Unit+3+The+Atom

The Atom The following unit discusses how our knowledge of the atom came to be, starting from its origin back as a simple hypothesis. It explains the discovery of atoms, how atoms are classified together, the properties of which are only found in all atoms, and the characteristics of atoms which differentiate atoms from each other. The discovery and analysis brought to the atom for many centuries has formed the very fabric of our understanding of all matter found in the universe today. Scientists are still building off of the most simplistic studies of the atom brought to the world through scientific pioneers such as John Dalton in order to further expand our already vast understanding of the subject in order to further our knowledge of the world.

Group #1 pages 101-103 Co-editor: Kevin Petterson Group Members: Pat Hodge, Michael Clarke, Andrew Whalen

- All matter is composed of atoms - Not visible to the naked eye Many scientists have argued over which model would be correct and this is a short video showing the different models: []
 * __EARLY MODELS OF THE ATOM__**
 * Atom-** The smallest particle of an element that retains its identity in a chemical reaction.

Dalton's Atomic Theory (Michael Clarke)
Democritus first suggested the existence of the atom but it took almost two millennia before the atom was placed on a solid foothold as a fundamental chemical object by John Dalton (1766-1844). Although two centuries old, Dalton's atomic theory remains valid in modern chemical thought. 1) All matter is made of atoms. Atoms are indivisible and indestructible. 2) All atoms of a given element are identical in mass and properties 3) Compounds are formed by a combination of two or more different kinds of atoms. 4) A chemical reaction is a**//rearrangement//** of atoms. || Modern atomic theory is, of course, a little more involved than Dalton's theory but the essence of Dalton's theory remains valid. Today we know that atoms can be destroyed via nuclear reactions but not by chemical reactions. Also, there are different kinds of atoms (differing by their masses) within an element that are known as "isotopes", but isotopes of an element have the same chemical properties. Many heretofore unexplained chemical phenomena were quickly explained by Dalton with his theory. Dalton's theory quickly became the theoretical foundation in chemistry.
 * || **__Dalton's Atomic Theory__**

media type="youtube" key="IdSUqsq1yY8" height="315" width="560"

Democritus's Atomic Philosophy: (Kevin Petterson) - He was among the first to suggest the existence of atoms - Democritus believed that atoms were indivisible and indestructable - His ideas didn't agree with later scientific theory, because they did not explain chemical behavior - His experiments also lacked chemical support. - The experiments were not based on the scientific method __Sizing Up the Atom:__ (Andrew Whalen)
 * A coin the size of a penny and composed of pure copper (Cu) illustrates Dalton's comcept of the atom
 * Copper atoms are very small
 * A pure copper coin the size of a penny contains about 2.4x10^22 atoms[[image:http://faculty.virginia.edu/consciousness/images/iron%20atoms%20electrons.jpg width="500" height="400" align="right"]]
 * By comparison, Earth's population is only about 6x10^9 people
 * There are about 4x10^12 times as many atoms in the coin as there are people on Earth
 * If you could line up 100,000,000 copper atoms side by side, they would produce a line only 1 cm long!
 * The radii of most atoms fall within the range of 5x10^-11 m to 2x10^-10 m
 * Despite their small size, individual atoms are observable with instruments such as scanning tunneling microscopes
 * Individual atoms can even be moved around and arranged in patterns
 * The ability to move individual atoms holds future promise for the creation of atomic-sized electronic devices, such as circuits and computer chips
 * This atomic scale, or "nanoscale," technology could become essential to future applications in medicine, communications, solar energy, and space exploration

Group #2 pages 104-107 Co-editor: Joe Geraghty Group Members: Nate Gallishaw, Andy Marcotte, Sean Doherty

Electrons: Pages 104-105 (Nathaniel Gallishaw) An **electron** is a negatively charged subatomic particle. Thomson experimented by passing an electric current through gases at low pressure. The gases were sealed inside glass tubes, which had metal disks at each end. These disks were called electrodes. One was positively charged (anode) and one was negatively charged (cathode). The **cathode ray** was the glowing beam that resulted. The cathode ray can be deflected by a magnet. It can also be deflected by the metal plates. A plate with a positive charge attracts the ray and a plate with a negative charge repels the ray. Because he knew that opposite charges attract and like charges repel, Thomson hypothesized that the cathode ray was composed of a stream of negatively charged particles moving at high speed. These particles were electrons, but he had originally called them corpuscles. Thomson discovered that the ratio of the charge of an electron to its mass was constant. This ratio did not depend on the type of gas in the cathode ray tube or the type of metal in the electrodes, so Thomson was able to conclude that electrons existed in the atoms of all elements. An electron has one unit of negative charge and has a mass that is 1/1840 the mass of a hydrogen atom.
 * J. J. Thomson** (1856-1940) was an English physicist that discovered the electron in 1897.
 * Robert A. Millikan** (1868-1953) was an American physicist that found the quantity of the charge carried by the electrons. He also calculated the mass of the electron using his own results and the results of Thomson. The values he had determined for electron charge and mass in 1916 are close to the accepted values in use today.

