Ch. 5: Electrons in Atoms

1. Ch. 5: Electrons in Atoms 5.1 Models of Atomstimeline, Bohr, Schrodinger, atomic orbitals 5.2 Electron Arrangement in Atomselectron configurations 5.3 Physics and the Quantum Mechanical Modellight, electromagnetic radiation, atomic spectra Vocab: quantum, atomic emission spectrum, ground state, photon

2. Review: Properties of Electric Charge Electric charge may be only (+) or (-). Opposite charges attract each other Like charges repel.

3. Review: Democritus (400 B.C.) and John Dalton, 1803 Dalton “Each element consists of a particular kind of atom. All atoms of a particular element are identical.” His atoms had a definite size…they were just “really small.” Democritus “Everything is made up of a few simple parts called atomos.” Atomos means “uncuttable” in Greek.He envisioned atomos as small, solid particles of many different sizes and shapes.

4. Review: Plum Pudding (Thomson) and Rutherford Model “Atoms contain negative particles called electrons and the rest of the atom is positively charged.” “The mass of the atom is contained in a tiny positive nucleus, which is surrounded by electrons moving at high speeds.”

5. 5.1 5.1 Models of the Atom The scale model shown is a physical model. However, not all models are physical. In fact, several theoretical models of the atom have been developed over the last few hundred years. You will learn about the currently accepted model of how electrons behave in atoms.

6. 5.1 This timeline needs to be in your notes! 5.1 Models of the Atom The Development of Atomic Models The timeline shoes the development of atomic models from 1803 to 1911. DALTON----THOMSON (PLUM PUDDING)—RUTHERFORD (NUCLEUS)

7. 5.1 The Development of Atomic Models The timeline shows the development of atomic models from 1913 to 1932. BOHR (ELECTRONS ORBIT)—DE BROGLIE (WAVE MOVEMENT)—SCHRODINGER (ELECTRON CLOUD)

8. 5.1 The Development of Atomic Models What was inadequate about Rutherford’s atomic model? • Rutherford’s atomic model could not explain the chemical properties of elements. • Rutherford’s atomic model could not explain why objects change color when heated. • Why objects when heated to higher and higher temp.first glow red, thenyellow, then white

9. Niels Bohr, 1913 (1911?) “The electrons circle the nucleus but are restricted to particular orbits, like the planets around the Sun.” *Proved this using calculations of different energies of the atom. 1) An electron is found only in specific circular paths, or orbits, around the nucleus; 2) Each orbit has a fixed energy called energy levels; 3) A quantum of energy is the amount of energy required to move an electron from one energy level to another

10. 5.1 The Bohr Model What was the new proposal in the Bohr model of the atom? • Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus. • Each possible electron orbit in Bohr’s model has a fixed energy. • The fixed energies an electron can have are called energy levels. • A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level.

11. The Bohr Model • Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus • Each orbit has a particular energy level; electrons must gain or lose energy to move from one energy level to the next • A quantum of energy is the amount of energy required to move an electron from one energy level to another energy level

12. 5.1 The Bohr Model • Like the rungs of the strange ladder, the energy levels in an atom are not equally spaced. • The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level.

13. Erwin Schrödinger, 1926 (1922?) “The electrons do not orbit the nucleus. Instead, their position is determined by probability. 95% of the time, they can be found in Bohr’s proposed orbits. But not always!” *Proved this using calculations. The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus (electrons are found in orbitals drawn as an “electron cloud”)

14. 5.1 The Quantum Mechanical Model What does the quantum mechanical model determine about the electrons in an atom? • The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. • Austrian physicist Erwin Schrödinger (1887–1961) used new theoretical calculations and results to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom. • The modern description of the electrons in atoms, the quantum mechanical model, comes from the mathematical solutions to the Schrödinger equation.

15. the Quantum Mechanical Model Mathematically determined by Erwin Schrodinger Determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus An atomic orbital is often thought of as a region of space in which there is a high probability of finding an electron. • Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found. • Different atomic orbitals are denoted by letters. The s orbitals are spherical, and p orbitals are dumbbell-shaped. (s, p, d, and f )

16. 5.1 Atomic Orbitals Four of the five d orbitals have the same shape but different orientations in space. • The numbers and kinds of atomic orbitals depend on the energy sublevel.

17. 5.1 The Quantum Mechanical Model • In the quantum mechanical model, the probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud. The cloud is more dense where the probability of finding the electron is high. The propeller blade has the same probability of being anywhere in the blurry region, but you cannot tell its location at any instant. The electron cloud of an atom can be compared to a spinning airplane propeller.

18. 5.1 Atomic Orbitals • The number of electrons allowed in each of the first four energy levels are shown here.

19. 6.2 Using the periodic table

20. 5.1 Section Quiz. 1. Rutherford's planetary model of the atom could not explain a) any properties of elements. b) the chemical properties of elements. c) the distribution of mass in an atom. d) the distribution of positive and negative charges in an atom.

