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UNIT 2: Atomic Theory
Part 1: Atom Basics
Part 2: Isotopes
Part 3: Nuclear Change
Part 4: Electron Configuration
Part 5: History of Atomic Theory
Unit Synapsis
In our second unit we will take a in depth look at atoms, and how atoms change during nuclear changes.
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 take a much deeper look at the various ways that atoms can change; in stars, in atomic bombs, in nuclear plants, and through natural decay processes.
In part 4 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.
Part 1: Atom BasicsThis part of the Unit is covered on pages 111-116 in your textbook
Part 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.
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+”
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-”
Very Simple Model of the Atom
Image Credit: encyclopedia Britannica
Another Very Simple Model of the Atom
Image Credit: googlesites.com
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 shaped more like footballs.
• 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 that charge out
• The strong nuclear force binds them 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.
Model of the Nucleus
Image Credit: mscl.msu.edu
Analogies for the Nucleus and 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”
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
VIDEOS!
Show the Video: Crash Course Chemistry #1 – The NucleusShow the Video: Just How Small is an Atom?Show the Video: Atoms and Fusion
Part 1 Additional Resources
• Crash Course Chemistry #1: The Nucleus
• TedED Video: Just How Small is the atom?
Part 2: IsotopesThis part of the Unit is covered on pages 115-121 in your textbook
Part 2: Isotopes / Objectives
“After this lesson I can…
• …define isotopes.
• …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”)
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) = 12C
• “Carbon-13” = (6 protons, 7 neutrons) = 13C
• “Carbon-14” = (6 protons, 8 neutrons) = 14C
• 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.
Additional Isotope Examples
“Nitrogen-14” = (7 proton, 7 neutrons) = 14N
“Nitrogen-15” = (7 protons,8 neutrons) 15N
“Oxygen-15” = (8 protons, 7 neutrons) = 15O
“Oxygen-16” = (8 protons, 8 neutrons) = 16O
“Oxygen-17” = (8 protons, 9 neutrons) = 17O
“Krypton – 71” = (36 protons, 35 neutrons) = 71Kr
“Krypton – 74” = (36 protons, 38 neutrons) = 74Kr
“Krypton – 76” = (36 protons, 40 neutrons) = 76Kr
• Note that Nitrogen-15 & Oxygen-15 have the same mass number but are different isotopes and different elements.
Isotope Examples
Image Credit: kaffee.50webs.com/Science/activities/Chem/Activity.Isotopes.Table.htm
Two Isotopes of Iodine
Image Credit www.sprawls.org/ppmi2/MATTER/
Isotope Example With Atomic Number Included
Image Credit: astronomy.swin.edu.au
Isotope Example With Atomic Number Included
Image Credit: highschoolpedia.com
Practice Problems: Isotopes and Mass Numbers
Directions: Write the Number of Protons & Neutrons for each Isotope below...
1) 25Al
2) 33P
3) Silver - 109
4) Uranium - 238
5) 181Au
6) 30S
7) 122Xe
8) Beryllium -9
9) 4He
10) Nickel - 60
11) 56Fe
12) 38Ar
13) 73As
14) Tellurium - 134
Practice Problems: Isotopes and Mass Numbers
Directions: Write the Isotope Symbol and Isotope 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
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: Al3+
• Iron atom with 23 electrons: Fe3+
• Chlorine atom with 18 electrons: Cl-
• Oxygen atom with 10 electrons: O2-
• Sodium atom with 10 electrons Na+
Practice Problems: Ions
Directions: Write down the ion (element symbol with 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
Directions: determine the number of protons, neutrons and electrons in each ion…
1) 40As3-
2) 29Si4+
3) 139I-
4) 71Se2-
Example with Explanation of Isotope or Atom Symbols
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”.
Mass Number and 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 12.000000.
• Understanding why there is a slight difference is really not too difficult. However, we are not going to dedicate class time to it because there are more important objectives to focus on.
• 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.
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.
