Arrangement of Electrons in Atoms

1. Arrangement of Electrons in Atoms The Development of a New Atomic Model

2. The Development of the Atom • Dalton’s model • Characteristics • Dalton’s atomic theory • Modifications to Dalton’s atomic theory • Thomson’s model • Characteristics • Grapes in jello • Plum pudding model

3. The Development of the Atom • Rutherford’s model • Gold foil experiment • Characteristics • Problems • What was the exact location of electrons? • Why did electrons not fall into the nucleus?

4. Properties of Light • The study of light revealed the relationship between it and an atom’s electrons. • Light is a form of electromagnetic radiation(EMR). • It is a form of energy. • Other forms of EMR are gamma rays, X-rays, microwaves, ultraviolet rays, infrared rays, and radio waves. • Speed is 3 x 108 m/s(c) in a vacuum. • It has wave-like features.

5. The Electromagnetic Spectrum

6. Wave Features • Repetitive nature. • Wavelength – distance from crest to crest(m, cm, nm)(l) • Frequency – how many crests pass a point in one second(cps, Hz)(f or n) • 1cps = 1Hertz • Crest - peak • Trough – valley

7. Wavelength and Frequency

8. The Photoelectric Effect • The photoelectric effect refers to the emission of electrons from a metal when light shines on the metal. • The effect depends on the frequency of the light striking the metal. If the frequency is too low, nothing happens regardless of time of exposure or intensity. A certain minimum frequency was required. • This was a problem because classical wave theory physics predicted that any frequency of light would cause the effect.

9. The Photoelectric Effect

10. The Particle Description of Light • Planck was studying the emission of light by hot objects(black body radiators). • He proposed that such objects emitted energy in small, specific packets of energy(quanta). • His relationship was E = hn or hf, where E is the energy in joules, h is Planck’s constant(6.626x 10-34 Js) and n or f is frequency.

11. The Particle Description of Light • Einstein expanded on Planck’s ideas and introduced the idea that EMR had a dual wave-particle nature. • According to him, light could travel in both wave motion and particle motion. • He called particles of light photons. Each photon carried a quantum of energy.

12. Einstein explains the photoelectric effect • Einstein proposed that EMR is absorbed by matter only in whole numbers of photons. In order for an electron to be ejected from a metal surface,the electron must be struck by a single photon possessing the minimum energy required to knock the electron loose. This energy corresponds to a threshhold frequency. If the frequency is too low, nothing happens. If the frequency is too high, more energetic electrons are emitted. If the intensity of the light is increased at the threshhold frequency, more electrons are emitted.

13. The Hydrogen Atom Line Emission Spectrum • When current is passed through a gas at low pressure, the potential energy of some gas atoms increases. • The lowest energy state of an electron is the ground state. • A state in which an atom has a higher potential energy than it has in its ground state is an excited state. • When an electron returns to the ground state, it emits absorbed energy in the form of light. When this light was passed through a prism, a bright line emission spectrum was received.

14. Hydrogen’s Line Emission Spectrum(Balmer series lines) In other areas of the EMR spectrum, the Lyman series lines(UV) and Paschen series lines(IR) are seen in the spectrum.

15. Hydrogen Atom Line Emission Spectrum • Classical physics predicted that the hydrogen atoms would be excited by any amount of energy and that their electrons could be in any energy state. • If that was true, then a continuous spectrum(rainbow) would be seen. • As shown, only specific frequencies of light are given off. • This was even the case in non-visible spectra from the infrared and ultraviolet regions of the EMR spectrum.

16. Bohr’s model of the hydrogen atom • In 1913, Bohr proposed a model of the hydrogen atom linking the atom’s electron with photon emission. • The electron can circle the nucleus in allowed orbits only. It has fixed energy. • It has lowest energy closest to the nucleus. As it moves away from the nucleus, it gets higher in energy. • There is a common analogy between Bohr’s atom and the rungs of a ladder.

17. What really happens when energy is absorbed and re-emitted?

18. The Quantum Model of the Atom • New developments – De Broglie • De Broglie wavelength l = h/mv • Heisenberg Uncertainty Principle • It is impossible to know both the location and the velocity of an electron or any other particle. • Schrodinger’s model • Quantum mechanical model • Considers probability of finding electron in a position. • Orbitals are what electrons travel in. • Energy sublevels correspond to orbitals. • Describes mathematically the wave properties of electrons and other very small particles.

19. Atomic Orbitals and Quantum Numbers • Energy levels designated by the principal quantum number, n, where n may be 1, 2, 3, 4, ……… • Energy sublevels correspond to different cloud shapes or regions of high electron probability. These are called atomic orbitals. They are designated by the azimuthal (or angular momentum) quantum number, l, where l may be 0, 1, 2, 3,…. n-1. • Example: If n = 3, then l may be 0, 1, 2.

20. Atomic Orbitals The “s” orbital is spherical. All main levels have an “s” orbital. “p” orbitals first appear in the 2nd main level. There are 3 of them, px, py, and pz. They are perpendicular to one another.

21. Atomic Orbitals “d” orbitals first appear in the 3rd main level. There are 5 of them, dxz, dyz, dxy, dx2-y2, and dz2.

22. Atomic Orbitals “f” orbitals first appear in the 4th main level. There are 7 of them. They are much more complex than the previous three types of orbitals.

23. Atomic Orbitals • The orbital is designated by the magnetic quantum number, ml, which describes the orientation around the nucleus. It has values from –l to 0 to +l. • Example: If l = 2, then ml can have the values -2, -1, 0, 1, 2. • There may be only 2 electrons in an orbital and they must be spinning in opposite directions according to Pauli Exclusion Principle. Thus the spin quantum number, ms, indicates this with values of –1/2 or +1/2.

24. Example question about quantum numbers • What is the set of quantum numbers associated with a 3d electron? • n= 3 • l = 2 • ml = -2, -1, 0, 1, 2 • ms = -1/2, +1/2

25. Important Points about Atomic Orbitals • The lowest energy level(closest to the nucleus) has only 1 sublevel and in that sublevel(“s”) is only one orbital. • The second energy level has 2 sublevels(“s” and “p”), the first of which has one orbital and the second has 3 degenerate orbitals. • The third energy level has 3 sublevels(“s”, “p”, and “d”) having 1, 3, and 5 orbitals respectively.

26. Important Points about Atomic Orbitals • Remember that each atomic orbital can have only 2 electrons. • The maximum number of electrons in any main energy level is 2n2. • Example: If n = 3, then the maximum number of electrons that can be placed there is 2(32) = 18. • The maximum number of orbitals in a given main energy level is n2.

27. Summary