1. COURSE NAME: CHEMISTRY 101 COURSE CODE: 402101-4 Chapter 6 Electronic structure of atoms

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3. 2 Dalton’s Atomic Theory Based on this work and his own experimentations, John Dalton developed an atomic theory:John Dalton 1. Each element is made up of tiny particles called atoms. 2. The atoms of a given elements are identical; atoms of the different elements are different. 3. Chemical compounds are formed when atoms combine with each other; a given compound always has the same relative numbers and types of atoms. 4. Chemical reactions involve reorganization of the atoms, changes in way they are bound together. The atoms themselves are not changed in a chemical reaction.

4. 3 Thomson Atomic Model (1903) Thomson discovered the electron through his work on cathode rays.

5. 5 Manometers Used to Measure Gas Pressures 4 Thomson postulated that: The –ve particles (electrons) were distributed in a uniform sea of positive charge; this was the plum pudding model as the electrons were embedded in the positive charge like plums in a plum pudding.plum pudding model The Plum Pudding Model of the Atom

6. 5 Rutherford Atomic Model (1911) To test Thomson model, he made the gold foil experiment. According to Thomson's model the mass of the atom was spread out throughout the atom. Then, if he shot high velocity alpha particles (helium nuclei) at an atom then the alpha particles will be deflected very little. Rutherford's gold foil experiment

7. 6 The Bohr's Theory Just like Rutherford he assumed that electrons rotate around the nucleus, but with the consideration of some ideas.Rutherford 1. An atom consists of a small, heavily positively charged nucleus around which electrons revolve in definite circular paths called orbits. 2. These orbits are associated with definite energies. These orbits are stable and called "stationary" orbits. 3. As long as the electron remains in a particular orbit its energy remains constant. 4. Each emission or absorption of radiation energy represents the electron transition from one stationary orbit to another.

8. 7 Lyman series n1 = 1 n2 = 2, 3, 4, 5, 6, 7… Balmer series n1 = 2 n2= 3, 4, 5, 6, 7…… Paschen series n1 = 3 n2= 4, 5, 6, 7…… Brackett series n1 = 4 n2= 5, 6, …… Pfund series n1 = 5 n2= 6, 7, ……

9. 9 8 Atomic Structure

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11. 11 10 Atomic Structure

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13. Quantum Numbers Describe the properties of the atomic orbitals and the electrons 1-principle quantum number 2-Secondary quantum number 3-Magnetic quantum number 4-Spin quantum number 12

14. Principle quantum number (n) - It is integral values: n = 1,2,3,4,…∞ -The average distance of the electron from the nucleus. - Greater n is greater the average distance, greater energy, larger the orbital. 13

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17. 16 Magnetic quantum number (m l ) Describes the orientation of the orbital in space

18. 17 Spin quantum number (m s ) Defines the spin of electron about its own axis

19. 18 Table of Allowed Quantum Numbers

20. 19 Rules Governing the Allowed Combinations of Quantum Numbers The three quantum numbers (n, l, and m) that describe an orbital are integers: 0, 1, 2, 3, and so on. The principal quantum number (n) cannot be zero. The allowed values of n are therefore 1, 2, 3, 4, and so on. The angular quantum number (l) can be any integer between 0 and n - 1. If n = 3, for example, l can be either 0, 1, or 2. For any n there are n values of l (the orbits or sub-shells). The magnetic quantum number (m) can be any integer between -l and +l. If l = 2, m can be either -2, -1, 0, +1, or +2. For any l there are (2l+1) values of m (orbitals). The spin quantum number (s) takes only two values which are +1/2 or -1/2, and any orbital can be occupied by maximum of two electrons of opposite spin.

21. 20 Build-up of the elements and electronic configuration Aufbau Principle The electrons occupy the lowest possible energy levels. In other words, orbitals are filled so that those of the lowest energy are filled first.

22. 21 Sequence of the energy Levels 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p and so on.

23. 23 22 Pauli Exclusion Principle According to Pauli Exclusion Principle, each orbital can accommodate a maximum of two electrons provided they have opposite spins. Hund’s Rule According to Hund’s rule, when we come to orbitals of equal energy such as the three p orbitals, we add single electron to each orbital with spins unpaired until each orbital contains one electron. Then, we make pairing.