Group #3 pages 110-116 Co-editor: Freddy Dwyer Group Members: Drew Humphrey, Matt Moshella

Pages 115 - 116: Matt Moschella This section focuses on the atomic mass and atomic mass unit (amu) So we know what mass is, the amount of matter an object contains. __**Atomic mass**__ is a weighted average mass of the atoms in a naturally occurring sample of an element. Ok so next important part is what gives the atom mass? atoms are made up of protons neutrons and electrons. Electrons are small, even compared to neutrons and protons, so they are impractical to use when measuring atomic mass. So yeah the atoms protons + its neutrons = the mass of the atom. Atomic mass is used for the elements. the scientists that determined the atomic mass wanted as little era as possible. to get accuracy many many many atoms of the same pure substance were looked at and individually given a mass. then all of the masses were averaged together to get the final atomic mass that is now on the periodic table. keep in mind that each atom is slightly different in mass and size so an average mass was a good idea. this average thing is put into fancy words and called relative abundance. to determine the atomic mass of some element, you need to know: the number of stable isotopes of the element, the mass of each isotope, and the natural percent abundance of each isotope. to calculate the atomic mass of an element, multiply the mass of each isotope by its natural abundance, expressed as a decimal, and then add the products. i think this video does a good job explaining how to find the atomic mass of an element.[|Atomic Mass Vidio] Pages 112-113- Freddy Dwyer It is crucial to be able to determine the composition of atoms. To do this you must determine the number of protons, electrons, and neutrons are in the atom. b.) Neon has an atomic number of 10 and a mass number of 20 c.) Sodium has an atomic number of 11 and a mass number of 23 now we know that the number of electrons is the same as the atomic numbers and the protons is the same. In order to determine the number of neutrons you must subtract the number of electrons (protons) from the mass number. So the number of neutrons in beryllium is 5, the number of neutrons in neon is 10 and lastly, sodium, is 12. __**ISOTOPES-**__ Atoms that have the same number of protons but differebt numbers of neutrons. Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. Despite these differences, isotopes are chemically alike because they have identical numbers of protons and electrons, which are the subatomic particles responsible for chemical behavior.There are 3 known isotopes of hydrogen has one proton in its nucleus. The most common hydrogen isotope has no neutrons. It has a mass number of 1 and is called hydrogen-1. the second has one neutron and i scalled hydrogen-2 and has a mass number of 2 it is called deuterium. the third isotope has 2 neutrons and has a mass number of 3 and is called hydrogen-3 or tritium.
 * __EXAMPLE-__** a.) Beryllium has an atomic number of 4 and a mass number of 9

Group #4 pages 127-132 Co-editor: Lauren Rossi Group Members: Alex Ortiz, Lindsey Trafford, Emily Healy


 * __THE BOHR MODEL-__** alex ortiz
 * __Niels Bohr-__** was a young danish physicist that believed Rutherford's model of a atom needed improvement.

- In 1913 he changed the model: the model included newer discoveries about how the energy of an atom changes when it absorbs or emits light. - Bohr proposed that an electron is found only in specific circular paths, or orbits, around the nucleus

- lowest rung= lowest energy level - to move from one energy level to another an electron must gain or lose just the right amount of energy. - energy levels in an atom are not equally spaced. - the higher the energy level occupies by an electron, the less energy it takes to move from that energy level to the next higher energy level - It still failed in many ways by: - explaining the energies absorbed - emitting by atoms with more than one electron media type="youtube" key="Ic8OnvRonb0" width="425" height="350" align="right" __**Quantum Mechanical Model**__ __**Atomic Orbitals- Emily Healy**__
 * energy levels-** fixed energies an electron can have (like rungs of a ladder)
 * quantum-** of energy is the amount of energy level to move an electron from one energy level to another energy level.
 * higher energy levels are closer together.
 * in 1926, Erwin Schrödinger, an Austrian physicist, used the equations and results of other scientists to devise and solve a mathematical equation describing the behavior of the electron in a hydrogen atom
 * the quantum mechanical model comes from the mathematical solutions to the Schrödinger equation
 * quantum mechanical model of the atom restricts the energy of electrons to certian values much like the Bohr model- unlike the Bohr model, however, this model does not involve an exact path the electron takes around the nucleus
 * it determines the the allowed energies an electron can have and how likely it is to find the electron in various locations aroun the nucleus- this is described by probability
 * probability of finding an electron within a certain volume of space surrounding the nucleus can be represented as a fuzzy cloud...the cloud is more dense where the probability of finding the electron is higher, and it is less dense where the probability of finding the electron is low
 * though it is unclear where the cloud eneds, there is at least a slight chance of finding the electron at a considerable distance from the nucleus
 * An atomic orbital is often thought of as a region of space in which there is a high probablility of finding an electron.
 * the energy levels of electrons in the quantum mechanical model are labeled by principal quantum numbers (n).
 * each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found.
 * s orbitals are spherical.
 * p orbitals are dumbbell-shaped.
 * d orbitals have a clover leaf shape.
 * the lowest principal energy level (n=1) has only one sublevel.
 * the second principal energy level (n=2) has two sublevels.
 * the third principal energy level (n=3) has three sublevels.
 * the fourth principal energy level (n=4) has four sublevels.

