’s orbit levels. • n: 1-7 represent the seven principal energy levels for electrons of all known elements Section 5-2

42. Atomic Orbitals(cont.) Principal levels have multiple sublevels The number of sublevels = n and relate the shape of the orbital. Section 5-2

43. Atomic Orbitals (cont.) Each energy sublevel relates to orbitals of different shape. Spacial orientation can increase the number of sublevels. Section 5-2

44. Atomic Orbitals (cont.) Each Orbital has a unique set of “Quantum Numbers” defining its energy and position. ‘n’ = Principal Quantum Number (size /energy) ‘l’= Angular Momentum Quantum Number. Shape of orbital. There are “n” shapes from 0 to n-1. Also referred to letters based on early spectral descriptions (0=s,1=p,2=d,3=f). ml = Magnetic Quantum Number. Orientation of orbital. Has values between –l and l . ms = Electron Spin Quantum Number. Electrons can spin in two directions: ½ and -½ or “up” and “down” Section 5-2

46. Quantum Numbers (cont.) • What are the quantum numbers for the hydrogen’s lone electron in its ground state? • n=1, l=0, ml=0, ms=1/2 • Usually written as: (1,0,0,1/2) • Only certain quantum numbers are allowed • If n=1, then can’t have l=3, for example. • Would the following be allowed? • (3,2,-3,1/2) • No: level n=3 can have ml=-2,-1,0,1,2 only • (2,1,-1,-1/2) • Yes Section 5-2

47. Quantum Numbers (cont.) • What is the maximum number of electrons with quantum number of (n=5, l=4)? • If l = 4, then can have 9 orbitals of ml = -4,-3,-2,-1,0,1,2,3,4 and within each orbital, can have two electrons of spin ½ and -½ • Thus there is room for 9x2 = 18 electrons • These 18 electrons will have quantum numbers starting with (5,4) • 5,4,-4, ½ • 5,4,-4, -½ • 5,4,-3, ½ • 5,4,-3, -½ • etc. Section 5-2

48. Section 5-2 Atomic Orbitals Summary • Principal quantum number(n) = sizeand energy of atomic orbitals. • n= 1-7 energy levels or “shells” • Subshells (l)= shape of atomic orbitals • Subshell quantum no.: 0, 1, 2, 3 • Subshell orbital letters: s, p, d, f • Orientation (ml): different directions increases the number of orbitals per level. • Number of orbitals for each subshell: s=1; p=3; d=5; f=7.

49. Assessment Heisenberg proposed the: A. Atomic Emission Spectrum B. Quantum Leap C. Wave theory of light D. Uncertainty Principle • A • B • C • D Section 5-2

50. Assessment Who proposed that particles could also exhibit wavelike behaviors? A.Bohr B.Einstein C.Rutherford D.de Broglie • A • B • C • D Section 5-2