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26. 1. How many protons and neutrons are in sulfur-33? 2 protons, 16 neutrons 16 protons, 31 neutrons 16 protons, 17 neutrons 15 protons, 16 neutrons 2. Elements differ from one another in their properties because they have different numbers of protons. electrons. neutrons. protons and neutrons.. 3. The element X has three isotopes: X-221, X-220, and X-218, with masses 220.9, 220.0, and 218.1 amu. The relative abundances of these three isotopes are 74.22%, 12.78%, and 13.00%, respectively. Calculate the atomic mass of element X. -219.7 amu -220.5 amu -220.42 amu -220.4 amu 26 25

27. 4. Dalton's atomic theory was an improvement over the ideas of Democritus because it -did not explain the chemical behavior of atoms. -explained how subatomic particles were distributed in an atom. -explained why atoms are electrically neutral. -was based on experimental evidence. 5. Which of the following statements about subatomic particles is correct? -The subatomic particle with the greatest mass is the electron. -A proton has a mass nearly equal to that of a neutron. -Protons, neutrons, and electrons have the same mass. -The subatomic particle with the least mass is the proton. 27 26

28. 6. According to the Rutherford atomic model, -positive charge is evenly spread throughout the atom. -the electrons occupy almost all the volume of an atom. -the nucleus is mostly empty space. -the neutrons are distributed around the nucleus. 7. An element has 22 protons, 20 electrons, and 26 neutrons. What isotope is it? -superscript 48 subscript 22 Ti -superscript 48 subscript 26 Fe -superscript 48 subscript 22 Fe -superscript 48 subscript 26 Ti 28 27

29. 8. Which of the following ideas is NOT part of Dalton's atomic theory? -All elements are composed of tiny particles called atoms. -Atoms of the same element are identical. -In chemical reactions, atoms of one element are changed into atoms of another element. -Atoms of different elements can physically mix together or chemically combine. 9. The relative charges of an electron, neutron, and proton are, respectively, −1, +1, and 0. +1, 0, and −1. −1, 0, and +1. 0, −1, and +1. 10. Which of the following is NOT an example of a subatomic particle? -protons. -molecules. -electrons. -neutrons 29 28

30. 11. What neutral atom is represented by the following configuration: 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p6, 5s2,4d10,5p5? Bromine Iodine Lead Zinc 12. What neutral atom is represented by the following configuration: 1s2,2s2,2p6,3s2,3p6,4s2,3d10,4p4? (1 point) Silicon Selenium Sulfur Silver 13. Which Aufbau diagram below correctly represents a Carbon atom in its ground state Third one First one Second one Third one First one Second one 30 29

31. 31 14. Which is the correct order for the Aufbau principle? 1s2 2s2 3s2 4s2 5s2 6s2 7s2 2p6 3p6 4p6 5p6 6p6 7p6 3d10 4d10 5d10 6d10 4f14 5f14 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 7s2 7p6 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6 15. The "up" and "down" arrows in electron orbital notation depict protons & neutrons electrons with opposite charges electrons with opposite spins protons & electrons 16. Which atomic sub-level will follow the 5d? 5f 6p 4f 6s 31 30

32. 32 17. Which of the following noble gas configurations is correct for chlorine? (Ar)2s2,2p5 (Ar)3s2,3p5 (Ne)2s2,2p5 (Ne)3s2,3p5 18. Which of the following states that electrons must fill up each available orbit before doubling up. Hund's Rule Pauli's Exclusion Heisenbery Uncertainity Aufbau Principle Atomic Theory 19. Which principle states that an electron must occupy n=1 if it is available before entering n=2 Hund's Rule Pauli Exclusion Principle Aufbau Principle Heisenberg Uncertainty Principle 32 31

33. 20. Which of the following is not a possible value of ml for an electron with l = 2? A)-1 B)0 C)+1 D)+2 E)+3 21. What is the total number of orbitals associated with the principle quantum number n = 2? A)1 B)2 C)3 D)4 E)None of the above. 22. The set of quantum numbers that correctly describes an electron in a 3p orbital is: A)n = 3; l = 0; ml = 0; ms = 0 B)n = 3; l = 2; ml = -2, -1, 0, 1, or 2; ms = +1/2 or -1/2 C)n = 3; l = 1; ml = -1, 0, or 1; ms = +1/2 or - 1/2 D)n = 4; l = 0; ml = -1, 0, or 1; ms = +1/2 or -1/2 E)None of the above. 33 32

34. 23. What is the maximum number of electrons that can be accommodated in the shell with n = 4? A)32 B)18 C)24 D)10 E)None of the above. 24. The electronic configuration and filling order of the element whose atomic number is 26 is: A);1s22s22p63s23p64s03d8 B)1s22s22p63s23p63d64s2 C)1s22s22p63s23p64s23d6 D)1s22s22p63s23p64s23d44p2 E)None of the above. 25. Using the noble gas core designation, which of the configurations below correctly describes the ground state electron configuration of Cu? A)[Ne]4s23d9 B)[Ar]4s23d9 C)[Kr]4s13d10 D)[Ar]4s13d10 E)None of the above. 34 33