Average Atomic Mass Calculation Example: Titanium
• 46Ti: 45.953 × .0825 = 3.791
• 47Ti: 46.692 × .0744 = 3.474
• 48Ti: 47.948 × .7372 = 35.347
• 49Ti: 48.948 × .0541 = 2.648
• 50Ti: 49.945 × .0518 = 2.587
• (3.791+3.474+35.347+2.648+2.587) = 47.87
Naturally Occurring Isotopes of Titanium
Isotope Atomic Mass % Abundance46Ti 45.953 8.25%47Ti 46.962 7.44%48Ti 47.948 73.72%49Ti 48.948 5.41%50Ti 49.945 5.18%
Average Atomic Mass Calculation Example: Neon
• 20Ne: 19.992 × .9048 = 18.089
• 21Ne: 20.994 × .0027 = .056
• 22Ne: 21.991 × .0925= 2.034
• (18.089+.056+2.034) = 20.18
Part 2 Additional Resources
• 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”
Part 3: Nuclear ChangeThis part of the Unit is covered on pages 860 to 884 in your textbook
(note: We won’t be covering all the concepts in those pages)
Part 3: Nuclear Change / Objectives
“After this lesson I can…• …categorize elements on the periodic table as:
• left over from the big bang• formed during the normal life cycle of stars• formed only in supernova explosions
• …explain & create an illustrative model nuclear fusion.”• …explain fission, including where it most commonly occurs
and the isotope that is used.” • …compare and contrast stable and unstable isotopes.”• …define half-life and explain general trends in half life.”• …use a graph of nuclides to determine if an isotope is stable or
unstable, as well as the type of decay an unstable isotope will undergo.”
• …determine the products and write equations for alpha decay, beta plus decay, and beta minus decay.”
• …explain decay chains and how they end.”• …identify equations that model fusion, fission, alpha decay,
beta plus decay, and beta minus decay.”• … create an illustrative model of radioactive decay
The Big Bang, Fusion, and Nucleosynthesis
• The universe began from a singularity we call the big bang approximately 13.8 billion years ago.
• After the big bang, it was about 100,000,000 years or so before things cooled down and the first stars formed.
• At this point in time, the only elements that existed in the universe were Hydrogen, Helium, and Lithium; none of the other elements on the periodic table existed.
• Almost all of it was Hydrogen-1, and to this day roughly 75% of all the matter in the universe is Hydrogen - 1.
• Almost all of the other elements on the periodic table are created inside the center of stars.
• The process by which stars create larger elements is called nuclear fusion or nucleosynthesis. Fusion is when smaller atoms combine to form larger ones.
• Small stars fuse hydrogen into helium and maybe some helium into lithium and beryllium before they die. Larger stars can fuse element up to Iron in their lifetime before they die.
• All the elements above Iron are formed in a supernova, which is when a massive star dies.
Examples of Fusion Equations
• There are three types of fusion reactions happening in our sun simultaneously. Products from the first step are used as reactants in the second step and products from the second step are used as reactants in the third step.
• Step 1: 1H + 1H 2H + ν + e+ + energy
• Step 2: 2H + 1H 3He + γ + energy
• Step 3: 3He + 3He 4He + 1H + 1H + energy
• Our sun will never do these reactions but here are equations for fusion of larger elements that happen in bigger stars:
• 3He + 4He 7Be
• 8Be + 4He 12C
• 16O + 4He 20Ne
• 24Mg + 4He 28Si
• Note that in all the fusion equations at least 1 product has a mass number larger than either of the reactants.
Fusion of Hydrogen-2 & Hydrogen-3 to form Helium-4
Image Credit: Wikipedia
Fusion of Helium to form Beryllium and then Carbon
Image Credit Wikipedia
VIDEOS!
Show the Video: Crash Course Big History #1 – The Big Bang ClipShow the Video: Crash Course Big History #2 – Fusion ClipShow the Video: Cosmos Clip: Fusion
Stable and Unstable Isotopes
• The opposite of Fusion is fission. In fission, larger atoms breakdown into smaller ones.
• Fission happens in unstable isotopes, especially Uranium-235.
• Isotopes can be stable or unstable. Stable isotopes (also called stable nuclides) are isotopes that will never change without an outside force acting on them. Stable isotopes will basically exist forever and never change into other atoms.
• Elements with atomic numbers up to 82 have at least 1 stable isotope except 43 and 61 (Technetium and Promethium). Elements 83 and above have no stable isotopes.
• Unstable isotopes (also called radioactive isotopes) are atoms that will change spontaneously without any outside force acting on them. Basically unstable isotopes will breakdown into another element at some point in time. When this happens it’s called radioactive decay.
• The time it takes for an unstable isotope to break down is called the isotopes half-life. Half-life is more narrowly defined as the time it takes half the atoms in a given sample to breakdown.
• For example: Carbon-14 has a half life of 5,730 years. So if you had a sample of 50,000 Carbon-14 atoms, then 5,730 years later you would only have 25,000 left.
Periodic Table Showing Isotope Data
VIDEOS!
