UNIT 2: Atomic Theory Section 1: Atom Basics Section 2: Isotopes Section 3: Electron Configuration Section 4: History of Atomic Theory

Transcription

1 UNIT 2: Atomic Theory Section 1: Atom Basics Section 2: Isotopes Section 3: Electron Configuration Section 4: History of Atomic Theory

2 UNIT 2 Synapsis In our second unit we will take a in depth look at atoms. We will start in parts 1 and 2 with some basic things that you may already know, and a few other basic things you may not know. Then we will explore how atoms of the same element can be different because of how many neutrons they have. In part 3 we will look at how the behavior of the electrons in an atom can be described by brushing the surface of quantum mechanics. Finally we will wrap up the unit by learning about some of the major discoveries and people that brought us to our current understanding of atoms.

3 Section 1: Atom Basics This part of the Unit is covered on pages in your textbook

4 Section 1: Atom Basics / Objectives After this lesson I can define an atom & its three parts. Identify an element when give a number of protons and a Periodic Table recall the charge, mass, symbol and location of all three sub atomic particles. give an analogy for the size of the nucleus compared to the rest of the atom. recall that the strong nuclear force and the neutrons are what overcome the repulsive force between the protons and holds the nucleus together in stable atoms.

5 Protons Atoms are the fundamental building block of matter and we learned previously that matter is basically everything. All atoms are 1 of 118 elements. An element is the type or identity of an atom and is determined by the number of protons the atom has. The number of protons an atom has is also called it s atomic number. Examples: 1 proton = hydrogen 8 protons = oxygen 26 protons = iron. 92 protons = uranium Protons are Located in the Nucleus (the center of the atom) Positively charged have a mass of 1 a.m.u (atomic mass unit). go by the symbol p + or H +

6 Neutrons & Electrons 98% of hydrogen atoms contain just 1 proton. Atoms of all the other elements and a small percentage of Hydrogen atoms have neutrons as well. Neutrons are Located in the nucleus Have a mass of 1 a.m.u. Do not have a charge Use the symbol n Finally there are the electrons. Electrons don t change what element an atom is, and don t really matter for nuclear change. They are the real role players during chemical change because they are shared or exchanged when atoms bond with each other. Electrons are Located outside the nucleus Have a mass so small (.0002 a.m.u.) we often say their mass is negligible Have a negative charge Use the symbol e -

7 Very Simple Model of the Atom Image Credit: encyclopedia Britannica

8 Another Very Simple Model of the Atom Image Credit: googlesites.com

9 The Nucleus You can think of the nucleus as being this tightly packed ball of protons and neutrons. Unlike the models we are going to use to describe the electrons later in the unit, this is actually very close to reality. heavier atoms with higher numbers of protons and neutrons are actually shaped more like footballs but you get the point. Students often wonder how positively charged protons could be tightly packed together given that positive charges repel one another. There are two explanations why: The neutrons help to spread the charge out The strong nuclear force binds the protons together The other major things you need to know about the Nucleus can be difficult to imagine because there is nothing in the big world that really compares to the reality of an atom and it s nucleus. The nucleus is roughly 100,000 times smaller than the atom as a whole contains all of the atoms mass. In other words, the nucleus is all of the mass and none of the volume.

10 Model of the Nucleus Image Credit: :mscl.msu.edu

11 Analogies for the Nucleus Compare to the Rest of the Atom If the atom were the size of ford field the nucleus would be the size of a marble If the atom were the size of a church the nucleus would be just barely visible. About the size of a spec of dust suspend in air If the atom were the size of new York city the nucleus would be about the size of an orange

12 The Nucleus & The Electron Cloud So what takes up all the space within an atom then? The answer is nothing. Atoms are almost completely empty space (about 99.99%) The electrons occupy the vast region outside of the nucleus that we call the electron cloud. They are very tiny though and do not come close to taking up all that space. They don t fill up the electron cloud, it s simply where they can be found

13 Video Time!!! Crash Course Chemistry #1 Video: The Nucleus TedED Video: Just How Small is an Atom?

14 Section 1 Additional Resources & Links Science Post Video: Basic Parts of the Atom prontons, neutrons, electrons, nucleus Fuse School Video: What is an Atom

