Understanding Electromagnetic Radiation and Particle Properties of Light

1. Chapter 9 Electrons in Atoms and the Periodic Table Helium gas He has 2 protons and if neutral how many electrons? 2006, Prentice Hall

2. CHAPTER OUTLINE

3. Blimps • blimps float because they are filled with a gas that is less dense than the surrounding air • early blimps used the gas hydrogen, however hydrogen’s flammability lead to the Hindenburg disaster • blimps now use helium gas, which is not flammable

4. Why is hydrogen gas diatomic but helium gas not? Because hydrogen is reactive and helium is inert. What makes hydrogen reactive? Recall that when elements are arranged in order of increasing atomic number (# of protons), certain sets of properties periodically recur -All group I elements have similar reactivity -All noble gases are inert -Alkali metals- +1; metals -Alkaline earth metals- +2; metals -Halogens- -1; nonmetals -Noble gases- 0; nonmetals We need a theory/model that would explain why these occur:

5. Classical View of the Universe • since the time of the ancient Greeks, the stuff of the physical universe has been classified as either matter or energy • we define matter as the stuff of the universe that has mass and volume • therefore energy is the stuff of the universe that does not have mass and volume • we know that matter is ultimately composed of particles, and the properties of particles determine the properties we observe • energy therefore should not be composed of particles, in fact the thing that all energy has in common is that it travels in waves

6. Light: Electromagnetic Radiation • light is a form of energy • light is one type of energy called electromagnetic radiation with electric and magnetic field components and travels in waves

7. Electromagnetic Waves • velocity = c = speed of light • its constant! = 2.997925 x 108 m/s (m•sec-1) in vacuum • all types of light energy travel at the same speed • amplitude = A = measure of the intensity of the wave, “brightness” • height of the wave • wavelength = l = distance between crests • generally measured in nanometers (1 nm = 10-9 m) • same distance for troughs or nodes • determines color • frequency = n= how many peaks pass a point in a second • generally measured in Hertz (Hz), • 1 Hz = 1 wave/sec = 1 sec-1

8. Electromagnetic Waves • Wavelength (λ) is the distance between any 2 successive crests or troughs. Amplitude Nodes

9. inversely proportional WAVES • Frequency (nu,) is the number of waves produced per unit time. • Wavelength and frequency are inversely proportional. • Speed tells how fast waves travel through space. As wavelength of a wave increases its frequency decreases 10.1

10. ELECTROMAGNETICRADIATION • Energy travels through space as electromagnetic radiation. This radiation takes many forms, such as sunlight, microwaves, radio waves, etc. • In a vacuum, all electromagnetic waves travel at the speed of light (3.00 x 108 m/s), and differ from each other in their frequency and wavelength.

11. ELECTROMAGNETICRADIATION • The classification of electromagnetic waves according to their frequency is called electromagnetic spectrum. • These waves range from -rays (short λ, high f) to radio waves (long λ, low f). Long wavelength Low frequency Short wavelength High frequency

12. ELECTROMAGNETICRADIATION Visible light is a small part of the EM spectrum Infrared waves have longer λ but lower  than visible light X-rays have longer λ but lower  than -rays 10.2

13. Particles of Light • Albert Einstein and other scientists in the early 20th century showed that wave properties do not completely explain electromagnetic radiation (EM) and showed that EM was composed of particle-like properties called photons • photons are particles of light energy • each wavelength of light has photons that have a different amount of energy • the longer the wavelength, the lower the energy of the photons

14. DUAL NATUREOF LIGHT • Scientists also have much evidence that light beams act as a stream of tiny particles, called photons. Red light has longer wavelength and less energy than blue light A photon of red light A photon of blue light

15. DUAL NATUREOF LIGHT • Scientists, therefore, use both the wave and particle models for explaining light. This is referred to as the wave-particle nature of light. • Scientists also discovered that when atoms are energized at high temperatures or by high voltage, they can radiate light. Neon lights are an example of this property of atoms.