Show the Video: FissionShow the Video: Crash Course Chemistry 39 – Fusion and FissionShow the Video: Cosmos Clip – Radioactivity and Half Lives
Simple Model of Fission
Image Credit: ap.smu.ca
Model of Fission Showing how it will Chain React
Image Credit: Jim Doyle
Unstable Isotopes and Radioactive Decay
• Half-life times of isotopes have a wide range.
• Some isotopes have half-lives in billions and trillions of years. Those atoms have been here since the earth formed 4.5 billion years ago.
• Other isotopes have half-lives of 100 or even 1,000 times less than a second. These isotopes only exist because we create them in particle accelerators, or because they are the products of other unstable isotopes decaying.
• General rules about half-lives
• The shorter the half-life, the more radioactive and more dangerous the isotope is.
• As atomic number increases beyond 82, half-life decreases for the most stable isotope of a given element.
• For elements that have stable isotopes, the half-lives of their unstable isotopes get shorter as you get “farther away” from the stable isotopes.
• A “graph of nuclides” is a useful reference tool that shows all the atoms we know of. Both the stable isotopes of an element and it’s unstable ones. They are commonly represented showing type of decay or length of half life.
Graph Nuclides Showing Type of Decay
Graph Nuclides Showing Length of Half Lives
Practice Problems: Isotopes and Mass Numbers
Directions: Use a Graph or Table of Nuclides to determine if each isotope is stable or unstable. If it is unstable, give the mode of decay. It is best to write the number of protons and neutrons out first.
1) Carbon-13
2) Oxygen – 19
3) Oxygen – 14
4) 238U
5) 65Fe
6) 57Cu
7) Tin – 120
8) 164Er
Radioactive Decay
• As mentioned on a previous slide when an unstable isotope breaks down through a natural process it’s called radioactive decay. In radioactive decay elements change into different elements. This is called transmutation.
• There are roughly 6 types of radioactive decay. 3 of these are pretty rare and 3 of them are much more common.
• The rare ones include:
• Proton emission: An atom kicks out a single proton
• Neutron emission: An atom kicks out a single neutron
• Spontaneous Fission: An atom splits apart it two ore more smaller atoms on it its own. It’s only different from the fission we discussed previously in that it happens naturally. Usually when the word fission is used it’s referring to human caused fission, most commonly that of Uranium-235 inside a nuclear reactor.
• We are not going to focus on these types of radioactive decay and they will not appear on quizzes or tests. They are only being mentioned for clarification because you will see them appear on some of the graphs and tables we will use in class.
Alpha Decay
• Alpha Decay is a type of radioactive decay that occurs most often in the isotopes of elements above Lead.
• In alpha decay, the nucleus of an atom ejects a Helium-4 atom with no electrons.
• “4He2+ ” is sometimes represented by the symbol: “α” which is the Greek letter “alpha.”
• Examples of Alpha Decay
• 238U 234Th + α
• 238U 234Th + 4He2+
• 222Ra 218Rn + 4He2+
• 222Ra 218Rn+ α
• Alpha Decay has the generic equation:
• Where
• X = unstable isotope or “parent isotope”
• Y = new isotope or “daughter isotope”
• Z = atomic number
• A = mass number
Alpha Decay Model #1
Image Credit: http://www.physicslovers.com/radioactivity/radioactive-decay
Alpha Decay Model #2
Image Credit: Study.com
Practice Problems: Alpha Decay
Directions: Write the alpha decay equation for the following parent isotopes. You can use either “α” or “4He2+” for the alpha particle.
1) Iridium – 167
2) 150W
3) Actinium – 221
4) 294Og
5) 259Lr
Beta Minus Decay “β-”
• Beta Minus Decay is most common when an isotope has more neutrons than stable isotopes of the same element.
• In Beta Minus Decay, a neutron in the parent isotope changes into a proton. An electron and anti-neutrino is ejected. It’s worth noting that in Beta Minus Decay the mass number does not change.
• The symbol for an electron is “e-” or “β-”and the symbol for an anti-neutrino is “ν”
• Examples of Beta Minus Decay:
• 121Pd 121Ag + e- + ν
• 75Zn 75Ga + e- + ν
• 25Ne 25Na + β- + ν
• Beta Minus Decay has the generic equation:
• Where
• X = unstable isotope or “parent isotope”
• Y = new isotope or “daughter isotope”
• Z = atomic number
• A = mass number
Practice Problems: Beta Minus Decay
Directions: Write the beta minus decay Equation for the following Isotopes. You can write the electron as “e-” or “β-”
1) 97Y
2) Boron – 15
3) Chromium – 68
4) 152Pr
5) 151Ba
Beta Plus Decay “β+” also called Positron Decay
• Beta Plus Decay is the exact opposite of Beta Minus Decay
• Beta Plus Decay is most common when an isotope has less neutrons than stable isotopes of the same element.
