2. Unit 4 – Atomic Structure Bravo – 15,000 kilotons

3. Atoms, Ions, and the Periodic Table What is an atom? It is smallest particle of an element that retains the elements properties. But how did we come to know all the information we have about these tiny particle?

4. Democritus (460-370 BC)

5. Matter is made of tiny, solid, indivisible particles which he called atoms (from atomos, the Greek word for indivisible). Different kinds of atoms have different sizes and shapes. Different properties of matter are due to the differences in size, shape, and movement of atoms. Democritus’ ideas, though correct, were widely rejected by his peers, most notably Aristotle (384-322 BC). Aristotle was a very influential Greek philosopher who had a different view of matter. He believed that everything was composed of the four elements earth, air, fire, and water. Because at that time in history, Democritus’ ideas about the atom could not be tested experimentally, the opinions of well-known Aristotle won out. Democritus’ ideas were not revived until John Dalton developed his atomic theory in the 19th century!

6. John Dalton (1766-1844)

7. –All matter is composed of extremely small particles called atoms. –All atoms of one element are identical. –Atoms of a given element are different from those of any other element. –Atoms of one element combine with atoms of another element to form compounds. –Atoms are indivisible. In addition, they cannot be created or destroyed, just rearranged.

8. Dalton’s theory was of critical importance. He was able to support his ideas through experimentation, and his work revolutionized scientists’ concept of matter and its smallest building block, the atom. Dalton’s theory has two flaws: –In point #2, this is not completely true. Isotopes of a given element are not totally identical; they differ in the number of neutrons. Scientists did not at this time know about isotopes. –In point #5, atoms are not indivisible. Atoms are made of even smaller particles (protons, neutrons, electrons). Atoms can be broken down, but only in a nuclear reaction, which Dalton was unfamiliar with.

9. Discovery of the Electron JJ Thomson (1856-1940)

10. Discovered the electron, and determined that it had a negative charge, by experimentation with cathode ray tubes. A cathode ray tube is a glass tube in which electrons flow due to opposing charges at each end. Televisions and computer monitors contain cathode ray tubes. Thomson developed a model of the atom called the plum pudding model. It showed evenly distributed negative electrons in a uniform positive cage. Diagram of plum pudding model:

11. Discovery of the Nucleus Ernest Rutherford (1871-1937) http://www.mhhe.com/physsci/chemistry/animations/chang_2e/rut herfords_experiment.swf

12. Discovery of the Nucleus Ernest Rutherford (1871-1937) Discovered the nucleus of the atom in his famous Gold Foil Experiment. Alpha particles (helium nuclei) produced from the radioactive decay of polonium streamed toward a sheet of gold foil. To Rutherford’s great surprise, some of the alpha particles bounced off of the gold foil. This meant that they were hitting a dense, relatively large object, which Rutherford called the nucleus.

13. Rutherford then discovered the proton, and next, working with a colleague, James Chadwick (1891-1974), he discovered the neutron as well.

14. Models of the Atom - Niehls Bohr

15. Developed the Bohr model of the atom (1913) in which electrons are restricted to specific energies and follow paths called orbits a fixed distance from the nucleus. This is similar to the way the planets orbit the sun. However, electrons do not have neat orbits like the planets. Diagram of Bohr model:

16. Quantum Mechanical Model

17. This is the current model of the atom. We now know that electrons exist in regions of space around the nucleus, but their paths cannot be predicted. The electron’s motion is random and we can only talk about the probability of an electron being in a certain region.

18. Modern Atomic Theory  Atoms of an element have a characteristic average mass which is unique to that element.  Atoms cannot be subdivided, created, or destroyed in ordinary chemical reactions. However, these changes CAN occur in nuclear reactions!  All matter is composed of atoms  Atoms of any one element differ in properties from atoms of another element

19. Discovery of the Electron In 1897, J.J. Thomson used a cathode ray tube to deduce the presence of a negatively charged particle. Cathode ray tubes pass electricity through a gas that is contained at a very low pressure.

20. Conclusions from the Study of the Electron  Cathode rays have identical properties regardless of the element used to produce them. All elements must contain identically charged electrons.  Atoms are neutral, so there must be positive particles in the atom to balance the negative charge of the electrons  Electrons have so little mass that atoms must contain other particles that account for most of the mass

21. Thomson’s Atomic Model Thomson believed that the electrons were like plums embedded in a positively charged “pudding,” thus it was called the “plum pudding” model.

22. Rutherford’s Gold Foil Experiment  Alpha particles are helium nuclei  Particles were fired at a thin sheet of gold foil  Particle hits on the detecting screen (film) are recorded

23. Rutherford’s Findings  The nucleus is small  The nucleus is dense  The nucleus is positively charged  Most of the particles passed right through  A few particles were deflected  VERY FEW were greatly deflected “Like howitzer shells bouncing off of tissue paper!” Conclusions:

24. Atomic Particles ParticleChargeMass #Location Electron0Electron cloud Proton+11Nucleus Neutron01Nucleus

25. The Atomic Scale  Most of the mass of the atom is in the nucleus (protons and neutrons)  Electrons are found outside of the nucleus (the electron cloud)  Most of the volume of the atom is empty space “q” is a particle called a “quark”

26. About Quarks… Protons and neutrons are NOT fundamental particles. Protons are made of two “up” quarks and one “down” quark. Neutrons are made of one “up” quark and two “down” quarks. Quarks are held together by “gluons”

27. Atomic Number Atomic number (Z) of an element is the number of protons in the nucleus of each atom of that element. Element# of protonsAtomic # (Z) Carbon66 Phosphorus15 Gold79

