Atomic+Structure+and+Electrons+in+Atoms

Atomic Structure and Electrons in Atoms. Period D Editor: Zac Boulerice Co-Editors: -Henry Dodge -Kevin McAllister -Gabe Hannawi Introduction To the Atom The Atom is everywhere. When we talk about matter itself we talk about matter being everywhere. In the air, the ground, and in us. We just can’t get away from this thing called matter. But what makes up matter? Atoms. Small, tiny, subatomic particles which can’t even be seen underneath a microscope. These are the things you will be learning about in this wiki.

Group 1: - Nick Romero (101) - Lauren Sachs (102) - Henry Dodge (103) - Ian Travis (104-105) - Alex Trombetta (106) - Erin Cropanese (107)

Pg 101. Early Models of the Atom by Nick Romero

Vocabulary.... ATOM: the smallest particle of an element that retains its identity in a chemical reaction. - not visible to the naked eye - made up of subatomic particles Early predictions... Many early scientists predicted the existence of the atom but did not posses the necessary equipment to actually discover it.

DEMOCRITUS'S ATOMIC PHILOSOPHY This greek philosopher proposed the idea of an atom four hundred years before Christ's birth. He thought that the atom was -indivisible -indestructible Atom comes from the word atomos, meaning "indivisible"
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[|Democritus Video]

Pg 102. Dalton's Atomic Theory By: Lauren Sachs · Modern process of discovery about atoms began with John Dalton · By using experimental methods, Dalton transformed Democritus's ideas on atoms into a scientific theory · Dalton studied the ratios in which elements combine to form Dalton's atomic theory 1. All elements/matter are composed of tiny particles called atoms 2. Atoms cannot be subdivided*  3. Atoms of the same element are identical*. 4. The elements of one element have different properties from those of any other element. 5. Atoms of different elements can physically mix together or can chemically combine in simple whole-number ratios to form compounds 6. Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element, however, are never changed into atoms of another element as a result of a chemical reaction (can't be created or destroyed). *=false (can be subdivided through nuclear reactions; isotopes are not identical)

Pg 103. Sizing up the Atom By: Henry Dodge

-John Dalton had an idea that if elements were repeatedly divided, eventually there would come a point when it could be divided no further.

-This stage, or final particle, is called an atom.

-Copper atoms are very small.

-Most atoms have a radius between 5 x 10-11 m and 2 x 10-10 m.

-Despite their small size, individual atoms can be observed with instruments such as scanning tunneling microscopes.

-With this technology, individual atoms can even be moved around and arranged in patterns.

-This atomic-scale could hold breakthrough promises in the areas of medicine, communications, solar energy, and space exploration.



Pg. 104-105. 4.2 Structure of the Nuclear Atom

By: Ian Travis

Though much of Dalton's theory is accepted today, there is one major change in modern chemistry: the atom can be divided into even simpler "subatomic" particles. In chemistry, the main subatomic particles that are studied are the electrons, protons, and neutrons

J. J. Thomson the English physicist discovered the electron through experimentation. A glass container containing low-pressure gas was fitted with metal disks called electrodes at each end. The electrodes were connected to a source of electricity. One disk became positively charged and was called the anode. The other became negatively charged and was called the cathode. Between the two electrodes, a glowing beam of electrons called a cathode ray appeared, traveling from the cathode to the anode.

Electron- A negatively charged subatomic particle that orbits the nucleus of an atom.

Cathode Ray- Beam of electrons that will travel from cathode to anode.

Cathode rays can be deflected by magnets are electrically charged plates. Negatively charged plates will repel a cathode ray, and a positively charged plate will attract it.

Thomson was able to infer and hypothesize that a cathode ray was a stream of negatively charged particles moving at high speed.

Pg 106. Protons, Neutrons, and the Atomic Nucleus By:Alex Trombetta Protons and Neutrons What is left after a hydrogen atom loses an electron? This question can be answered with four simple ideas about electrical charges.

1.) Atoms have no net electric charge; they are electrically neutral. 2.) Electric charges are carried by particles of matter 3.)Electric charges always exist in whole number multiples of a single basic unit; that is there are no fractions of charges. 4.) When a given number of negatively charged particles combines with an equal number of positively charged particles, an electrically neutral particle is formed.

Using this information you can determine that one unit of positive charge remains after the losing of one electron in a hydrogen atom.

Evidence of this positively charged particle was found in 1886 by Eugen Goldstein. To do this he observed a cathode-ray and found rays traveling in the opposite directions to the actual cathode rays. He determined that they ate positively charged which are now called protons with a mass 1840 times that of an electron. ￼

Later another subatomic particle was discovered by James Chadwick. This particle was the neutron which is a subatomic particle with no charge but with a mass similar to that of a proton. ￼

Particle Symbol Charge Relative Mass Actual Mass Electron e- 1- 1/840 9.11x10^-28 Proton p+ 1+ 1 1.67x10^-24 Nuetron n0 0 1 1.67x10^-24

The Atomic Nucleus When first discovered the configuration of the subatomic particles in an atom was unsure and scientists became curious and had trouble because of the small size of these particles. Originally thought by many, including JJ Thompson, electrons were evenly distributed in the atom filled uniformly with protons. The “plum pudding model” created by Thomson proved “short-lived” after the successful work of Thomson’s student Ernest Rutherford. <span style="font-family: 'Arial','sans-serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt;">￼ (JJ Thomson: The Electron) (James Chadwick: The Neutron) (Eugen Goldstein: The Proton)