21. 5.1 Section Quiz. 2. Bohr's model of the atom proposed that electrons are found a) embedded in a sphere of positive charge. b) in fixed positions surrounding the nucleus. c) in circular orbits at fixed distances from the nucleus. d) orbiting the nucleus in a single fixed circular path.

22. 5.1 Section Quiz. 3. What is the lowest-numbered principal energy level in which p orbitals are found? a) 1 b) 2 c) 3 d) 4

23. 5.2 5.2 Electron Arrangement in Atoms If this rock were to tumble over, it would end up at a lower height. It would have less energy than before, but its position would be more stable. You will learn that energy and stability play an important role in determining how electrons are configured in an atom.

24. 5.2 Electron Configurations The ways in which electrons are arranged in various orbitals around the nuclei of atoms are called electron configurations.Three rules—the aufbau principle, the Pauli exclusion principle, and Hund’s rule—tell you how to find the electron configurations of atoms. You will need to know electron configurations of elements with atomic number 1-20: ex: calcium is 1s2 2s2 2p6 3s2 3p6 4s2

25. 5.2 Electron Configurations Aufbau Principle According to the aufbau principle, electrons occupy the orbitals of lowest energy first. In the aufbau diagram below, each box represents an atomic orbital. 1s 2s, 2p 3s, 3p 4s, 3d, 4p 5s, 4d, 5p 6s, 4f, 5d, 6p 7s, 5f, 6d, 7p

26. 5.2 Electron Configurations • Pauli Exclusion Principle According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins; that is, the electron spins must be paired. • Hund’s Rule Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.

27. 5.2 Electron Configurations • According to the aufbau principle, electrons occupy the orbitals of lowest energy first. • According to the Pauli exclusion principle, an atomic orbital may describe at most two electrons. To occupy the same orbital, two electrons must have opposite spins • Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible.

28. Write the electron configuration and orbital filling for a. Li b. Mg c. Si Li: atomic number 3 1s22s1 Electron Configuration and Orbital Filling Practice b. Mg: atomic number 12 1s22s22p63s2 1s22s22p63s23p2 • Si: atomic number 14

29. Electron Configurations in Groups Group 1A elements -- there is only one electron in the highest occupied energy level. • Group 4A elements -- there are four electrons in the highest occupied energy level. noble gases are the elements in Group 8A

30. 5.2 Exceptional Electron Configurations Why do actual electron configurations for some elements differ from those assigned using the aufbau principle? • Some actual electron configurations differ from those assigned using the aufbau principle because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations. • Exceptions to the aufbau principle are due to subtle electron-electron interactions in orbitals with very similar energies. • Copper has an electron configuration that is an exception to the aufbau principle. It is 1s22s22p63s23p64s13d10 Instead of 1s22s22p63s23p64s23d9

31. 5.2 Section Quiz. 1. Identify the element that corresponds to the following electron configuration: 1s22s22p5. a) F b) Cl c) Ne d) O

32. 5.2 Section Quiz. 2. Write the electron configuration for the atom N. a) 1s22s22p5 b) 1s22s22p3 c) 1s22s1p2 d) 1s22s22p1

33. 5.2 Section Quiz. 3. The electron configurations for some elements differ from those predicted by the aufbau principle because the a) the lowest energy level is completely filled. b) none of the energy levels are completely filled. c) half-filled sublevels are less stable than filled energy levels. d) half-filled sublevels are more stable than some other arrangements. `

34. 5.3 5.3 Physics and the Quantum Mechanical Model Neon advertising signs are formed from glass tubes bent in various shapes. An electric current passing through the gas in each glass tube makes the gas glow with its own characteristic color. You will learn why each gas glows with a specific color of light.

35. 5.3 Light • How are the wavelength and frequency of light related? • The amplitude of a wave is the wave’s height from zero to the crest. • The wavelength, represented by  (the Greek letter lambda), is the distance between the crests. • The frequency, represented by  (the Greek letter nu), is the number of wave cycles to pass a given point per unit of time. • The SI unit of cycles per second is called a hertz (Hz). • The wavelength and frequency of light are inversely proportional to each other.

36. 5.3 Light The product of the frequency and wavelength always equals a constant (c), the speed of light.

37. 5.3 Light • According to the wave model, light consists of electromagnetic waves. • Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma rays. • All electromagnetic waves travel in a vacuum at a speed of 2.998  108 m/s. • Sunlight consists of light with a continuous range of wavelengths and frequencies. • When sunlight passes through a prism, the different frequencies separate into a spectrum of colors. • In the visible spectrum, red light has the longest wavelength and the lowest frequency.