15 Section 2: Isotopes This part of the Unit is covered on pages in your textbook

16 After this lesson I can define isotopes. Section 2: Isotopes / Objectives determine the number of protons & neutrons in isotopes write isotope names and symbols (mass numbers) determine the number of electrons in an ion. distinguish mass numbers, atomic mass, and average atomic mass. calculate the average atomic mass of an element when give the percent abundance, atomic mass, and the mass numbers of it s naturally occurring isotopes (when given the Table of Isotopic Masses and Natural Abundances )

17 Video Time!!! Veratsium Video: What are Atoms and Isotopes?

18 Isotopes Earlier it was noted that the number of protons in an atom s nucleus determines what element that atom is, but atoms can have a different number of neutrons and still be the same element. An isotope is the type or version of an element and is determined by the number of neutrons. Isotopes are different atoms of the same element. Isotopes have a mass number which indicates the total number of protons & neutrons. If the isotope is given by name, the mass number is written after the dash. If the isotope is given by symbol, the mass number is written as a superscript to the left of the symbol. Isotope examples: Carbon-12 = (6 protons, 6 neutrons) = 12 C Carbon-13 = (6 protons, 7 neutrons) = 13 C Carbon-14 = (6 protons, 8 neutrons) = 14 C Sometimes the number of protons is written below the mass number. This is redundant because the number of protons is given by the element name or symbol but there are examples in a few slides.

19 Additional Examples of Isotopes Nitrogen-14 = (7 proton, 7 neutrons) = 14 N Nitrogen-15 = (7 protons,8 neutrons) 15 N Oxygen-15 = (8 protons, 7 neutrons) = 15 O Oxygen-16 = (8 protons, 8 neutrons) = 16 O Oxygen-17 = (8 protons, 9 neutrons) = 17 O Krypton 71 = (36 protons, 35 neutrons) = 71 Kr Krypton 74 = (36 protons, 38 neutrons) = 74 Kr Krypton 76 = (36 protons, 40 neutrons) = 76 Kr Note that Nitrogen-15 & Oxygen-15 have the same mass number but are different isotopes and different elements.

20 Isotope Image #1 Image Credit: :kaffee.50webs.com/science/activities/chem/activity.isotopes.table.htm

21 Isotope Image #2 Image Credit:

22 Isotope Image with Atomic Number Included Image Credit: astronomy.swin.edu.au

23 Isotope Image #4 Image Credit: highschoolpedia.com

24 Practice Problems: Decoding Isotopes Directions: Write the number of protons & neutrons for each isotope below 1) 25 Al 2) 33 P 3) Silver ) Uranium ) 181 Au 6) 30 S 7) 122 Xe 8) Beryllium -9 9) 4 He 10) Nickel ) 56 Fe 12) 38 Ar 13) 73 As 14) Tellurium - 134

25 Practice Problems: Writing Isotopes Directions: Write the isotope symbol and name for the given atom 1) 15 protons, 16 neutrons 2) 27 protons, 32 neutrons 3) 79 protons, 100 neutrons 4) 79 protons, 102 neutrons 5) 31 protons, 36 neutrons 6) 8 protons, 8 neutrons 7) 100 protons, 165 neutrons 8) 78 protons, 99 neutrons 9) 34 protons, 42 neutrons

26 Ions Atoms usually have the same number of electrons as they do protons and are thus neutral overall. Atoms that have more or less electrons than they do protons are called ions. When an atom is an ion it has a charge written as a superscript on the right side of the element s symbol. Examples: Aluminum atom with 10 electrons: Al 3+ Iron atom with 23 electrons: Fe 3+ Chlorine atom with 18 electrons: Cl - Oxygen atom with 10 electrons: O 2- Sodium atom with 10 electrons Na +

27 Example and Explanation of Isotope or Atom Symbols

28 Practice Problems: Ions Directions: Write down the element symbol & charge 1) 15 protons, 18 electrons 2) 2 protons, 0 electrons 3) 20 protons, 18 electrons 4) 23 protons, 18 electrons 5) 28 protons, 26 electrons 6) 35 protons, 36 electrons 7) 7 protons, 10 electrons

