Unit 2 contd

# Unit 2 contd

## Unit 2 contd

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##### Presentation Transcript

1. Unit 2 contd Chapter 6 and 7.1-7.6

2. Behavior of electrons and atoms Wave Nature of Light Models of the Atom Bohr Model Quantum Mechanical Model Atomic Orbitals Electron Configurations Periodic Properties of Elements

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

4. Electronic Structure of the Atom • 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

5. Electronic Structure of the Atom • 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.  study light!

6. The Wave Nature of Light • So, to understand electronic structure, we must learn about light. • Light is a type ofelectromagnetic radiation • The nature of electromagnetic radiation: wavelike characteristics (like water waves). • Describing waves: • The distance between corresponding points on adjacent waves is the wavelength ().

7. Waves • The nature of electromagnetic radiation: wavelike characteristics (like water waves). Describing waves: • The distance between corresponding points on adjacent waves is the wavelength ().

8. Waves • The number of complete wavelengths passing a given point per unit of time is the frequency(). • For waves traveling at the same velocity, the longer the wavelength, the smaller the frequency.

9. Electromagnetic Radiation • All radiation moves through vacuum at same speed: Speed of light c = 3.00 x 108 m/s c = ln • = wavelength in nm or m n = frequency in ‘per second’ (s-1) High frequency  small/short wavelength Long wavelength  low frequency

10. Think • What is the frequency of green light with a wavelength of 520 nm? c = ln n = c/l • KGOU broadcast at a frequency of 106.3 MHz (megahertz, 1 MHz = 106 s-1). What is the wavelength of this radiation? c = ln l = c/n

11. The electromagnetic spectrum: Wavelengths of g rays: atomic nuclei Wavelengths of radio waves: football field

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

13. The Nature of Energy • Red-hot object is cooler than white-hot object….why? T and l of radiation • Max Planck explained it by assuming that energy comes in packets called quanta. • Quantum:the smallest quantity of energy that can be emitted (released) or absorbed as electromagnetic energy

14. Quantized Energy and Photons • Planck (1900) proposed that the energy of a single quantum is directly proportional to its frequency: E = hn where E = energy n = frequency h = Planck’s constant (6.63x10-34 Joule second or Js)

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

16. 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

17. 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 = hn) is very small. • Gaining or losing such a small amount of energy is: • insignificant on macroscopic objects • very significant on the atomic level

18. 1905: Einstein used Planck’s quantum theory to explain photoelectric effect. “photocells”

19. 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.

20. 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 = hn

21. 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.

22. Think! A laser emits light with a frequency, n, of 4.69 x 1014 s-1. What is the energy of one photon of the radiation from this laser? E = hn h =6.63 x 10-34 Js E = 6.63 x 10-34 Js x 4.69 x 1014 s-1 = 3.11 x 10-19 J

23. More practice What is the energy of one photon of yellow light with a wavelength of 589 nm? E = hn and c = ln n = c/l So… E = h c/l h = 6.63 x 10-34 Js and c = 3.00 x 108 m/s E = 3.37 x 10-19 J

24. 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.

25. Models of Atomic Structure • Scientists initially thought of the atom as a “microscopic solar system.” • electrons orbiting the nucleus • Unit 2 suggested that the atom has a tiny positively charged nucleus with a diffuse “cloud” of electrons surrounding it. • need better understanding of the nature of this “cloud” of electrons.

26. Atomic Models • Two models are used to explain the behavior and reactivity of atoms and ions. Bohr Model And Quantum Mechanical Model

27. The Nature of Energy Another mystery in the early 20th century involved the emission spectra observed from energy emitted by atoms and molecules.

28. The Nature of Energy • For atoms and molecules one does not observe a continuous spectrum, as one gets from a white light source. • Only a line spectrum of discrete wavelengths is observed.

29. 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?

30. The Bohr Model of the AtomNiels Bohr: 1922 Nobel Prize in physics for model of hydrogen atom According to the Bohr Model of the atom: • Electrons move in circular orbitsaround 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”)

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

32. 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 Ångstroms for the hydrogen atom

33. 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

34. 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.

35. 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

36. 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.

37. Bohr Model of the Atom • Niels Bohr (Nobel Prize 1922): • electron in a hydrogen atom • Classical physics do not apply in the atom. • 2. Electrons orbit the nucleus, but only orbits of certain energies are allowed. • 3. Electrons can change from one allowed orbit to another, but that the change will either require energy or produce energy. The energy involved is often in the form of light. n=4 n=3 n=2 Increasing energy n=1 An electron in the lowest energy level (ground state) of H.

38. Sodium Lamp The characteristic yellow light in a sodium lamp is the result of electrons in the high-energy “3p” orbital falling back to the lower-energy “3s” orbital. What does that mean?

39. 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

40. Heisenberg (Nobel Prize 1932) • If electron has properties of both a particle and a wave (dual nature), it is impossible to know the exact position and momentum (mass times speed) simultaneously. Uncertainty principle Bottom line: Electrons don’t move in well-defined circular orbits around the nucleus. Only important for masses as small as an electron!!

41. Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. It is known as quantum mechanics.

42. Quantum Numbers • Solving the Schrödinger wave equation gives a set of “wave functions”, called orbitals, and their corresponding energies. • Each orbital describes a spatial distribution of electron density. • An orbital is described by a set of three quantum numbers. = a way of expressing the probability of finding an electron at a particular location in space

43. Orbitals • An orbital: • describes a specific distribution of electron density in space • has a characteristic energy • has a characteristic shape • is described by three quantum numbers: n, l, ml • can hold a maximum of 2 electrons • Note:A fourth quantum number (ms) is needed to completely describe each electron in an orbital

44. Quantum Numbers • 1) Principal quantum number (n): • integral values n = 1, 2, 3, 4, .. • indicates the average distance of the electron from the nucleus • as n increases, the average distance from the nucleus increases n=4 n=3 n=2 Increasing energy n=1

45. Quantum Numbers 2) Angular momentum quantum number (l) • integral values l = 0, 1, 2,….(n-1) Example: If n = 4, then l = 0, 1, 2, or 3. • defines the shape of the orbital • The value for l from a particular orbital is usually designated by the letters s, p, d, f, and g: 01 2 3 4 s p d f g Value of l Letter used

46. Translate Quantum Numbers into Orbitals • An orbital with quantum numbers of n = 3 and l = 2 would be a 3d orbital • An orbital with quantum numbers of n = 4 and l = 1 would be a 4p orbital

47. Quantum Numbers 3) Magnetic Quantum Number (ml) • The magnetic quantum number describes the three-dimensional orientation of the orbital. • Allowed values of mlare integers ranging from - l to l: − l ≤ ml≤ l. • Therefore, on any given energy level, there can be up to 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, etc.

48. Orbitals • Orbitals with the same value of n form a shell. • Different orbital types within a shell are subshells. −l ≤ ml ≤ l 0…(n-1) Can draw orbital diagrams

49. s Orbitals n = 1, 2, 3, 4,…… l = 0, 1, 2….(n-1) where 0=s, 1=p, 2=d, 3=f ml = - l ….0…+ l • The value of lfor s orbitals is 0 and therefore ml=0 • They are spherical in shape. • The radius of the sphere increases with the value of n. • found in all shells of an atom

50. s Orbitals Observing a graph of probabilities of finding an electron versus distance from the nucleus, we see that s orbitals possess n−1 nodes, or regions where there is 0 probability of finding an electron.