Chapter 11 Modern Atomic Theory

1. Chapter 11Modern Atomic Theory

2. Rutherford’s Atom • The concept of a nuclear atom (charged electrons moving around the nucleus) resulted from Ernest Rutherford’s experiments. • Question left unanswered: how are electrons arranged and how do they move?

3. Rutherford’s Atom (cont.)

4. Electromagnetic Radiation • Electromagnetic radiation is given off by atoms when they have been excited by any form of energy, as shown in flame tests. • Flame tests: Sample:Color: LiCl SrCl2 NaCl KCl CaCl2

5. Electromagnetic Waves • Velocity = c = speed of light • 2.997925 x 108 m/s • All types of light energy travel at the same speed. • Amplitude = A = measure of the intensity of the wave, i.e.“brightness”

6. Electromagnetic Waves (cont.) • Wavelength =  = distance between two consecutive peaks or troughs in a wave • Generally measured in nanometers (1 nm = 10-9 m) • Same distance for troughs • Frequency = = the number of waves that pass a point in space in one second • Generally measured in Hertz (Hz), • 1 Hz = 1 wave/sec = 1 sec-1 • c =  x 

7. Types of Electromagnetic Radiation

8. Planck’s Discovery • Showed that for certain applications light energy could be thought of as particles or photons

9. Emission of Energy by Atoms/Atomic Spectra • Atoms that have gained extra energy release that energy in the form of light.

10. Atomic Spectra • Line spectrum: very specific wavelengths of light that atoms give off or gain • Each element has its own line spectrum, which can be used to identify that element. • This is the basis of atomic absorption spectroscopy.

11. Atomic Spectra (cont.) • Hydrogen atoms have several excited state energy levels. • Different colors are produced when the excited atoms return to the ground state. • The line spectrum of hydrogen must be related to energy changes in the atom.

12. Atomic Spectra (cont.) • The atom is quantized, i.e. only certain energies are allowed. Continuous levels Quantized levels

13. Bohr’s Model (Niel’s Bohr 1885-1911) • Energy of the atom is quantized • An atom can only have certain specific energy states called quantum levels or energy levels. • When an atom gains energy, an electron “moves” to a higher quantum level. • When an atom loses energy, the electron “moves” to a lower quantum level. • Lines in a spectrum correspond to the difference in energy between the levels.

14. Bohr’s Model (cont.) • Ground state: minimum energy of an atom • The ground state of hydrogen corresponds to having its one electron in the n=1 level • Excited states: energy levels higher than the ground state

15. Bohr’s Model (cont.) • Distances between energy levels decrease as the energy increases • 1st energy level can hold 2 electrons, the 2nd level 8 electrons, the 3rd 18 electrons, etc.

16. Problems with the Bohr Model • Only explains hydrogen atom spectrum (and other 1-electron systems). • Neglects interactions between electrons. • Assumes circular or elliptical orbits for electrons (which is not true).

17. Wave Mechanical Model of the Atom • Experiments later showed that electrons could be treated as waves: Louis De Broglie • The quantum mechanical model treats electrons as waves and uses wave mathematics to calculate probability densities of finding the electron in a particular region in the atom. • Schrödinger Wave Equation

18. Orbitals and Energy Levels • Solutions to the wave equation give regions in space of high probability for finding the electron. These are called orbitals. • Each principal energy level contains one or more sublevels. Sublevels are made up of orbitals.

19. Orbitals and Energy Levels (cont.)

20. Atomic Sublevels & Orbitals • Each type of sublevel has a different shape each and energy. s orbital shape: p orbital shapes: • Each sublevel contains one or more orbitals.

21. Atomic Sublevels & Orbitals (cont.)

22. Pauli Exclusion Principle • No orbital may have more than 2 electrons. • Electrons in the same orbital must have opposite spins. • s sublevel holds 2 electrons (1 orbital) • p sublevel holds 6 electrons (3 orbitals) • d sublevel holds 10 electrons (5 orbitals) • f sublevel holds 14 electrons (7 orbitals)

23. Sublevels and Orbitals nSublevels : Types of Orbitals (and numbers) 1 1s(1) 2 2s(1) 2p(3) 3 3s(1) 3p(3) 3d(5) 4 4s(1) 4p(3) 4d(5) 4f(7) • For a multiple-electron atom, build-up the energy levels, filling each orbital in succession from lowest to highest. • Degenerate orbitals: orbitals with the same energy • e.g. Each p sublevel has 3 degenerate p orbitals

24. Orbital Filling 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 7s 7p Therefore, the order of filling is:

25. Electron Configurations • For a set of degenerate orbitals, fill each orbital half-way first before pairing the electrons. • Electron configurations show how many electrons are in each sublevel of an atom – describes where the electrons are. - The electron configuration for a ground state Li atom is: - The electron configuration for a ground state N atom is:

26. Electron Configurations (cont) What is the electron configuation of: • potassium • cobalt

27. Electron Configurations (cont.) • Valence shell: highest energy level • Electrons in the valence shell are called valence electrons. • Core electrons: electrons not in the valence shell • Often use symbol of previous noble gas in brackets to represent core electrons: phosphorus: 1s22s22p6 3s23p3 = core valence

28. Electron Configuration and the Periodic Table • Elements in the same column on the periodic table have: • Similar chemical and physical properties • Similar valence shell electron configurations • same numbers of valence electrons • same orbital types, but different energy levels Be: Mg:

29. s1 p1 p2 p3 p4 p5 1 2 3 4 5 6 7 p6 d1 d2d3 d4 d5 d6 d7 d8 d9 d10 f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14 s2

31. Orbital diagrams Orbital diagrams (or box diagrams) show orbitals as boxes grouped by sublevels with arrows representing the electrons. Sulfur has an electron configuration of : The orbital diagram for sulfur is:

32. Atomic Properties and the Periodic Table

33. Metallic Character: Metals • Metals • Malleable & ductile • Shiny, lustrous • Conduct heat and electricity • Form cations in solution • Lose electrons in reactions - oxidized

34. Metallic Character: Metalloids • Metalloids -Also known as semi-metals -Show some metal and some nonmetal properties

35. Metallic Character: Nonmetals • Nonmetals -Brittle in solid state -Form anions and polyatomic anions -Gain electrons in reactions - reduced

36. Metallic Character (cont.) • The ease of losing an electron varies as follows: Cs > Rb > K > Na > Li • The same trend is seen in the Group 2 metals:

37. Trend in Ionization Energy • Minimum energy needed to remove a valence electron from an atom in the gas phase: M(g) → M+(g) + 1e- • The lower the ionization energy, the easier it is to remove the electron. • Metals in general have low ionization energies and nonmetals relatively high ones. • Ionization energy decreases down the group. • Valence electron is farther from the nucleus • Ionization energy increases left to right across the period.

38. Trend in Atomic Size