1. The Atom Mr. Sackman South Dade Senior High 2010

2. History The Greeks Dalton J.J. Thompson Millikan Rutherford Chadwick Bohr

3. History The Greeks were the first to attempt to describe matter and atoms Philosophers were intellectual thinkers of the time and anything that they said many believed without argument The Greeks first classified matter as Earth, Wind, Water, and Fire Their ideas were creative, however, there was no way to test their theories at the time; another reason many just accepted what philosophers said They believed that matter could endlessly, meaning infinitely, be divided into smaller and smaller pieces with no end

4. History Democritus – First to propose that matter isn’t infinitely divisible – Believed matter was made of tiny particles called atomos (atoms) – Believed atoms could not be created, destroyed, or further divided – Matter is composed of empty space through which atoms move

5. History – Atoms are solid, homogenous, and indivisible – Different kinds of atoms have different shapes and sizes – The differing properties of matter are due to the size, shape, and movement of atoms – Changes in matter can only be caused by changes in grouping of atoms and not from changes in the atoms themselves – His thinking was way ahead of his time and some of his ideas still hold

6. History Aristotle (384 B.C.-322B.C.) – One of the most influential minds of his time – Gained wide acceptance for his view on nature – What ever he stated most accepted, or believed, to be fact or true – He rejected atomic theory all together simply because of his own ideas didn’t agree – His major argument was that matter isn't empty space through which atoms move, he didn’t believe nothingness could exist

7. History – Democritus was unable to answers challenges to his ideas paving the way for Aristotle's beliefs – Democritus' ideas were then eventually thrown out – Aristotle’s theory was accepted and he threw out the existence of atoms altogether – Because of Aristotle’s influence the answer to the question of the acceptance, or denial, of atoms went unchallenged for ~2000 years

8. History John Dalton (1766-1844) – Finally thousands of years later someone attempted to describe the atom – He marked the beginning of modern atomic theory – Science now allowed for the study of matter and attempted to prove the existence of atoms – Proposed new atomic theory in 1803 – Some of his theories were the same as Democritus

9. History -Dalton’s Atomic Theory states the following -All matter is composed of extremely small particles called atoms -All atoms of the same element are identical and atoms of different elements differ completely from others -Atoms cannot be created, destroyed, or divided -Different numbers of atoms combine in simple whole number ratios to form compounds and in chemical reactions atoms are separated, combined, or rearranged

10. History Recall the law of conservation of mass? Dalton's theory easily explains this law by stating that atoms are only separated or rearranged in reactions which would neither create or destroy atoms The law of definite proportions states that no matter how large the sample is a compound is always composed of the same elements in the same proportion by mass. For example water is always 11% H and 89% O no matter how large or small the sample

11. Dalton had used some form of technology and the new aged science to refine Democritus' theory Dalton observed and recorded numerous reactions making careful observations, and measurements, as he performed his experiments, does this process sound familiar?

12. History – Using Lavoisier's, and Proust’s, ideas he had come up with his own – Is his theory completely correct? – How did his theory differ from that of the Greeks? – What do you think gave him the advantage of acceptance of his theory at that time?

13. Defining the Atom What is the actual definition of an atom? What does an atom look like? Can you picture something that is so small you can’t see?

14. Defining the Atom Suppose you decide one day you want to grind your pure silver necklace down, how far could you grind? Could you eventually grind down far enough that you reach something that is not divisible or visible? Does every smaller and smaller piece you grind retain the properties of what you are grinding down?

15. Defining the Atom Exactly how small is an atom? – The book gives a very good example and it is Consider the size of the population in the world which in 2000 was about 6 billion or 6,000,000,000 now compare that to how many Cu atoms are in a penny which is 29,000,000,000,000,000,000,000 This is almost 5 billion times more copper atoms than people

16. Defining the Atom Now take the diameter of that same penny, 1.20X10 -10 m If one were to place six billion copper atoms side by side, this is the same as the world population, the line of copper atoms would be less than the length of a meter stick

17. Defining the Atom Can one actually see an atom with technology these days? Now can you begin to see, or picture, the size and existence of atoms?

18. The electron Everyone here knows what an electron is, or do you? How do we prove the existence of an electron Did we set out to prove there was such a particle called the electron or was it accident? Curiosity sparked the investigation between electrical charge and matter By accident, one day, Henry Crookes noticed a flash of light from one of his tubes he created while working in a dark laboratory

19. The electron These flashes were the result of something striking a light producing coating applied at one end of a cathode tube Further investigation showed that that a stream appeared to flow from the cathode to the anode This device led to the one of the most important social developments of all time, T.V. (old school ones) and computer monitors. (also old school ones) Pictures on these screens are just formed when radiation from the cathode strikes light producing chemicals that coat the backside of a screen producing an image

20. The electron Research showed this stream of light was actually a ray of particles not just some invisible rays The particles were shown to carry a negative charge when a magnetic either deflected or attracted the stream. How could you prove this with only a magnet?