• In Beta Plus Decay, a proton in the parent isotope changes into a neutron. A positron (anti-electron) and a neutrino is ejected. It’s worth noting that in Beta Plus Decay the mass number does not change.
• The symbol for a positron is “e+” or “β+” and the symbol for a neutrino is “ν”.
• Examples of Beta Plus Decay:
• 48Mn 48Cr + e+ + ν
• 73Kr 73Br + e+ + ν
• 105In 105Cd + e+ + ν
• Beta Plus Decay has the generic equation:
• Where
• X = unstable isotope or “parent isotope”
• Y = new isotope or “daughter isotope”
• Z = atomic number
• A = mass number
Practice Problems: Beta Plus Decay
Directions: Write the beta plus decay equation for the following parent isotopes. You can write the positron as “e+” or “β+”
1) Europium – 137
2) 71Br
3) Gold – 187
4) 100Rh
5) 65Ge
Beta Plus and Minus Decay Model
Image Credit: https://education.jlab.org/glossary/betadecay.html
Decay Chains
• When a isotope radioactively decays, most of the time the daughter isotope is another unstable isotope. This means that unstable atom is going to radioactively decay sooner or later as well, whatever it decays into may very well be another unstable isotope.
• This process, sometimes continues for more than a dozen unstable isotopes before a stable isotope is reached. It can take trillions of years for that to happen or several minutes.
• This process is called a decay chain and there are models of it on the next slide
• Don’t confuse a decay chain with a chain reaction.
• A chain reaction is what happens when Uranium-235 undergoes rapid, self-sustaining fission inside nuclear power plants or in nuclear bombs.
Decay Chain Model of Thorium-232
Image Credit: Wikipedia
Decay Chain Model of Uranium-238
Image Credit: Berkeley.edu
VIDEOS!
Show the Video: Crash Course Chemistry 38: Radioactive Decay Show the Video: Cosmos Clip – Radioactivity and Half Lives
Final Notes on Nuclear Change
• Beta Plus and Beta Minus Decay are easy to get mixed up. One good way to remember the difference between the two.
• Beta Minus means a negative particle or an electron will get emitted
• Remember that there is always 1 particle of matter and 1 particle of anti-matter.
• Some isotopes don’t always decay via the same mode each time. For example, Bismuth-214 sometimes alpha decays and sometimes Beta Minus Decays. The “Table of Nuclides by Type of Decay” shows the most common type of decay for a given isotope if there is more than one way it will go.
• Earlier it was mentioned that elements above Iron are formed in supernovas. This is only true up to element 94. 95 and above are all man made in particle accelerators. These elements are called synthetic elements.
Part 3 Additional Resources
• Crash Course Big History #1 Video: The Big Bang• Note: Watch from 5:00 to 12:40
• Crash Course Big History #2 Video: Exploring the Universe• Note: Watch from 8:20 to 11:30
• Fission Video
• Crash Course Chemistry #39 Video: Fusion and Fission• Note: Watch from 3:25 to the end and ignore binding energy
• Tyler Dewitt’s Video: Nuclear Fission
• Crash Course Chemistry #38 Video: Radioactive Decay• Note: Watch from the start of the video to 7:02
• Tyler Dewitt’s Video: Alpha Decay
• Tyler Dewitt’s Video: Beta Decay (Beta Minus Decay)
• Tyler Dewitt’s Video: Positron Decay (Beta Plus Decay)
• Khan Academy Video: Types of Decay
Part 4: Electron ConfigurationThis part of the Unit is covered on pages 146 to 162 in your textbook
Part 4: Electron Configuration / Objectives
“After this lesson I can…
• …recall the four major facts about electrons mentioned in lecture.
• …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 gasshorthand
• …write and identify electron configurations using orbital diagrams.
Quantum Mechanical Model of Electrons
• Understanding the electrons in an atom is way, way more difficult than understanding the nucleus of an atom and how it can change. 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.
Some facts about electrons
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) Electron are only allowed in specific regions of the electron cloud called atomic orbitals (terrible name given that they don’t orbit).
4) 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.
VIDEOS!