Group #5 pages 133-136 Co-editor: Emily Mills Group Members: Sejal Batra, Abby Williams, Monika Maczuga

__**ELECTRON ARRANGEMENT IN ATOMS ~ Emily Mills **__
 * Energy and stability play an important role in determining how electrons are configured in an atom.

**__Electron Configuration__**
//* Electron configurations// are the ways in which electrons are arranged in various orbitals around the nuclei of atoms.
 * Three rules that tell you how to find the electron configuration of atoms:
 * 1) The Aufbau principle
 * 2) The Pauli exclusion principle
 * 3) Hund's rule

**Aufbau Principle[[image:http://0.tqn.com/d/chemistry/1/0/j/g/econfiguration.jpg width="161" height="342" align="right"]]**

 * Electrons occupy the orbitals of lowest energy first.
 * The orbital for any sublevel of a principle energylevel are always equal energy levels.
 * The Aufbau diagram shows the energy levels of the various atomic orbitals. The orbitals of greater energy are higher on the diagram.

Page 135:Abby Williams


 * an oxygen atom has 8 electrons
 * the lowest energy orbital (1s) contains 1 electron and a second of opposite spin
 * next orbital (2s) has 1 electron and a second of opposite spin
 * one electron then occupies three 2p orbitals of equal energy, while the last pairs with one of the electron of the 2p orbital
 * the other 2p orbitals are only half full with one electron


 * Exceptional Electron Configuration **- Monika Maczuga
 * You can obtain correct electron configuration for the elements up to vanadium (atomic number 23) by following the aufbau diagram for orbital filling.
 * Exceptions to the aufbau principal are due to subtle electron-electron interactions in orbitals with very similar engeries.
 * __**Copper**__ and **__chromium__** have electron configuration that are an exception to the aufbau principle.

If you were to continue the aufbau fashion, you would assign chromium and copper the following incorrect configurations:
 * Cr ** 1s2 2s2 2p6 3s2 3p6 3d4 4s2 **/ Cu** 1s2 2s2 2p6 3s2 3p6 3d9 4s2

The correct electron configurations are as follows:

(The above arrangements give chromium a half-filled //d// sublevel and copper a filled //d// sublevel) ü Some actual electron configurations differ from those assigned using the aufbau principal because half-filled sublevels are not as stable as filled sublevels, but they are more stable than other configurations.
 * Cr ** 1s2 2s2 2p6 3s2 3p6 3d5 4s1 **/ Cu** 1s2 2s2 2p6 3s2 3p6 3d10 4s1
 * Filled energy levels are more stable than partially filled sublevels
 * This tendency overcomes the small difference between the energies of the 3d and 4s sublevels in copper and chromium.

//How are the energy differences and exceptions to the aufbau principal related?//

For a video explaining these exceptions please see: [|http://www.youtube.com/watch?v=7MuSuySi0Q0&feature=player_detailpage#t=2s]

Group #6 pages 138-145 Co-editor: Kendall Lavin-Parsons Group Members: Kelsey Persechini, Katie Manis __Atomic Spectra:__ __An Explanation of Atomic Spectra:__
 * passing an electric current through a gas in a neon tube energizes the electrons of the atoms of the gas, and causes them to omit light.
 * when atoms absorb energy electrons move ino higher energy levels and these electrons lose energy by emitting light when they return to lower energy levels.
 * each specific frequency of visible light emitted corresponds to a particular color.
 * when the light passes the frequencies of light emitted by an element seperate into discrete lines to give the **atomic emission spectrum** by the element.
 * each discrete line in an emission spectrum corresponds to one exact frequency of light emitted by the atom.
 * no two elements have the same emission spectrum.
 * In the Bohr model the lone electron in the hydrogen can have only specific energies.
 * the lowest possible energy is called **g**r**ound state.**
 * in the ground state the electrons principal quantum number is 1.
 * the light emitted from an electron moving from a higher to a lower energy level has a frequency directly proportional to energy change of the electron.
 * lyman series are the lines at the UV end of the hydrogen spectrum. n=1
 * the lines in the visible spectrum are the Balmer series. n=2
 * paschen series are the transition from n=2 to n=3.
 * n=4 and n=5 exists too.
 * the quantum mechanical model is based on the description of the motion of material objects as waves.

__** Light **__ As the wave lenght increases frequency decreases electomagnetic spectrum ~ Kendall Lavin-Parsons
 * light consists of waves not particles that was previouslt suggested by Isaac Newton
 * each complete wave cycle starts at zero
 * to ge to its highest point it raises above zero
 * to get to its lowest point it falls below zero
 * AMPLITUDE= the waves height from zero to the crest
 * WAVELENGTH= the distance between the crests
 * FREQUENCY= the number of wave cycles to pass a given point per unit of time
 * frequency is usually measured with cycles per second
 * HERTZ (Hz)= the SI unit 0f cycles per second