22. The electron J.J. Thompson (1856-1940) – Began a series of cathode ray tube experiments when using Crooke’s technology – He calculated both the magnetic and electrical fields and found the mass to charge ratio – When comparing this ratio to other known ratios he concluded that this charged particle was actually lighter than a hydrogen atom – What did the mass of this particle being less than that of the smallest atom prove?

23. The electron – When changing the matter that filled the tube the results were the same, what does this mean? – His theory went unaccepted for some time as many still believed the atom is indivisible, Dalton’s theory – What did Thompson just prove and demonstrate, how important was this at this time?

24. The Electron Millikan (1856-1940) – Determined, and proved, the charge of an electron to be negative in 1909 – His technique, and set up, was so accurate that his value found in 1909 still only has an error of ~1.0% – Charge was determined to be that of a single charge, meaning,-1 – Now knowing the charge, and the mass to charge ratio discovered my Thompson, he calculated the mass to be 9.1 X10 -28 g which is 1/1840 the mass of a hydrogen atom; what does this mean?

25. The Proton Matter is in a electrically neutral state most of the time so how are atoms neutral if they carry a particle with a negative charge? – There must be a positive charge as well – J.J. Thompson proposes plum pudding model describes the atom to be a spherical shape of uniformly distributed positive charge spherically around the atom with the electrons packed inside.

26. The Proton The plum pudding model doesn’t hold for long when Rutherford comes around In 1911, Ernest Rutherford, simply interested on how alpha particles interacted with matter, began a series of experiments He conducted experiments to see if alpha particles were deflected if passed through a thing sheet of gold foil, this experiment was a breakthrough in atomic theory

27. The Proton Knowing Thompsons model Rutherford expected only minor deflections of alpha particles Believed that if there was any deflection it would be due to the collision, or near collision, of a negatively charged electron He also believed that the positive charge was so uniform throughout it wouldn't deflect the massive alpha particles His results were stunning and opposite of what he expected

28. As you saw in the pictures of this gold foil experiment some of the rays went directly through the foil, some were deflected a little, and some deflected straight back to the source, what did this indicate? The ones that passed right trough were not interacting with the atom as they passed through empty space The rays deflected were due to the interaction of the positively charge nucleus of atoms, the closer the ray to the nucleus the great the deflection

29. He proposed that the atom was mostly empty space with a central, and extremely, densely packed nucleus which contained all the positive charge

30. The Nucleus An atom is mostly empty space so how big, or small is an atom? We already studied the size of an atom, but these particles are smaller than the atom itself, aren't they? If the nucleus were the size of a dot on an exclamation point than the mass of that nucleus would be that of ~70 automobiles If the atom had the diameter of a football field the nucleus would be the size of only a nickel What does this mean?

31. The Neutron Now putting all the results together we discover the mass of a electron, and that of a proton, did not account for the entire mass of the lightest known atom, what does this mean? The answer came with the discovery of the neutron James Chadwick, a student of Rutherford discovered another particle in the nucleus with no charge that had a mass equal to that of a proton both ~1.675X10 -24 g, or 1 amu, much larger than the mass of an electron which is nearly equal mass to the mass of a proton, both subatomic particles are given a mass of 1amu ParticleSymbolLocationRelative electrical charge Relative Mass (amu) Actual Mass Electrone-e- In the space around some place around the nucleus 1-1/18409.11X10 -28 ProtonP+P+ Nucleus1+11.673X10 -24 NeutronN0N0 Nucleus011.675X10 -24

32. How Atoms Differ Not long after Rutherford a man named, Henry Moseley, demonstrated that atoms of each different element had a unique positive charge in the nucleus, what does this mean? The number of protons defines, or identifies, the element and is called the atomic number(Z), each proton has a charge of +1, the opposite of the electron, and has a mass of 1amu Since the mass of a neutron, and proton, are equal and more massive than an electron both particles added together for the mass number(A) and the electron isn't

33. How Atoms Differ How could on calculate the number of neutrons from these two numbers? If the mass number is equal to the number of protons plus neutrons, and the atomic number is the number of protons, all we have to do is subtract the atomic number from the mass number, A-Z What are the letters for each element on the periodic table, why is the first always capitalized and the second always lower case?