29 Practice Problems: Ions Directions: Determine the number of protons, neutrons, & electrons in each ion 1) 45 As 3-2) 29 Si 3) 139 I - 4) 71 Se 2-5) 33 S 2-6) 57 Fe 3+ 7) 35 Cl - 8) 17 O 2-9) 36 Ar

30 Percent Abundance Some isotopes of a given element are more common the others. For example 98.9% of all carbon atoms on planet earth are Carbon-12, only 1% are Carbon-13, & less than.1% are Carbon-14. The percentage of an isotope for a given element as it naturally occurs on earth is called it s natural abundance or percent abundance. Most elements have at least 2 isotopes that occur naturally, some have just 1, other have 5 or more. 10 is the highest (Tin if your curious) You can see all the naturally occurring isotopes for all the elements in a table called Table of Isotopic Masses and Natural Abundances. The vast majority of isotopes that naturally occur are known as stable isotopes. Basically that means the nucleus is going to exist forever as far as we can tell. Unstable isotopes will break down and most elements have a dozen or more unstable isotopes with very few occurring naturally. These isotopes radioactively decay or undergo fission. Studying these isotopes is a branch of chemistry called nuclear chemistry. Nuclear chemistry will not be covered in Chemistry A or B this year. However, there are a few really cool short videos that talk about nuclear chemistry at the end of this section.

31 Difference between Mass Number & Atomic Mass The mass number of an isotope is it s total number of protons and neutrons. The atomic mass of an isotope is it s true mass that we have determined through repeated experiments with very precise instruments. You d think these numbers would be identical, but they are not. All isotopes have a very slight difference between mass number and atomic mass. The one exception to this is Carbon-12 which has a atomic mass of exactly Understanding why there is a slight difference is not in our learning objectives. You can find the actual atomic mass for isotopes that naturally occur on the same table that lists natural abundances. It s worth noting that the difference between atomic mass and mass number is so slight, that many educators and instructional videos you find on the internet will not even mention it.

32 Average Atomic Mass The Average Atomic Mass of an element, which is the number that appears below atomic number on your periodic table, is the average atomic mass of the natural occurring isotopes of an element. It takes into account each isotopes percent abundance. If an element only has 1 naturally occurring isotope (and many do) than the average atomic mass is just the true atomic mass of it s only naturally occurring isotope. Using the Table of Isotopic Masses and Percent Abundances you should be able to calculate the average atomic mass that appears for each element on the periodic table. There are two examples on the next two slides.

33 Average Atomic Mass Calculation Example: Titanium Naturally Occurring Isotopes of Titanium Isotope Atomic Mass % Abundance 46 Ti % 47 Ti % 48 Ti % 49 Ti % 50 Ti % 46 Ti: = Ti: = Ti: = Ti: = Ti: = ( ) = 47.87

34 Average Atomic Mass Calculation Example: Neon 20 Ne: = Ne: = Ne: = ( ) = 20.18

35 Section 2 Additional Resources & Links Tyler Dewitt s Video: What are Isotopes Tyler Dewitt s Video: Isotope Notation Tyler Dewitt s Video: Atomic Mass: An Introduction Tyler Dewitt s Video: How to Calculate Atomic Mass: Practice Problems YouChemTutorials Video: Calculating Average Atomic Mass

36 Section 3: Electron Configuration This part of the Unit is covered on pages 146 to 162 in your textbook

37 Section 3: Electron Configuration / Objectives After this lesson I can recall the 4 sub-shells or sub-levels and the maximum number of electrons they can hold. create the electron configuration cheat chart. write and identify electron configurations for elements and mono-atomic ions write and identify electron configurations using noble gas shorthand write and identify electron configurations using orbital diagrams. recall the four major facts about electrons mentioned in lecture.

38 Writing Electron Configuration This is the order that the electrons fill the orbitals in: 1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p So the electron configuration for Ognesson, which has 118 electrons is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6 The electron configuration for Carbon, which has 6 electrons is: 1s 2 2s 2 2p 2 The electron configuration for Iron, which has 26 electrons is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 The superscript on the sub-shell indicates the number of electrons in that subshell. Iron doesn t fill the 3d subshell just like Carbon doesn t fill the 2p subshell. Electrons follow the aufbau principal and fill up each subshell before moving on to the next one. There is a cheat chart to help you remember the order in which the subshells fill up in 2 slides.