16. Light’s Relationship to Matter He • Atoms can acquire extra energy, but they must eventually release it • When atoms emit energy, it always is released in the form of light • However, atoms don’t emit all colors, only very specific wavelengths • in fact, the spectrum of wavelengths can be used to identify the element Hg

17. When an atom absorbs energy it reemits it as light emission spectrum of Hydrogen White Light Source Emits at every wavelength (all colors) also called a continuous spectrum H produces its own unique and distinctive emission spectrum

18. Spectra

19. The Bohr Model of the Atom • The Nuclear Model of the atom does not explain how the atom can gain or lose energy • Neils Bohr developed a model of the atom to explain the how the structure of the atom changes when it undergoes energy transitions • Bohr’s major idea was that the energy of the atom was quantized, and that the amount of energy in the atom was related to the electron’s position in the atom • quantized means that the atom could only have very specific amounts of energy 1885 to1962 Nobel Prize in Physics in 1922

20. The Bohr Model of the AtomElectron Orbits • the Bohr Model, electrons travel in orbits around the nucleus • more like shells than planet orbits • the farther the electron is from the nucleus the more energy it has Rank the electrons from highest to lowest energy e- e- e- e- e-

21. The Bohr Model of the AtomOrbits and Energy • each orbit has a specific amount of energy • the energy of each orbit is characterized by an integer - the larger the integer, the more energy an electron in that orbit has and the farther it is from the nucleus • the integer (whole #s), n, is called a quantum number Bohr orbits are like steps in a ladder. It is possible to be on one step or another, but it is impossible to be between steps.

22. BOHR MODELOF ATOM • Neils Bohr, a Danish physicist, studied the hydrogen atom extensively, and developed a model for the atom that was able to explain the line spectrum. • Bohr’s model of the atom consisted of electrons orbiting the nucleus at different distances from the nucleus, called energy levels. • In this model, the electrons could only occupy particular energy levels, and could “jump” to higher levels by absorbing energy.

23. BOHR MODELOF ATOM • The lowest energy level is called ground state, and the higher energy levels are called excited states. • When electrons absorb energy through heating or electricity, they move to higher energy levels and become excited. energy

24. BOHR MODELOF ATOM • When excited electrons return to the ground state, energy is emitted as a photon of light is released. • The color (wavelength) of the light emitted is determined by the difference in energy between the two states (excited and ground). Lower energy transition give off red light Higher energy transition give off blue light

25. The Bohr Model of the AtomGround and Excited States • The lowest amount of energy hydrogen’s one electron can have corresponds to being in the n = 1 orbit –this is the ground state • when the atom gains energy, the electron leaps to a higher energy orbit –this is the excited state • the atom is less stable in an excited state, and so it will release the extra energy to return to the ground state • either all at once or in several steps

26. The Bohr Model of the AtomHydrogen Spectrum • every hydrogen atom has identical orbits, so every hydrogen atom can undergo the same energy transitions • however, since the distances between the orbits in an atom are not all the same, no two leaps in an atom will have the same energy • the closer the orbits are in energy, the lower the energy of the photon emitted • lower energy photon = longer wavelength • therefore we get an emission spectrum that has a lot of lines that are unique to hydrogen

27. The Bohr Model of the AtomHydrogen Spectrum Which e- has longer wavelength and lower energy (red, violet or blue-green)?

28. The Bohr Model of the AtomSuccess and Failure • the mathematics of the Bohr Model very accurately predicts the spectrum of hydrogen • however its mathematics fails when applied to multi-electron atoms • it cannot account for electron-electron interactions • a better theory was needed

29. QUANTUM MECHANICALMODEL OF ATOM • In 1926 Erwin Shrödinger created a mathematical model that showed electrons as both particles and waves. This model was called the quantum mechanical model. • This model predicted electrons to be located in a probability region called orbitals. • An orbital is defined as a region around the nucleus where there is a high probability of finding an electron. 1887 to 1961 Nobel Prize in Physics in 1933

30. Orbits vs. OrbitalsPathways vs. Probability Orbital – acts as a wave and particle thus could end up anywhere in a probability map Orbit – acts as a particle and follows a well-defined path

