Electronic Structure and Periodic Properties

1. Electronic Structure and Periodic Properties Wave Nature of Light Models of the Atom Bohr Model Quantum Mechanical Model Atomic Orbitals Electron Configurations Periodic Properties of Elements

2. Electronic Structure of Atoms--Introduction • Elements in the same group exhibit similar chemical and physical properties. • Alkali Metals: • soft • very reactive • metal • Noble Gases • gases • inert (unreactive) • Why???

3. Electronic Structure of Atoms--Introduction • When atoms react, their electrons interact. • The properties of elements depend on their electronic structure. • the arrangement of electrons in an atom • number of electrons • distribution of electrons around the atom • energies of the electrons

4. Electronic Structure of Atoms--Introduction • Understanding the nature of electrons and the electronic structure of atoms is the key to understanding the reactivity of elements and the reactions they undergo. • Much of our knowledge of the electronic structure of atoms came from studying the ways elements absorb or emit light.

5. The Wave Nature of Light • Light is a type ofelectromagnetic radiation • a form of energy with both electrical and magnetic components Wavelength (l) the distance between successive peaks Frequency (u) the number of complete wavelengths that pass a given point in 1 sec l x u = c c = 3.00 x 108 m/s (speed of light)

6. The Wave Nature of Light The electromagnetic spectrum:

7. The Wave Nature of Light • Different types of electromagnetic radiation have different properties because they have different u and l. • Gamma rays • wavelength similar to diameter of atomic nuclei • Hazardous • Radio waves • wavelength can be longer than a football field

8. Quantized Energy and Photons • Classical physics (mechanics) suggests that both electromagnetic radiation and matter can have any energy: A car rolling down a hill can have any potential energy (energy of position) depending on its position on the hill.

9. Quantized Energy and Photons • Classical mechanics is not correct, however. • Max Planck suggested the idea that energy is transferred in “packets” called quanta(plural). • Quantum:the smallest quantity of energy that can be emitted or absorbed as electromagnetic energy

10. Quantized Energy and Photons • Planck proposed that the energy of a single quantum is directly proportional to its frequency: E = hu where E = energy u = frequency h = Planck’s constant (6.63x10-34 J-s)

11. Quantized Energy and Photons • According to Planck’s theory,energy is always emitted or absorbed in whole number multiples of hu(i.e hu, 2hu, 3hu) • According to Planck’s theory, the energy levels that are allowed are ‘quantized.’ • restricted to certain quantities or values

12. Quantized Energy and Photons • In order to understand quantized energy levels, compare walking up (or down) a ramp versus walking up (or down) stairs: • Ramp:continuous change in height • Stairs:quantized changed in height • You can only stop on the stairs, not between them

13. Quantized Energy and Photons • If Planck’s quantum theory is correct, why don’t we notice its effects in our daily lives? • Planck’s constant is very small (6.63 x 10-34 J-s). • A quantum of energy (E = hu) is very small. • Gaining or losing such a small amount of energy is: • insignificant on macroscopic objects • very significant on the atomic level

14. Quantized Energy and Photons • In 1905 Einstein used Planck’s quantum theory to explain the photoelectric effect. • Light shining on a clean metal surface causes the surface to emit electrons. • The light must have a minimum frequency in order for electrons to be emitted.

15. Quantized Energy and Photons • Einstein explained these results by assuming that the light striking the metal is a stream of tiny energy packets of radiant energy(photons). • The energy of each photon is proportional to its frequency. E = hu

16. Quantized Energy and Photons • When a photon strikes a metal surface: • Energy is transferred to the electrons in the metal • If the energy is great enough, the electron can overcome the attractive forces holding it to the metal. • Any extra energy above the amount required to “free” the electron simply increases the kinetic energy of the electron.

17. Quantized Energy and Photons • Einstein’s explanation of the photoelectric effect led to a dilemma. • Is light a wave or does it consist of particles? • Currently, light is considered to have both wave-like and particle-like properties. Matter also has this same dual nature.

18. Atomic Models • Two models are used to explain the behavior and reactivity of atoms and ions. • Bohr model • Quantum mechanical model

19. High voltage H2 Bohr Model • Bohr developed an atomic model that explained the line spectrum observed for the hydrogen atom. • When an electrical current is passed thru a sample of H2 (g), energy is transferred to the H2 molecules. • The molecules are broken up. The H atoms absorb energy and “jump” to a higher energy level.

20. The Bohr Model of the Atom The H atoms “relax” back to their original energy level by giving off the absorbed energy as electromagnetic radiation. High voltage H2

21. The Bohr Model of the Atom The light is analyzed in a spectrometer by separating it into its different colors. High voltage H2

22. The Bohr Model of the Atom The separated colors are recorded as spectral lines. High voltage H2 Atomic spectrum

23. The Bohr Model of the Atom • The spectrum of atomic hydrogen consists of a series of discrete lines such as the ones shown previously. • Why would an atom emit only certain frequencies of light and not all of them?

24. The Bohr Model of the Atom According to the Bohr Model of the atom: • Electrons move in circular orbits around the nucleus. • Energy is quantized: • only orbits of certain radii corresponding to certain definite energies are allowed • an electron in a permitted orbit has a specific energy (an “allowed energy state”)

25. The Bohr Model of the Atom • The allowed orbits have specific energiesgiven by the formula: En = (-RH) 1 where n = 1, 2, 3… n2 RH = Rydberg constant = 2.18 x 10-18 J • n is called the principal quantum number

26. The Bohr Model of the Atom • Each orbit in an atom corresponds to a different value of n. • As n increases the radius of the orbit increases (i.e. the orbit and any electrons occupying it are further from the nucleus) • n=1 is the closest to the nucleus • 0.529 Angstroms for the hydrogen atom

27. The Bohr Model of the Atom • The energy of the orbit is lowest for n=1 and increases with increasing n. • Lower energy = more stable • Lower energy = more preferred state

28. The Bohr Model of the Atom • The lowest energy state of an atom is called the ground state. • n = 1 for the electron in a H atom • When an electron has “jumped” to a higher energy orbit (i.e. n = 2, 3, 4…) it is considered to be in an excited state.

29. The Bohr Model of the Atom • To explain the line spectrum for hydrogen, Bohr assumed that an electron can “jump” from one allowed energy state to another. • Energy absorbed e- “jumps” to higher energy state • e- “relaxes” back to a lower energy state energy is emitted

30. n=4 n=3 n=2 energy n=1 The Bohr Model of the Atom

31. The Bohr Model of the Atom • Since the energies of the orbits in an atom are quantized, transitions from one allowed orbit to another involves only specific amounts of energy. DE = Ef - Ei

32. The Bohr Model of the Atom • Since E = hu, the energy of the light emitted can have only specific values. • Therefore the u of the light can have only specific values as well. • So, the line spectrum for each element will be unique and will depend on the “allowed” energy levels in that element.

33. The Bohr Model of the Atom • The Bohr model effectively explains the line spectra of atoms and ions with a single electron • H, He+, Li2+ • Another model is needed to explain the reactivity and behavior of more complex atoms or ions • Quantum mechanical model