Chapter 4 Electrons In Atoms

1. Chapter 4 • Electrons • In • Atoms

2. Chapter 4 • Section 1 • New Atomic • Model

3. Objectives • Explain the mathematical relationship among speed, wavelength and frequency of electromagnetic radiation. • Discuss the dual wave-particle nature of light. • Describe the photoelectric effect. • Describe the Bohr model of the atom.

8. Rutherford Model Was an improvement over previous models. Helped to explain the positively charged nucleus. It did not explain where the atom’s negatively charged electrons are located in space around the nucleus.

9. Light and Electrons • To begin to grasp the nature of electrons, examining the nature of light is necessary. • We will begin by first introducing some properties of light. • We will then see how these properties are related to the properties of the electron.

10. Properties of Light Light behaves as waves and has wave-like properties. Electromagnetic Radiation – a form of energy that exhibits wavelike behavior as it travels through space. Kinds of electromagnetic radiation include visible light, X rays, ultraviolet and infrared light, microwaves and radio waves.

11. Properties of Light Electromagnetic Spectrum – includes all forms of electromagnetic radiation.

12. Electromagnetic Spectrum

13. Properties of Light All forms of electromagnetic radiation move at a constant speed. 3.0 x 108 m/s This is considered the speed of light.

14. A significant feature of waves is its repetitive nature. • Waves can be characterized by two features: • Wavelength(l) - the distance between corresponding points on adjacent waves. • The units for wavelength are meters, centimeter and nanometer depending on the form of electromagnetic radiation.

15. Wavelength

16. Frequency (n)– defined as the number of waves that pass a given point in a specific time, usually one second (hertz - Hz).

17. Frequency and wavelength are related by the following equation: • c = ln • c = speed of light • l = wavelength • = frequency

18. c = ln • Because c is the same for all electromagnetic radiation, the product ln • is a constant. • is inversely proportional to n • As the wavelength (l) of light increase, its frequency (n) decreases, and vice versa.

20. The Photoelectric Effect • Photoelectric Effect – refers to the emission of electrons from a metal when light shines on the metal. • When light strikes a metal, no electrons were emitted if the light’s frequency was below a certain minimum. • Wave theory of light predicted any frequency of light could eject an electron.

21. Photoelectric Effect Experiment

22. The Photoelectric Effect • The explanation for the photoelectric effect is attributed to German physicist Max Planck. • Planck proposed that objects emit energy in small, specific amounts called quanta. • Quantum – the minimum quantity of energy that can be lost or gained by an atom. • E = hn

23. The Photoelectric Effect • E = hn • E = energy • n = frequency • h = Planck’s constant – 6.626 x 10-34 J-s

24. The Photoelectric Effect • This energy can also be related to its wavelength by the following equations: • E = hnand c = ln • to get: • E = hc l

25. The Photoelectric Effect Albert Einstein expanded on Planck’s theory by explaining that electromagnetic radiation has a dual wave-particle nature. Light can also be thought of as a stream of particles. Each particle of light carries a quantum of energy.

26. The Photoelectric Effect Einstein called these particles photons. Photon – a particle of electromagnetic radiation having zero mass and carrying a quantum of energy. The energy of a particular photon depends on the frequency of radiation: Ephoton= hn

27. The Photoelectric Effect Summary: Light has both wave properties (l and n) and particle (photons) properties. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing the minimum energy (frequency and wavelength).

28. Atom Line Emission Spectrum Ground State – the lowest energy state of an atom. Excited State – A state in which an atom has a higher energy than it has in its ground state.

29. When an excited atom returns to its ground state, it gives off energy that it gained in the form of electromagnetic radiation. • E2 • Excited state energy • Electromagnetic radiation • Electric current • E1 Ground state energy

30. When an electric current was passed through a tube containing hydrogen gas, a pink glow of light was emitted. • When this pink emitted light was passed through a prism, it was separated into a series of specific wavelengths of visible light. • The bands of light were part of what is known as hydrogen’s line-emission spectrum.

31. Hydrogen Atom Line Emission Spectrum

32. Why has the hydrogen atoms given off only specific wavelengths of light? • Scientists had expected to observe the emission of a continuous range of wavelengths of electromagnetic radiation, that is a continuous spectrum. • Attempts to explain this observation led to a new theory of the atom call Quantum Theory.

33. Whenever an excited hydrogen atom falls back from an excited state to its ground state, it emits a photon of radiation. The energy of this photon is: Ephoton = hn This energy is equal to the difference in energy between the atom’s excited state (E2) and its ground state (E1). E2 – E1 = Ephoton = hn

34. Energy difference between ground and excited state

35. The fact that hydrogen atoms emit only specific wavelengths of light indicated that the energy differences between the atom’s energy states were fixed. • This suggested that the electron of a hydrogen atom exists only in very specific energy states.

36. Bohr Model of the Hydrogen Atom • Niels Bohr, a Danish physicist explained the line spectrum of hydrogen in 1913. • His model combined the concepts of Planck and Einstein. Ephoton = hn • Bohr assumed the atom contained a nucleus and that the electrons circled the nucleus in circular orbits.

37. Bohr Model The three postulates of the Bohr model: The electron in the hydrogen atom may only occupy orbits of certain radii that correspond to certain discrete energies. While an electron is in an allowed energy orbit, it does not radiate energy and it remains in that orbit without crashing into the nucleus.

38. Bohr Model • 3) An electron may move from one energy state to another by absorbing or releasing energy. The energy needed is the difference between one energy level and another and is equal to a photon, • Ephoton = hn

39. Bohr Model

40. Bohr Model • Ephoton= hn • By knowing the wavelengths from the hydrogen atom line emission spectrum, Bohr could solve for the energy of the photon using the above equation. • This energy (Ephoton) represents the difference in energy between the different orbits of the hydrogen atom.

41. Bohr Model While the Bohr model works well for hydrogen, it does have its limitations: It did not work well with atoms with more than one electron. It does not account for electron-electron repulsions. Additional electron-nucleus interactions present problems.

42. Classwork Section Review, page 97 Questions 1-5

43. Lab Demo • light experiments with various gases

44. Homework Page 118 -119 Questions 1, 6, 9, 31, 33

45. Chapter 4 • Section 3 • Electron Configurations

46. Objectives • List the atomic orbitals of an atom. • List the total number of electrons needed to fully occupy each main energy level. • State the Aufbau principle, the Pauli Exclusion principle and Hund’s rule. • Write the electron configuration for any element.

47. Atomic Orbitals • Quantum Mechanical Model • A more complex, highly mathematical model was developed to explain observations of atoms containing more than one electron. • This model works for all the elements and not just for hydrogen as in the Bohr model.

48. Electronic Configuration – describes the arrangement of electrons in an atom. • Because atoms of different elements have different number of electrons, a distinct electron configuration exists for each element.

49. The electrons will assume arrangements that have the lowest possible energies. Ground State Configuration – the lowest energy arrangement of the electrons for each element.

50. Atomic Orbitals Bohr Model – the orbit of the electron was circular around the nucleus. In the quantum mechanical modelthe simple circular orbit was replaced with orbitals (electron clouds) of various shapes in which an electron is likely to be.