39 Electron Configuration of Sulfur with Parts Labeled

40 Writing Electron Configuration Cheat Chart

41 Steps for Writing Electron Configuration Step 0: Write the cheat chart on a piece of scrap paper Step 1: Determine the total number of electrons. This is the atomic number in the case of neutral atoms. Step 2: If your dealing with an ion, add or subtract electrons accordingly. Step 3: Fill up the orbitals according to the cheat chart until you run out of electrons. Step 4: Counting up all the electrons in your configuration to confirm you have included all the electrons from Steps 1 or 2.

42 Practice Problems: Writing Electron Configuration Directions: Write the electron configuration for the following. 1) O 2) P 3) Mo 4) Fe 5) I - 6) Cu 7) Mg 2+

43 Writing Electron Configuration Using Noble Gas Short-Hand When the electron configuration is very long a short hand method is available using noble gases. Only noble gases can be used to short hand or abbreviate. When abbreviating with noble gases you just stick in the noble gas with brackets around it for however many electrons that noble gas has. Examples are below: Example #1: Silicon s full version is 1s 2 2s 2 2p 6 3s 2 3p 2 However, the short hand could be written [Ne]3s 2 3p 2 Example #2: Iron s full version is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 However, a short hand for Iron could be written [Ar]4s 2 3d 6 Example #3: Rubidium s full version is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 However, a short hand for Rubidium could be written [Kr]5s 1 Example #4: Bismuth s full version is 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 10 4f 14 6p 3 However, a short hand for Bismuth could be written [Xe]6s 2 5d 10 4f 14 6p 3 The more electrons an Atom has the more sense it makes to abbreviate because electron configurations can bet pretty long

44 Periodic Table Showing Noble Gases In Yellow

45 Practice Problems: Electron Configuration w/ Noble Gas Short Hand Directions: Write the electron configuration for the following using noble gas short hand. 1) O: 1s 2 2s 2 2p 4 2) P: 1s 2 2s 2 2p 6 3s 2 3p 3 3) Mo: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 4 4) Fe: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 5) I: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 6) Cu: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 7) Ni 2+ : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6

46 Hund s Rule & the Pauli Exclusion Principal Electron configuration does not show the specific orbital the electrons are in because p sub-shells actually contain up to 3 orbitals d sub-shells actually contain up to 5 orbitals f sub-shells actually contain up to 7 orbitals A more accurate method of showing the distribution of electrons in an atom are orbital diagrams. In orbital diagrams, each orbital is represented by a box or line. The electrons are represented by arrows. Orbital diagrams introduce two more important facets of quantum mechanics: Hund's rule: basically states that electrons will occupy a new orbital if one is available at the same energy level before occupying a orbital that already has an electron in it. Pauli exclusion Principal: Basically states that no more than 2 electrons can occupy an orbital. Furthermore, when two electrons occupy the same orbital, they must spin in opposite directions.

47 Writing Electron Configuration Using Orbital Diagrams What Hund s rule and Pauli s principal translate too when using orbital diagrams is that you add 1 electron to each box before doubling up. When you do double up, you make sure the arrows that represent the electrons are pointed in different directions. When creating orbital diagrams, you will want to write out the full electron configuration first (not the noble gas short hand) and then make your boxes and arrows s sub-shells should have 1 box or line p sub-shells should have 3 boxes or lines d sub-shells should have 5 boxes or lines f sub-shells should have 7 boxes or lines

48 Orbital Diagrams (Each Box = 1 Orbital) Image Credit: Thoughtco.com

49 Orbital Diagram Examples Image Credit: Opentextbc.ca

50 Practice Problems: Orbital Diagrams Directions: Draw the Orbital Diagrams for the Elements Below 1) C 2) Ni 3) Fe +

51 Video Time!!! CCC #5 Video: The Electron Cosmos Clip Video: The Electron TedED Video: Particles and Waves The Central Mystery of Quantum Mechanics

