The Atom and Its Properties

# The Atom and Its Properties

## The Atom and Its Properties

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##### Presentation Transcript

1. The Atom and Its Properties Chapter 4 – Nucleus Chapter 5 – Electron Configuration

2. Chapter 4 Objectives Describe an atom’s structure and differentiate among the particles that make it up. Identify the numbers associated with elements and explain their meaning . Realize that the number of protons in a nucleus defines an element. Calculate the average atomic mass given isotopes and relative abundance

3. Chapter 4 Vocabulary Chapter 4.3 • Atomic Number • Isotope • Mass Number • AMU (Atomic Mass Unit) • Average Atomic Mass Chapter 4.1 • Dalton’s Atomic Theory • Atom Chapter 4.2 • Electron • Nucleus • Proton • Neutron

4. Modern View of the AtomThe nucleus is where the protons and neutrons are located and contain most of the atom’s mass.

5. Protons, Neutrons and Electrons

6. Sometimes Atomic Symbols are Displayed as:

7. Isotope Examples 24 52

8. What’s all this amu business? To simplify a system of indicating atomic masses since protons and neutrons have such extremely small masses, scientists have assigned thecarbon-12 atom a mass of exactly 12 atomic mass units. (amu) The mass of 1 amu (1/12 the mass of carbon-12) is very nearly equal to the mass of a single proton or neutron but not the same. 1 amu = 1.66 x 10-24 grams

9. Isotopes and Mass Number • Your text (p. 119) shows how to calculate the mass number for Cl given the % abundance of the isotopes. • Let’s do this for another element: Li • 6Li is 7.59 % abundant; 6.015 amu • 7Li is 92.41% abundant; 7.015 amu • Method 1: Use percentages. Think of this as a sample of 100 atoms.

10. In Tabular Form

11. Average Atomic Mass • The average mass of an atom is found by weighting the natural abundances of its isotopes. • Lithium (Method 2): Change % to fraction. • 6Li 6.015 amu 7.59% = 0.0759 • 7Li 7.015 amu 92.41% = 0.9241 Mass (amu) Fracabund Mass share Avg mass = 6.015 amu x 0.0759 = 0.46 amu 7.015 amu x 0.9241 = 6.48 amu 6.94 amu/atom

12. Displayed on Periodic Table

13. Electrons in Atoms Chapter 5

14. Chapter 5 Objectives Compare wave and particle matters of light See how frequency of light emitted by an atom is unique to that atom Compare and contrast the Bohr and quantum mechanical models of the atom Express the arrangements of electrons in atoms through orbital notations, electron configurations, and electron dot structures

15. Chapter 5 Vocabulary Chapter 5.2 • Ground State • Quantum Number • Quantum Mechanical Model of the Atom • Atomic Orbital • Principal Quantum Number • Principal Energy Level • Energy Sublevel Chapter 5.1 Electromagnetic Radiation Wavelength Frequency Amplitude Electromagnetic Spectrum Quantum Photoelectric Effect Photon Atomic Emission Spectrum

16. Wave Nature of Light Electromagnetic radiation displays wavelike behavior as it travels through space Waves can be described by several common characteristics

17. Characteristics of a Wave • Waves transfer energy • Properties of waves: • Frequency (ν – pronounced ‘nu’) - Number of vibrations per unit time – Hz (cycle/second) • Wavelength (λ) - Distance between points on two consecutive waves • Speed of wave is Frequency x wavelength Speed = ν x λ Frequency is the number of waves that hit this point in one second. Amplitude λ

18. Electromagnetic Spectrum Note: All EM Radiation travels at 3.00 x 108 m/s

19. Electromagnetic Spectrum The speed of light (3.00  108 m/s) is the product of it’s wavelength and frequency c = λν.

20. The Electromagnetic Spectrum – all light is energy

21. Electromagnetic Spectrum • Gamma Rays – Highest frequency, shortest wavelength. Can pass through most substances • X Rays – Lower frequency than Gamma rays. Can pass through soft body tissue but can’t pass through bone. • Ultraviolet (UV) Rays – Part of sunlight that causes sunburn

22. Electromagnetic Spectrum • Visible Light – Sensitive to our eyes. Allows us to see color • Infrared – Less energy and longer wavelength than visible light. Felt as heat given off a heater or near a fire • Radio Waves – Lowest frequencies on the EM spectrum. Used by radio and over-the-air TV.