31. QUANTUM MECHANICALMODEL OF ATOM • Based on this model, there are discrete principal energy levels within the atom. • Principal energy levels are designated by n. • The electrons in an atom can exist in any principal energy level. As n increases, the energy of the electron increases

32. QUANTUM MECHANICALMODEL OF ATOM • Each principal energy level is subdivided into sublevels. • The sublevels are designated by the letterss, p, d and f. • As n increases, the number of sublevels increases. 10.7, 10.8

33. QUANTUM MECHANICALMODEL OF ATOM • Within the sublevels, the electrons are located in orbitals. The orbitals are also designated by the letters s, p, d and f. • The number of orbitals within the sublevels vary with their type. s sublevel = 1 orbital = 2 electrons p sublevel = 3 orbitals = 6 electrons d sublevel = 5 orbitals = 10 electrons f sublevel = 7 orbitals = 14 electrons An orbital can hold a maximum of 2 electrons

34. How does the 1s Subshell Differ from the 2s Subshell

35. Probability Maps & Orbital Shapes Orbitals

36. Probability Maps & Orbital Shapep Orbitals

37. Probability Maps & Orbital Shaped Orbitals

38. ELECTRONCONFIGURATION • The distribution of electrons into the various energy shells and subshells in an atom’s ground state is called its electron configuration • The electrons occupy the orbitals from the lowest energy level to the highest level (Aufbau Principal). • The energy of the orbitals on any level are in the following order: s < p < d < f. • Each orbital on a sublevel must be occupied by a single electron before a second electron enters (Hund’s Rule).

39. ELECTRONCONFIGURATION • Electron configurations can be written as: Number of electrons in orbitals 2 p6 Principal energy level Type of orbital

40. ELECTRONCONFIGURATION • Another notation, called the orbital notation is shown below: Electrons in orbital with opposing spins Principal energy level 1 s Type of orbital

41. Filling an Orbital with Electrons • each orbital may have a maximum of 2 electrons with opposite spins • Pauli Exclusion Principle • electrons spin on an axis • generating their own magnetic field • when two electrons are in the same orbital, they must have opposite spins • so there magnetic fields will cancel

42. 1s 1s ELECTRONCONFIGURATION ↑ H 1s1 Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. ↑ ↓ He 1s2 Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins.

43. ↑ ↑ ↑ ↑ ↓ ↓ ↓ ↓ 2p 2p 1s 1s 2s 2s ELECTRONCONFIGURATION ↑ B 1s22s22p1 Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. ↑ ↓ ↑ C 1s22s22p2 The second p electron of carbon enters a different p orbital than the first p due to Hund’s Rule.

44. ↑ ↑ ↓ ↓ 2p 1s 2s 3s ELECTRONCONFIGURATION ↑ ↓ ↑ ↓ ↑ ↓ Ne 1s22s22p6 The last p electron for neon pairs up with the last lone electron and completely fills the 2nd energy level. ↑ Na 1s22s22p6 3s1 Sodium has 11 electron. The first 10 will occupy the orbitals of energy levels 1 and 2. core electrons valence electron

45. ELECTRON CONFIGURATION • As electrons occupy the 3rd energy level and higher, some anomalies occur in the order of the energy of the orbitals. • Knowledge of these anomalies is important in order to determine the correct electron configuration for the atoms.

46. ELECTRON CONFIG. & PERIODIC TABLE 10.15

47. ELECTRON CONFIG. & PERIODIC TABLE • The horizontal rows in the periodic table are called periods. The period number corresponds to the number of energy levels that are occupied in that atom. • The vertical columns in the periodic table are called groups or families. For the main-group elements, the group number corresponds to the number of electrons in the outermost filled energy level (valence electrons).

48. ELECTRON CONFIG. & PERIODIC TABLE One energy level 4 energy levels 3 energy levels

49. ELECTRON CONFIG. & PERIODIC TABLE 3 valence electrons 1 valence electron 5 valence electrons

50. ELECTRON CONFIG. & PERIODIC TABLE • The valence electrons configuration for the elements in periods 1-3 are shown below. • Note that elements in the same group have similar electron configurations. 10.15