Chapter5Section2Quantumtheoryandtheatom (1).pptx

Chemistry

1. Chapter 5 : Electrons in Atoms Section 2: Quantum theory and the atom Grade 10 Pages 146 - 155

2. Objectives: Compare the Rutherford, Bohr, and quantum models of an atom. Explain how the wavelengths of light emitted by an atom provide information about electron energy levels. List the four quantum numbers, and describe their significance. Write the electron configuration of an atom by using the Pauli exclusion principle and the the aufbau principle.

3. What makes up matter ? An atom is the smallest particle and is the building block of everything in the universe.

4. Who developed the atomic theory? In 1808, John Dalton published an atomic theory , stating that all matter is made up of atoms that cannot be created, divided, or destroyed . This theory also stated that all atoms of a certain element are identical, but they differ from atoms of all other elements. Every substance is made up of atoms combined in certain ways.

5. In 1897, J. J. Thomson’s experiments provided evidence that atoms contain negatively charged particles , which were later called electrons . In 1909, Ernest Rutherford’s experiment suggested that atoms have a nucleus —a small, dense center that has a positive charge. Rutherford later found that the nucleus is made up of smaller, positively charged particles that he called protons . Who developed the atomic theory?

6. Rutherford suggested that electrons, like planets orbiting the sun, revolve around the nucleus in circular or elliptical orbits . Rutherford’s model could not explain why electrons did not crash into the nucleus . The Rutherford model of the atom was replaced only two years later by a model developed by Niels Bohr . Who developed the atomic theory?

7. Who developed the atomic theory? Niels Bohr suggested a model in which electrons move around the nucleus in circular paths, with each path at a certain distance from the nucleus and in specific energy levels .

8. According to Bohr’s model , electrons can be only certain distances from the nucleus. Each distance corresponds to a certain quantity of energy that an electron can have. An electron that is as close to the nucleus as it can be is in its lowest energy level . The farther an electron is from the nucleus, the higher the energy level that the electron occupies . The difference in energy between two energy levels is known as a quantum of energy. Who developed the atomic theory?

9. What is the current atomic theory ? In 1924, Louis de Broglie pointed out that the behavior of electrons according to Bohr’s model was similar to the behavior of waves In 1932, James Chadwick discovered that the nucleus contains uncharged particles called neutrons . Instead, the current theory suggests that electrons move within an area around the nucleus called the electron cloud .

10. The current atomic model In this model, electrons are located in orbitals , regions around a nucleus that correspond to specific energy levels . Orbitals are regions where electrons are likely to be found. Orbitals are sometimes called electron clouds because they do not have sharp boundaries.

11. Electron behavior The present-day model of the atom takes into account both the particle and wave properties of electrons .

12. Electron behavior Each electron move from a particular energy level to a lower energy level will release light of a different wavelength. Light is an electromagnetic wave . Red light has a low frequency and a long wavelength . Violet light has a high frequency and a short wavelength. The frequency and wavelength of a wave are inversely related.

13. Electron behavior

14. Electrons and energy An electron can move from a low energy level to a high energy level by absorbing energy . Electrons at a higher energy level are unstable and can move to a lower energy level by releasing energy . This energy is released as light that has a specific wavelength.

15. Electrons and energy An electron in a state of its lowest possible energy, is in a ground state. The ground state is the lowest energy state of a quantized system If an electron gains energy, it moves to an excited state. An excited state is a state in which an atom has more energy than it does at its ground state. An electron in an excited state will release ( emit ) a specific quantity of energy as it quickly “falls” back to its ground state.

16. Quantum Numbers The present-day model of the atom is also known as the quantum model . To define the region in which electrons can be found, scientists have assigned four quantum numbers that specify the properties of the electrons . A quantum number is a number that specifies the properties of electrons.

17. Principal Quantum Number (n) The principal quantum number , symbolized by n , indicates the main energy level occupied by the electron. Values of n are positive integers , such as 1, 2, 3, and 4. As n increases , the electron’s distance from the nucleus and the electron’s energy increases .