23. An electromagnetic wave has a frequency of 6.0 x 104 Hz. Convert this frequency into its corresponding wavelength. Which region of the EM spectrum does this correspond to? It’s a radio wave (~103 meters) l x n = c l = c/n • = 3.00 x 108m/s • 6.0 x 104 /s l = 5.0 x 103m 7.1

24. Practice Problems Answers • What is the frequency of green light, which has a wavelength of 520 nm. • A radio station broadcasts at 94.7 MHz. What is the wavelength of the broadcast?

25. Nature of Light Max Planck (1858-1947) studied the different light emitted from heated objects Matter can only gain or lose energy in small specific amounts

26. Nature of Light A quantum is the minimum amount of energy that can be gained or lost by an atom The energy of EM radiation is proportional to its frequency (E αν)

27. Photons Albert Einstein (1879-1955) proposed that while a beam of light had wavelike characteristics, it also can be thought of as a stream of tiny particles (or bundles of energy) called photons Each photon carries a quantum of energy

28. Particle Nature of Light The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.

29. When copper is bombarded with high-energy electrons, X rays are emitted. Calculate the energy (in Joules) of the X-rays if their frequency is 1.95 x 1018 Hz. E = h x n E = 6.63 x 10-34 (J•s) x 1.95 x 1018 /s E = 1.29 x 10 -15 J 7.2

30. Ch. 5.2 – Quantum Theory of the Atom Each element has only certain specific frequencies of light that are emitted when atoms absorb energy and become excited Where do we see this? fireworks neon signs stars

31. Hydrogen Spectrum

32. Balmer Plot In 1885, Johann Balmer observed the lines of the spectrum fit this surprisingly simple formula: Where n1 =2 and n2 = 3, 4, 5, etc.

33. Balmer Plot RH is the slope of this line, 1.0972 x 107 m-1

34. Electronic Energy Transitions Neils Bohr (1885-1962) proposed the model the hydrogen atom (1913) to explain the discreet nature of the hydrogen spectrum.

35. Electronic Energy Transitions • Neils Bohr’s model the atom (1913) • Electrons exist only in discrete, “allowable” energy levels • Energy is involved in moving electrons from one energy level to another • Principal quantum number (n) - specifies the electron’s major energy level • The lowest energy is n=1, the next lowest in n=2, etc.

36. Bohr’s Model of the Atom (cont’d) Bohr suggested that an electron moves around the nucleus in only certain allowed circular orbits. n = 2 n = 1

37. Energy Absorption/Emission

38. Atomic Emission Spectra

39. Origin of Line Spectra Balmer series

40. Quantum or Wave Mechanics Schrodinger applied idea of e- behaving as a wave to the problem of electrons in atoms. He developed the WAVE EQUATION Solution gives set of math expressions called WAVE FUNCTIONS. Treated electrons as wavelike particles that became the Quantum Mechanical Model of the Atom. E. Schrodinger 1887-1961

41. Waves Wave motion: wave length and nodes“Quantization” in a standing wave

42. Hydrogen Atom Solution Where: a0 is the Bohr Radius given by a0 = 4πεoh2/me2 Generalized Laguerre Polynomial m here is quantum number Constant = 2.18 x 10-18 J m is mass of electron

43. Atomic Orbitals-Hydrogen

44. Orbitals • No more than 2 e- assigned to an orbital • Orbitals grouped in s, p, d (and f) sublevels s orbitals p orbitals d orbitals

45. s orbitals p orbitals d orbitals s orbitals p orbitals d orbitals No. orbs. 1 3 5 No. e- 2 6 10

46. Energy Levels and Sublevels • Sublevels are grouped in energy level. • Each energy level has a number called thePRINCIPAL QUANTUM NUMBER, n which indicates relative size and energy of the orbitals • Row on PT indicates n