38. 5.3 Light Sunlight consists of light with a continuous range of wavelengths and frequencies. • When sunlight passes through a prism, the different frequencies separate into a spectrum of colors. • In the visible spectrum, red light has the longest wavelength and the lowest frequency. • ROYGBIV (red, orange, yellow, green, blue, indigo, violet)

39. 5.3 Atomic Spectra A prism separates light into the colors it contains. When white light passes through a prism, it produces a rainbow of colors. When atoms absorb energy, electrons move into higher energy levels. The atom becomes unstable so the electrons then lose energy by emitting light, or photons, when they return to lower energy levels.

40. Atomic Spectra The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element. When light from a helium lamp passes through a prism, discrete lines are produced. Mercury Nitrogen The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron.

41. An Explanation of Atomic Spectra In the Bohr model, the lone electron in the hydrogen atom can have only certain specific energies. • When the electron has its lowest possible energy, the atom is in its ground state. • Excitation of the electron by absorbing energy raises the atom from the ground state to an excited state. • A quantum of energy in the form of light (photon) is emitted when the electron drops back to a lower energy level.

42. 5.3 Quantum Mechanics How does quantum mechanics differ from classical mechanics? • CLASSICAL: In 1905, Albert Einstein successfully explained experimental data by proposing that light could be described as quanta of energy. • The quanta behave as if they were particles. • Light quanta are called photons. • Ex. Electron absorbing energy from ground state and then releasing quanta of energy to return to ground state • QUANTUM: In 1924, De Broglie developed an equation that predicts that all moving objects have wavelike behavior.

43. 5.3 Quantum Mechanics Today, the wavelike properties of beams of electrons are useful in magnifying objects. The electrons in an electron microscope have much smaller wavelengths than visible light. This allows a much clearer enlarged image of a very small object, such as this mite. • Classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves.

44. 5.3 Quantum Mechanics The Heisenberg uncertainty principle states that it is impossible to know exactly both the velocity and the position of a particle at the same time. Once you’ve located the particle, the particle has moved. • This limitation is critical in dealing with small particles such as electrons. • This limitation does not matter for ordinary-sized object such as cars or airplanes. Therefore… classical mechanics adequately describes the motions of bodies much larger than atoms, while quantum mechanics describes the motions of subatomic particles and atoms as waves.

45. 5.3 Section Quiz. 1. The lines in the emission spectrum for an element are caused by a) the movement of electrons from lower to higher energy levels. b) the movement of electrons from higher to lower energy levels. c) the electron configuration in the ground state. d) the electron configuration of an atom.

46. 5.3 Section Quiz. 2. Spectral lines in a series become closer together as n increases because the a) energy levels have similar values. b) energy levels become farther apart. c) atom is approaching ground state. d) electrons are being emitted at a slower rate.

47. What you need to know: Chapter 5 • Atomic people— Democritus, Dalton, Thomson, Rutherford, Bohr, Schrodinger, Einstein/de Broglie • Atomic orbitals (Schrodinger)—energy levels, orbital shapes and location • Electron configurations and orbital filling diagrams (Aufbau, Pauli exclusion, Hund) • Electromagnetic radiation (visible spectrum, ROYGBIV, size of wavelength vs. energy, emission spectra • *concept of atoms absorbing a quantum of energy, moving to higher energy level, and emitting photon upon returning to lower energy level, or ground state; see the photon in terms of color or light emission; quanta of energy obtained by high temperature or high voltage • Trial by Fire and Emission Spectra Labs—light “seen” by heat and high voltage • Vocabulary—quantum mechanical model, atomic orbitals, electromagnetic radiation, atomic emission spectrum, quanta (quantum), ground state, photon, classical mechanics, energy levels, electron configuration, orbital filling,

48. Matter Anti-Matter *Paul Dirac & Carl Anderson, 1928-30 Extra stuff: something to think about… • Dirac postulated that a “positively charged electron” should exist. • Anderson discovers a particle carrying a positive charge of the same magnitude as an electron. He names it the “positron”. (He also wanted to rename the electron the “negatron”.

49. *What is Antimatter? • For some reason (scientists are still working on why), the universe is made of what we have called “matter.” The rules of matter include: • Electrons have a negative charge. Protons have a positive charge. • The universe also contains what we have called “antimatter.” The rules of antimatter include: • Electrons have a positive charge. Protons have a negative charge. • Antimatter is very rare: • Positrons are produced in a certain type of radioactive decay. • Whenever a antimatter particle collides with its matter particle “partner”, they annihilate, releasing energy. • Other anti-matter particles have been found, including the anti-proton and anti-neutron.

50. Murray Gell-Mann & George Zweig, 1964 • “Protons & neutrons are made up of smaller particles called quarks.” (*“Quark” comes from a James Joyce poem.) • There are 6 “flavors” of quarks now known: up, down, top, bottom, strange, and charm. • These were discovered using particle accelerators: colliding particles at high speeds and studying the “shrapnel” of matter and energy.