Modern Atomic Theory and the Periodic Table

1. Modern Atomic Theory and the Periodic Table

2. Electromagnetic Spectrum Increasing wavelength Increasing energy and increasing frequency

3. Wavelength • Wavelength – lambda, l • Distance from a peak in one wave to a peak in another wave • Measured in nanometers; meters

4. Frequency • Frequency, nu, n • The number of waves that pass a certain point in one second • Measured in 1/s or s-1 or Hertz (Hz)

5. Electromagnetic Radiation • Radiowaves • Microwaves – • Are the right wavelength to excite water molecules (get them moving around = heat). • Ultraviolet light – sunlight • X-rays • Gamma rays

6. Velocity, c • Distance a wave travels in meters in one second. • m/s or ms-1 • Velocity= wavelength x frequency • c = l n • c = speed of light; all electromagnetic radiation travels 3.00 x 108 m/s in a vacuum • Inverse proportionality between lambda and nu; n = 1/l

7. Energy • Energy associated with wavelengths and frequencies can be calculated. • E is energy • c is the speed of light 3.00 x 108 m/s • h is Planck’s constant 6.6256 x 10-34 Js • l is wavelength • n is frequency • E = hn • E = hc/l

8. Continuous spectrum vs. line spectrum • White light – when shown through a prism contains all of the colors in the visible region of the electromagnetic spectrum blending from one color to another. This is called a continuous spectrum. All wavelengths of visible light are present. • Atoms emit light – when excited by high voltage and shown through a prism has distinct colored lines associated with particular atoms.

9. Continuous spectrum Line Spectrum

10. Why line spectra? • Electrons orbit the nucleus in orbits or shells. These shells have different energy associated with them. • An electron can be excited (energized) causing it to jump up to a higher energy level. • It will not stay in this excited state for very long and will relax (fall back) to a lower level . When it relaxes to a lower level it emits energy in the form of light. • This energy emitted is called a photon (an energy packet). • The wavelengths emitted are discreet. Each element has its own unique spectrum.

11. Rainbow – H2Odroplets suspended in the atmosphere – droplets act like a prism which separates colors into all of the colors in the visible spectrum • Neon signs – noble gases – Ne, Ar, Kr

12. Niels Bohr • Simple model to explain line spectrum. • Orbits are specific, fixed distances from the nucleus. • Only worked for Hydrogen • Energy of each orbit is given a quantum number, n. This number was also fixed or quantized. Energy levels similar to a ladder – impossible to be between the steps so discreet energy levels.

17. Hydrogen gas excited by electrical voltageEach color represents a transition from different energy levels thus different wavelengths and different colors of light

18. 410 nm 434 nm 656 nm 486 nm

19. When an electron absorbs energy it is excited and is raised to a higher energy level. It goes from ground state to an excited state. • When an electron relaxes it loses energy and goes to a lower energy level and emits energy (photon) equal to the energy level difference • The energy is light.

20. Spectrum of Iron

21. Quantum Theory • More than just the simple model given to us by Bohr. • Energy sublevels or orbitals associated with each principal energy level (n)

22. Principal Quantum Numbers (n) • Principal energy levels are assigned values called principal quantum numbers • Electrons in the lower levels (n=1) are closer to the nucleus • Electrons in higher levels are farther from the nucleus • Period numbers on the periodic table also relate to principal quantum number

23. Subshells • Shells (principal energy levels) are divided into subshells • These are identified by s, p, f, d … • Number of subshells in an energy level equals principal quantum number • n = 1 has one subshell – s • n = 2 has two subshells – s, p • n = 3 has three subshells – s, p, d • n = 4 has four subshells = s, p, d, f

25. Shape of orbitals • s subshells are spherical • p subshells have 3 orbitals. • They are shaped like a dumbbells and each resides in a different plane, x,y,z • d subshells have 5 orbitals. • Complicated shapes. See page 209.