52 Quantum Mechanical Model of Electrons Now that we know how to write electron configurations & orbital diagrams, lets talk just a little bit about what electron configurations actually mean. First let me say that understanding the electrons in an atom is way, way more difficult than understanding the nucleus of an atom. The primary reason for this is that electrons behave in ways that have no similarities to the macroscopic world. As Neil Degrasse Tyson puts it electrons do not correspond to ordinary human experience, common sense is no help here at all. Just the word we use to describe the electrons in an atom quantum mechanics can be discouraging. Furthermore, the behavior of electrons is so complex that even today we still do not have a complete understanding of it. Despite this, at least part of the quantum mechanical model can be described in a fairly simple way that students often find is easier than it sounds. This is the whole point of orbital diagrams and electron configurations.

53 Atomic Orbital Basics Each atomic orbital in an atom consists of three parts: the principal quantum number. This is also referd to as the energy level or shell and there are 7 shells each represent by the numbers 1 7 The angular momentum. This is also referred to as the sub shell. There are four different subshells: s sub-shell that can hold 2 electrons p sub-shell that can hold 6 electrons d sub-shell that can hold 10 electrons f sub-shell that can hold 14 electrons And finally the orbital itself. The each individual orbital has a specific designation (for example orbitals within the p subshell have the desgination 3 px, 3 py, & 3 pz ). Knowing the designation of the specific orbitalis not in our objectives. Each atomic orbital can hold up to two electrons. Ultimately atomic orbitals represent a region of space outside the nucleus where there is a 95% probability of finding the electron that is in that orbital. Some of the actual shapes of the orbitals are on the next slide.

54 Orbital Shapes Around the Nucleus where Electrons Can Be Found s orbitals Spherical shape p orbitals Dumbell shape Image Credit: Thoughtco.com

55 Some Facts you Should Know 1) Electrons do NOT orbit the nucleus like planets orbit the sun. It s okay for beginning chemistry students to think of it that way though and we will model atoms this way in class. 2) Electrons have the properties of both standing waves and particles. 3) Electrons can move between these orbitals when they release or absorb light or energy, but they can not exist in the space between these orbitals. 4) It is impossible to know the exact position and speed of an electron at the same time (Heisenberg uncertainty principal).

56 Periodic Table Showing Principal Energy Levels & Sub-Shells Image Credit: Thoughtco.com

57 The Aufbau Principal How many energy levels and subshells there are, and how many electrons are in each sub-shell, can be written out using electron configuration. Electron configuration basically shows how the electrons are distributed within the electron cloud of the atom. It shows what energy level and subshell the electrons are in (not the specific orbital though that is what orbital diagrams are for). Writing electron configuration is easy because of the aufbau principal. The aufbau principal basically states that electrons fill atomic orbitals of the lowest energy level before moving on to the next one. To put it another way, the atoms will fill the atomic orbitals in the exact same order every time (there are exceptions to this but we will not worry about those exceptions).

58 Video Time!!! TedED Video: What is the Heisenberg Uncertainty Principle TedED Video: The Uncertain Location of Electrons TedED Video: Why Glass is Transparent

59 Section 4: History of Atomic Theory This part of the Unit is covered on pages 110 to 114 in your textbook

60 Section 4: History of Atomic Theory / Objectives After this lesson I can summarize how our model of the atom has changed over the last 2,500 years citing the four major contributors mentioned in lecture and their experiments. draw a picture of the Greek model of the atom. draw a picture of the Plum Pudding model of the atom and label the parts draw a picture of the Rutherford model of the atom and label the parts. explain, describe, and summarize Rutherford s Gold Foil experiment. draw a picture of the Bohr model of the atom and label the parts.