Show the Video: Crash Course Chemistry 5: The ElectronShow the Video: Cosmos Clip – Atomic Orbitals
Atomic Orbital Basics
• Each atomic orbital consists of three parts; the principal quantum number or energy level, the angular moment, which we are going to call the “sub-shell” or “sub-level,” and the orbital itself.
• There are 7 “energy levels” and they are numbered 1 – 7
• There are 4 “sub-shells”
• “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
• Finally each specific orbital has a crazy designation which we are not going to get into. What you do need to know is that each atomic orbital can hold 2 and only 2 electrons. That means that…
• s sub-shells only have 1 orbital
• p sub-shells have 3 orbitals
• d sub-shells have 5 orbitals
• f sub-shells have 7 orbitals.
• The energy level and the sub-shells are why the periodic is arraigned the way it’s arraigned.
Periodic Table Showing the Principal Energy Levels & “Sub-shells”
Image Credit: Thoughtco.com
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 they are in (not the specific orbital though).
• 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).
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:
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d107p6
• The electron configuration for Carbon, which has 6 electrons is:
1s2 2s2 2p2
• The electron configuration for Iron, which has 26 electrons is:
1s2 2s2 2p6 3s2 3p6 4s2 3d6
• 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.
• As mention before, 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 on the next slide.
Electron Configuration of Sulfur with Parts Labeled
Writing Electron Configuration Cheat Chart
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.
Practice Problems: Writing Electron Configuration
Directions: Write the electron configuration for the following elements or ions.
1) O
2) P
3) Mo
4) Fe
5) I
6) Cu
7) Mg2+
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.
• Helium with 2 electrons is used for elements with 3 to 9 electrons
• Example: Fluorine could be written [He]2s22p5
• Neon with 10 electrons is used for elements with 11 to 17 electrons
• Example: Silicon could be written [Ne]3s23p2
• Argon with 18 electrons is used for elements with 19 to 35 electrons
• Example: Iron could be written [Ar]4s23d6
• Krypton with 36 electrons is used for elements with 37 to 53 electrons
• Example: Rubidium could be written [Kr]5s1
• Xenon with 54 electrons is used for elements with 55 to 85 electrons
• Example: Bismuth could be written [Xe]6s25d104f146p3
• Radon with 86 electrons is used for elements with 87 to 117 electrons
Periodic Table Showing Noble Gases in Yellow
Practice Problems: Electron configuration w/ Noble Gas Short-hand
Directions: Write the electron configuration for the following elements or ions using noble gas short-hand.
1) O: 1s2 2s2 2p4
2) P: 1s2 2s2 2p6 3s2 3p3
3) Mo: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d4
4) Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
5) I: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5
6) Cu: 1s2 2s2 2p6 3s2 3p6 4s2 3d9
7) Ni2+: 1s2 2s2 2p6 3s2 3p6 4s2 3d6
Hund’s Rule and the Pauli Exclusion Principal
• As mentioned earlier 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.
Writing Electron Configuration Using Orbitals 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
Orbitals Diagrams (Each Box = 1 Orbital)
Image Credit: Thoughtco.com
Orbital Diagram Examples
Image Credit: Opentextbc.ca
Orbital Diagram Examples
Image Credit: Oklahoma State
Practice Problems: Orbital Diagrams
Directions: Draw the Orbital Diagrams for the elements below
1) C
2) Ni
3) Fe3+
Model of the Electron Configuration of Silicon
Image Credit: Wikipedia
Model of the Electron Configuration of Manganese
Image Credit: Wikipedia
Model of the Electron Configuration of Terbium
Image Credit: Wikipedia
Model of the Electron Configuration of Platinum
Image Credit: Quora.com
Model of the Electron Configuration of Oganesson
Image Credit: Wikipedia
Part 4 Additional Resources
• Crash Course Chemistry Video #5: The Electron
Part 5: History of Atomic TheoryThis part of the Unit is covered on pages 110 to 114 in your textbook
Part 5: 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.
History of Atomic Theory Overview
• For the first four parts 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. Dozens 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
~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.
Democritus’s Model of the Atom
1897 - 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.
JJ’s Plumb Pudding Model
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.
Rutherford’s Nucleus Centered Model of the atom
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.
Bohr Model of atom
Part 5 Additional Resources & Links
• Tyler Dewitt’s Video: Model of the Atom Timeline
• Crash Course Chemistry # 37 Video: History of Atomic Chemistry
• TedED Video: 2,400 year search for the Atom
• 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.