18. The angular momentum quantum number (l) The main energy levels can be divided into sublevels. These sublevels are represented by the angular momentum quantum number, l . which is ( n-1 ) This quantum number indicates the shape or type of orbital that corresponds to a particular sublevel. A letter code is used for this quantum number: Orbitals

19. Magnetic quantum number (m l ) The magnetic quantum number , symbolized by m , is a subset of the l quantum number . It also indicates the numbers and orientations of orbitals around the nucleus . The value of m takes whole-number values, depending on the value of l. l (n-1) n 1 1 2 2 3 3 4 M value (-l  +l)

20. Spin quantum number ( m s ) The spin quantum number , indicates the orientation of an electron’s magnetic field relative to an outside magnetic field. The spin quantum number is represented by: A single orbital can hold a maximum of two electrons , which must have opposite spins .

21. Quantum model

22. Quantum model

23. Orbital Notation

24. Pauli exclusion principle In 1925 the German chemist Wolfgang Pauli established a rule is known as the Pauli exclusion principle. The Pauli exclusion principle states that two particles of a certain class cannot be in the exact same energy state. This means that no two electrons in the same atom can have the same four quantum numbers.

25. Aufbau principle The aufbau principle states that electrons fill orbitals that have the lowest energy first. Aufbau is the German word for “building up.” The smaller the principal quantum number , the lower the energy. Within an energy level, the smaller the l quantum number , the lower the energy. So, the order in which the orbitals are filled matches the order of energies. 1s < 2s < 2p < 3s < 3p

26. Electron Configurations ( Hund’s rule ) Electron orbitals are filled according to Hund’s Rule . Hund’s rule states that orbitals of the same n and l quantum numbers are each occupied by one electron before any pairing occurs. Orbital diagram for sulfur: The arrangement of electrons in an atom is usually shown by writing an electron configuration .

27. 1 2 3 4 5 6 7 6 7 1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A 5A 6A 7A 8A group # = # valence (outside) e- d p f s Row = # shells

28. Electron Configurations

29. Electron Configuration 1 s 1 row # shell # possibilities are 1-7 7 rows subshell possibilities are s, p, d, or f 4 subshells group # # valence e- possibilities are: s: 1 or 2 p: 1-6 d: 1-10 f: 1-14 Total e- should equal Atomic # What element has an electron configuration of 1s 1 ?

30. Order of Electron Subshell Filling: It does not go “in order” 1s 2 2s 2 2p 6 3p 6 4p 6 5p 6 6p 6 7p 6 3s 2 4s 2 5s 2 6s 2 7s 2 3d 10 4d 10 5d 10 6d 10 4f 14 5f 14 1s 2 2s 2 2p 6 3p 6 3s 2 4s 2 4p 6 5s 2 3d 10 5p 6 6s 2 4d 10 6p 6 7s 2 5d 10 4f 14 7p 6 6d 10 5f 14

31. Electron Configurations Based on the quantum model of the atom, the arrangement of the electrons around the nucleus can be shown by the nucleus’s electron configuration. Example: sulfur has sixteen electrons . Its electron configuration is written as: 1s 2 2s 2 2p 6 3s 2 3p 4 Two electrons are in the 1s orbital, two electrons are in the 2s orbital, six electrons are in the 2p orbitals, two electrons are in the 3s orbital, and four electrons are in the 3p orbitals.

32. Electron Configurations (Shorthand Notation) Each element’s configuration builds on the previous elements’ configurations. To save space, one can write this configuration by using a configuration of a noble gas : neon, argon, krypton, and xenon The neon atom’s configuration is 1s 2 2s 2 2p 6 , so the electron configuration of sulfur is: [Ne ] 3s 2 3p 4

33. Sample Problem A: Write the electron configuration for an atom whose atomic number is 20.  atomic number = number of protons = number of electrons = 20 The electron configuration for an atom of this element is written as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 This electron configuration can be abbreviated as follows : [ Ar ]4s 2

34. Sample Problem B: Lithium: find the element on the periodic table what is the period number? how many shells? what is the group number? how many valence electrons? what subshell(s) does Li have? what is the electron configuration? atomic # = 3 2 2 1 1 s 1s 2 2s 1

35. Sample Problem C: Boron: find the element on the periodic table what is the row #? how many shells? what is the group #? how many valence electrons? what subshell(s) does B have? what is the electron configuration? atomic # = 5 2 2 3 3 p 1s 2 2s 2 2p 1

36. Electron Configuration of ions