61 Video Time!!! Crash Course Chemistry # 37 Video: History of Atomic Chemistry TedED Video: 2,400 year search for the Atom

62 History of Atomic Theory Overview For the first three section of this unit we have looked at Atomic Theory; what humanity has come to know and understand about atoms through verifiable observations and experiments. Now we are going to take a brief look at some of the major players who have helped us arrive at our current understanding. The first thing worth mentioning when it comes to Atomic Theory is that no one person really developed atomic theory on their own like Darwin did with evolution. Hundreds of people made noteworthy contributions, most of them living within the last 100 years or so. We are only going to look at four people who made major revisions to atomic theory and how our model of the atom changed because of them: Democritus JJ Thompson Ernest Rutherford Niels Bohr

63 ~400 B.C.E.: Democritus s Atom Democritus was a Greek philosopher who first came up with the idea of atoms. He thought of atoms as indivisible, solid, and spherical. Democritus was NOT a scientist and did not conduct experiments to test his ideas like true scientists do. In fact, science as we define it today did not even exist back then. You might call Democritus s experiment a thought experiment though. He came up with the idea of atoms simply by thinking about what happens if you keep cutting an object in half over and over again. He reasoned that if you keep cutting or dividing an object sooner or later you would get to a point where it could not be cut or divided any further. Atomos in Greek means that which cannot be split. We know today the idea of a of atoms as tiny, indivisible solid spheres is far from true, but we sometimes still model them that way. For example when looking at the structures of molecules or compounds each individual atom is represented by a colored sphere.

64 Democritus s Model of the Atom

65 JJ Thompson discovers the electron JJ Thompson was a scientist who discovered electrons through experiments he did with cathode ray tubes. He also discovered that these negatively charged particles coming out of atoms were 2,000 times less massive than the atom itself. Like many scientific discoveries, this led to more questions. JJ s major dilemma was that since he knew atoms where neutral overall, how could the positive charges be accounted for? Thompson hypothesized that atoms were a lot like plumb pudding, which was a popular desert at the time. He thought that the negatively charged electrons (plums) were spread throughout a positively charged spherical mass (pudding) To use a desert you are more familiar with, he thought atoms were like chocolate chip cookies. The chocolate chips were the electrons and the cookie dough was a positively charged mass. This was a reasonable conclusion given that protons and neutrons were yet to be discovered. However, as we now know it was very inaccurate and the plum pudding model has virtually no use today. Beyond discovering the electron, JJ was the first to realize that atoms could be subdivided, they could be broken down into smaller parts

66 JJ s Plumb Pudding Model

67 1907 Ernest Rutherford Discovers the Nucleus Ernest Rutherford s gold foil experiment is perhaps the most famous experiment in the history of chemistry; it is part of the reason why Rutherford got an element named after him. Rutherford discovered the nucleus and that atoms are mostly empty space. He did this by shooting alpha particles at a piece of gold foil. Because most of the alpha particles passed through the foil, but a small amount bounced back, he concluded that atoms must have a dense, positively charged center, or nucleus, but for the most part were empty space. Rutherford didn t know that the nucleus could be broken down into a specific number of protons and neutrons, or that protons and neutrons even existed for that matter. But he did realize atoms had a nucleus that contained almost all the atoms mass, that it was several thousand times smaller than the atom itself, and that it was positively charged. The discovery of protons came two years later, and shortly after that explanations for the behavior of the electrons (something that continues to evolve to this day) would come as well.

68 Rutherford s Nucleus centered Model

69 1913 Niels Bohr Proposes a New Model to Explain the Electrons While Rutherford showed that the electrons of an atom were outside the nucleus, he was not able to explain the behavior of those electrons. Unlike Rutherford and his gold foil or Thompson and his cathode ray tube, Bohr did not have a famous experiment of his own. What Bohr did was explain the behavior of electrons based on existing observations and mathematics. He hypothesized that: Electrons travel around the nucleus in fixed orbits. These orbits were regions outside the nucleus where the electrons were allowed to be. They could move between these orbits when they released or absorbed light, but could not exist in the space between. Each orbit can only hold a certain number of electrons For his contribution to atomic theory Bohr received a Nobel prize and eventually had a element named after him. Later scientists and physicists would build upon his ideas.

70 Bohr s Planetary Model

71 Section 4 Additional Resources & Links Tyler Dewitt s Video: Model of the Atom Timeline Tyler Dewitt s Video: Discovery of the Nucleus: Rutherford s Gold Foil Experiment. Khan Academy s Video: Rutherford s Gold Foil Experiment. McGraw Hill Video: The Gold Foil Experiment Simulator Program for the Gold Foil Experiment.