26. The maximum number of electrons each principal quantum energy level can hold can be determined by 2n2, where n is the principal quantum number • n = 1, one type of orbital - 1s, 2 electrons • 2(1)2 = 2 electrons • n = 2, two types of orbitals – 2s, 2p • 2s holds 2 electrons, three 2p which each hold 2 electrons • 2(2)2 = 8 • n = 3, three types of orbitals, 3s, 3p, 3d • One 3s, three 3p, five 3d; 2(3)2 = 18 electrons • n = 4, four types of orbitals, 4s, 4p, 4d, 4f • One 4s, three 4p, five 4 d, seven 4 f; 32 electrons

27. Energy levels, sublevels and orbitals • Electrons reside in shells (energy levels) • Shells are assigned number = principal quantum number (n) • Energy of electrons increase as n increases • Energy levels correspond to the principal quantum number • Sublevels are the types of orbitals in each energy level. Assigned s, p, d, f etc. • Orbitals are where electrons reside. No more than 2 electrons can reside in one orbital • Electrons enter the lowest energy orbitals first –s<p<d<f

28. Pauli exclusion principle • Each principal energy level has types of orbitals (s, p, d, f) associated with it. • Each orbital can hold 2 electrons. • Electrons appear to spin on an axis. Each electron can spin in only two directions. • Denoted by • An atomic orbital can only hold two electrons with opposite spins.

29. Electron Configuration, Orbital Filling,Valence electrons

30. Electrons are indicated by arrows: ↑ or ↓. • Each arrow direction represents one of the two possible electron spin states.

31. Electron Configuration • The electron configuration of an atom indicates which orbitals hold electrons and how many are present. • The outermost shell (energy level) of an atom holds what are known as valence shell. It is the electrons in this outmost shell that will take part in bonding (ionic or covalent) between two atoms of the representative elements. • The electrons in the valence shell are called valence electrons.

32. Things to remember • 1. Looking at neutral atoms for now • 2. No more than 2 electrons can occupy one orbital • 3. Electrons occupy the lowest energy level orbital available. s<p<d<f • 4. Each orbital in a sublevel must be occupied by a single electron before a second can enter

34. He Helium has two electrons. Both helium electrons occupy the 1s orbital with opposite spins. ↑ H 1s1 Hydrogen has 1 electron. It will occupy the orbital of lowest energy which is the 1s. ↑ ↓ 1s2

35. Filling the 2s Sublevel

36. ↑ ↓ ↑ ↓ Be 1s 2s The 2s orbital fills upon the addition of beryllium’s third and fourth electrons. ↑ Li 1s22s1 1s 2s The 1s orbital is filled. Lithium’s third electron will enter the 2s orbital. ↑ ↓ 1s22s2

37. Filling the 2p Sublevel

38. ↑ ↓ ↑ ↓ C 1s 2s 2p The second p electron of carbon enters a different p orbital than the first p electron so as to give carbon the lowest possible energy. ↑ ↓ ↑ ↓ N 1s 2s 2p The third p electron of nitrogen enters a different p orbital than its first two p electrons to give nitrogen the lowest possible energy. ↑ ↓ ↑ ↑ ↓ B 1s22s22p1 1s 2s 2p Boron has the first p electron. The three 2p orbitals have the same energy. It does not matter which orbital fills first. ↑ ↑ 1s22s22p2 ↑ ↑ ↑ 1s22s22p3

39. ↑ ↓ ↑ ↓ O 1s 2s 2p There are four electrons in the 2p sublevel of oxygen. One of the 2p orbitals is now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↓ F 1s 2s 2p There are five electrons in the 2p sublevel of fluorine. Two of the 2p orbitals are now occupied by a second electron, which has a spin opposite to that of the first electron already in the orbital. ↑ ↓ ↑ ↑ 1s22s22p4 ↑ ↓ ↑ ↓ ↑ 1s22s22p5

40. ↑ ↓ ↑ ↓ Ne 1s 2s 2p There are 6 electrons in the 2p sublevel of neon, which fills the sublevel. ↑ ↓ ↑ ↓ ↑ ↓ 1s22s22p6

41. Write the Electron configuration for the following elements. • B • N • Na • Cl

42. Electron Structure and the Periodic Table • Relate electron configuration to the periodic table.

43. Br, Rb, C, Ar

44. Shorthand notation for electron configuration • Use the noble gas symbol in front of the element. • Br • Kr • Al

45. Periodic trends • Families or groups all have similar chemical properties • Note: all elements in a group or family will have the same number of valence electrons in its outer shell • The number of valence electrons in the representative elements is the same as the Roman numeral on top of the column.