34. Isotopes If the mass of a neutron, and a proton, both equal 1 amu why are the masses on the periodic table not whole numbers? The answer is because of isotopes When you take a sample of an element the sample contains the element plus any of its isotopes Isotopes are the same element but have different number of neutrons causing the masses to be different Why couldn’t there be isotopes of the same element with different numbers of protons?

35. Isotopes JJ Thompson discovered one day that he two separate samples of neon gas, the same element, but they had different masses, how could this be? The answer again is isotopes, if the number of neutrons is different the mass will be as well.

36. Isotopes The reason the masses on the periodic table are not whole numbers is also because of isotopes As stated before a sample of any given element will contain its isotopes as but one of them will be more abundant than the others and expressed in a percent of all isotopes called % abundance If you take the % abundance, for each isotope, and multiply it by the isotopes mass then sum the values for all isotopes you will get the weighted atomic mass which is what is expressed on the periodic table.

37. Isotopes Lets demonstrate calculation of the weighted atomic mass Element X has 4 different isotopes what is its weighted atomic mass using the data below (this is not a real element, remember the % abundance must always be changed to a decimal when multiplying Isotope 1, 67.0% abundance, mass = 1.03 amu Isotope 2, 3.00% abundance, mass = 1.23 amu Isotope 3, 17.00% abundance, mass = 1.09 amu Isotope 4, 13% abundance, mass =1.10 amu

38. (.6700)(1.03) + (.0300)(1.23) + (.1700)(1.09) + (.1300)(1.03) = you figure it out and ask me

39. Isotope X-6 has a mass of 6.015 amu and a percent abundance of 7.5%, isotope X-7 has a mass of 92.5%, and a % abundance of 92.5 what element is this on the periodic table using method just shown? Again try it on your own then ask me if you are correct

40. Atomic Models Planetary Bohr Model Which model is correct?

41. As we already know the atom has a centrally located nucleus with all the protons and neutrons inside the nucleus And the atom has mostly empty space with electrons orbiting in some place around the nucleus at any given time but is the end of the description of the atomic model?

42. Atomic Model How many electrons can actually orbit an atom in the same area, or orbit? Is there a certain way electrons are placed, or found, to be around the atom at any given time?

43. Atomic Model The next model proposed, after the plum pudding model, is called the planetary model, as its name suggest you expect to find the electrons orbiting the nucleus but all electrons have the same orbit This means if I were to take sodium (Na), I would draw all the protons and neutrons in the nucleus and all 11 electrons in one orbit around the nucleus, is this accurate?

45. Instead of saying we have electrons around the nucleus in an orbit we can say in an energy level Energy levels are where the electrons have the highest probability of being found when attempting to locate them.

46. Atomic Model Niels Bohr, a student of Rutherford, proposed his own model in 1913 his model was quantized,which you will learn for now very basics, you will learn the reasoning behind the Bohr Model in more detail later when we study light and quantized energy

47. Atomic Structure Bohr proposed that electrons can only be placed in certain energy levels that orbit in certain circular orbits around an atom and that each of these energy levels can only hold a certain number of electrons The first energy level is the closest to the nucleus, and every energy level added to the atom causes the atomic radius to increase

48. Atomic Structure Now know that the first 5 energy levels can hold the following number of electrons: 2,8,18,32,50 To find the number of electrons use the following formula 2n 2, where n is the number of the energy level Now take sodium (Na) and draw the Bohr Model for an atom. You will need 3 energy levels for this one, meaning, three orbits around the nucleus each one successively larger in diameter going away from the nucleus, draw the model now

49. The electrons cannot choose any orbit they wish. They are restricted to orbits with only certain energies. Electrons can jump from one energy level to a higher one, or from a higher one to a lower one, only when a specific amount of energy is absorbed or emitted

50. There is no in between energy levels, the electrons can only be in the ground state or an excited state. For example think of a ladder and the rungs, or steps. The steps of the ladder are energy levels of an atom and your feet are the electrons. The lowest step is the ground state and the higher steps are excited states Can you stand in between, the rungs, or steps, of the ladder,? This is the same concepts are electrons and energy levels?

51. That exact, specific, or required, amount of energy, enough to place them in a higher energy level, or lower energy level is called a quantum, plural quanta Electrons can be promoted from the lowest energy state to a higher energy state as long as they have absorbed the same quantum of energy that the higher state possesses

52. When an electron goes down from that excited state, back to the ground state, it has to give off that same amount, quanta, of energy because energy can not be destroyed When that electron drops back down to a lower energy state it gives off that energy as a photon which can be seen as light The amount of energy that the photon contains is just the difference in energy between the excited state and the lower state

53. Atomic Structure Draw the following atoms according to the planetary model and the Bohr Model Na, Mg, Br, H, He, Ne, Hg, Cu, Fe, Xe, C, Rb