28. Mass Number Mass number is the number of protons and neutrons in the nucleus of an isotope. Mass # = p + + n 0 Nuclidep+p+ n0n0 e-e- Mass # Oxygen -10 -3342 - 3115 8818 Arsenic753375 Phosphorus153116

29. Isotopes Isotopes are atoms of the same element having different masses due to varying numbers of neutrons. IsotopeProtonsElectronsNeutronsNucleus Hydrogen–1 (protium) 110 Hydrogen-2 (deuterium) 111 Hydrogen-3 (tritium) 112

30. Atomic Masses IsotopeSymbolComposition of the nucleus % in nature Carbon-12 12 C6 protons 6 neutrons 98.89% Carbon-13 13 C6 protons 7 neutrons 1.11% Carbon-14 14 C6 protons 8 neutrons <0.01% Atomic mass is the average of all the naturally isotopes of that element. Carbon = 12.011

31. Isotopes Isotopes are atoms of an element with the same number of protons but different numbers of neutrons. Change in # of n 0 = Change in the Mass # Most elements on the periodic table have more than one naturally occurring isotope. There are a couple of ways to represent the different isotopes. One way is to put the mass after the name or symbol: Carbon-12 or C-12 Another way is :

32. Determining Average Atomic Mass The atomic mass on the periodic table is determined using a weighted average of all the isotopes of that atom. In order to determine the average atomic mass, –convert the percent abundance to a decimal –multiply it by the mass of that isotope –values for all the isotopes are added to together to get the average atomic mass Formula: (Mass 1 * Abundance 1 ) + (Mass 2 * Abundance 2 ) + (Mass 3 * Abundance 3 ) + etc….

33. Example of Average atomic mass calculation Given: 12 C = 98.89% at 12 amu 13 C = 1.11% at 13.0034 amu Calculation: (98.89%*12 amu) + (1.11%*13.0034 amu) = (0.9889*12 amu) + (0.011*13.0034 amu) = 12.01 amu

34. Now you try one: Neon has 3 isotopes: Neon-20 has a mass of 19.992 amu and an abundance of 90.51%. Neon-21 has a mass of 20.994 amu and an abundance of 0.27%. Neon-22 has a mass of 21.991 amu and an abundance of 9.22%. What is the average atomic mass of neon? The answer is: (0.9051 * 19.992 amu) + (0.0027 * 20.994 amu) + (0.0922 * 21.991 amu) = 20.179 amu Now compare this mass for Neon to the mass on the periodic table!

35. The Mole 1 dozen = 1 gross = 1 ream = 1 mole = 12 144 500 6.02 x 10 23 There are exactly 12 grams of carbon-12 in one mole of carbon-12.

36. Avogadro’s Number 6.02 x 10 23 is called “Avogadro’s Number” in honor of the Italian chemist Amadeo Avogadro (1776-1855). Amadeo Avogadro I didn’t discover it. Its just named after me!

38. MOLE MOLAR MASS6.02 x 10 23 particles Molecular MassMolecules Atomic Mass (really amu’s)Atoms Formula MassFormula Units

39. Calculations with Moles: Converting moles to grams How many grams of lithium are in 3.50 moles of lithium? 3.50 mol Li = g Li 1 mol Li 6.94 g Li 24.29

40. Calculations with Moles: Converting grams to moles How many moles of lithium are in 18.2 grams of lithium? 18.2 g Li = mol Li 6.94 g Li 1 mol Li 2.62

41. Calculations with Moles: Using Avogadro’s Number How many atoms of lithium are in 3.50 moles of lithium? 3.50 mol Li = atoms Li 1 mol Li 6.02 x 10 23 atoms Li 2.11 x 10 24

42. Calculations with Moles: Using Avogadro’s Number How many atoms of lithium are in 18.2 g of lithium? 18.2 g Li = atoms Li 1 mol Li6.02 x 10 23 atoms Li 1.58 x 10 24 6.94 g Li1 mol Li (18.2)(6.022 x 10 23 )/6.94

43. Nuclear Symbols Element symbol Mass number (p + + n o ) Atomic number (number of p + )

44. Types of Radioactive Decay  alpha production (  ): helium nucleus  beta production (  ): 2 4 He  1 0 e 2+

45. Alpha Radiation Limited to VERY large nucleii.

46. Beta Radiation Converts a neutron into a proton.

47. Types of Radioactive Decay  gamma ray production (  ):  positron production :  electron capture: (inner-orbital electron is captured by the nucleus) 1 0 e

48. Types of Radiation

49. Deflection of Decay Particles Opposite charges_________ each other. Like charges_________ each other. attract repel

50. Nuclear Stability Decay will occur in such a way as to return a nucleus to the band (line) of stability.

51. Half-life Concept

52. Sample Half-Lives

53. A Decay Series A radioactive nucleus reaches a stable state by a series of steps

54. Nuclear Fission and Fusion Fusion: Combining two light nuclei to form a heavier, more stable nucleus. Fission: Splitting a heavy nucleus into two nuclei with smaller mass numbers.

55. Energy and Mass Nuclear changes occur with small but measurable losses of mass. The lost mass is called the mass defect, and is converted to energy according to Einstein’s equation:  E =  mc 2  m = mass defect  E = change in energy c = speed of light Because c 2 is so large, even small amounts of mass are converted to enormous amount of energy.

56. Fission

57. Fission Processes A self-sustaining fission process is called a chain reaction.

58. A Fission Reactor

59. Fusion