Pg 107 by Erin Cropanese

Rutherford's Gold-Foil Experiment [] · 1911 he decided to test the current theory of atomic structure · he used massive alpha particles and shot a narrow beam of them at a very thin gold foil · he expected them to pass through easily with only slight reflection because of positive charge though to be spread out in the gold atoms · the results were that the great majority of the alpha particles passed straight through the gold atoms without deflection · also, a small fraction of them bounced straight back toward the source and the others bounced off at large angles

Nucleus: the tiny central core of an atom and is composed of protons and neutrons
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 * Rutherford Atomic Model
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">Rutherford proposed a new theory of the atom
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">he proposed that the atom is mostly empty space (explains the lack of deflection of most of the particles)
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">he concluded that all positive charge and almost all the mss are concentrated in a small region that has enough positive charge to account for the deflection of some of the alpha particles-> called the NUCLEUS
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">he proposed that the atom is mostly empty space (explains the lack of deflection of most of the particles)
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">he concluded that all positive charge and almost all the mss are concentrated in a small region that has enough positive charge to account for the deflection of some of the alpha particles-> called the NUCLEUS
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">atom model known as the nuclear atom
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">protons and neutrons located in the nucleus
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">the electrons are distributed around the nucleus and occupy almost all the volume of the atom
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">nucleus is tiny compared to the whole atom
 * <span style="font-family: Calibri; font-size: 12pt; line-height: normal; margin: 0in 0in 10pt; tabstops: list .5in; tabstops: list .5in;">his model turned out to be incomplete and has since been revised to explain chemical properties of elements

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Group 2: <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Jess Mahoney (110-111) <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Andrew Ware (112-113) <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Kevin McAllister (114-116) (Co-editor) <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Joe Hatch (127-129) <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Chris Hughes (130-132)

Pg 110- 111. 4.3 Distinguishing Among Atoms By Jess Mahoney · Atomic Number o Definition- the number of protons in the nucleus of an atom of an element o Protons and neutrons are in the nucleus, electrons surround the nucleus o Identifies an element o An atom is electrically neutral when the number of protons= the number of electrons o Elements are different because they contain different numbers of protons.
 * <span style="display: block; line-height: normal; margin: 0in 0in 0pt; mso-add-space: auto; mso-element-anchor-horizontal: margin; mso-element-anchor-vertical: paragraph; mso-element-frame-hspace: 9.0pt; mso-element-top: 6.8pt; mso-element-wrap: around; mso-element: frame; mso-height-rule: exactly; text-align: center;"> Atoms of the First Five Elements ||
 * Name || Symbol || Atomic # || Protons || Neutrons || Mass # || # of electrons ||
 * Hydrogen || H || 1 || 1 || 0 || 1 || 1 ||
 * Helium || He || 2 || 2 || 2 || 4 || 2 ||
 * Lithium || Li || 3 || 3 || 4 || 7 || 3 ||
 * Beryllium || Be || 4 || 4 || 5 || 9 || 4 ||
 * Boron || B || 5 || 5 || 6 || 11 || 5 ||

· Mass Number o Definition: the total number of protons and neutrons in an atom o Most of the mass of an atom is concentrated in its nucleus o Also depends on protons and neutrons (amount) § Example: helium atom has 2 protons and 2 neutrons. · Mass number=4 Carbon atom has 6 protons and 6 neutrons · Mass number=12 o With atomic and mass numbers you can determine an element’s composition. o “The number of neutrons in an atom is the difference between the mass number and the atomic number.” o Number of neutrons=mass number – atomic number <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt; text-indent: -0.25in;"> <span style="display: block; height: 17.95pt; left: 0px; margin-left: 306pt; margin-top: 9pt; position: absolute; text-align: left; width: 135pt; z-index: 8;">  Pg 112-113. Isotopes By: Andrew Ware
 * <span style="display: block; padding-bottom: 4.35pt; padding-left: 7.95pt; padding-right: 7.95pt; padding-top: 4.35pt;"> By Jess Mahoney ||

-Isotopes are atoms that have the same number of protons but different numbers of neutrons. -Because isotopes of an element have different numbers of neutrons, they also have different mass numbers. -despite differences, isotopes are chemically alike because they have identical numbers of protons and electrons. -There are 3 known isotopes of hydrogen. Each isotope of hydrogen has one proton in its nucleus. -The most common hydrogen isotope has a mass number of 1 (hydrogen-1 or hydrogen). The second isotope has one neutron and a mass of 2 (hydrogen-2 or deuterium). The third is called hydrogen-3 or tritium.


 * <span style="display: block; padding-bottom: 4.35pt; padding-left: 7.95pt; padding-right: 7.95pt; padding-top: 4.35pt;"> By Andrew Ware ||

<span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt; tab-stops: 55.5pt; tabstops: 55.5pt;">Pgs. 114-116. Atomic Mass <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt; tab-stops: 55.5pt; tabstops: 55.5pt;">By: Kevin McAllister o The Masses of protons and neutrons (1.67*10 to the -24) are extremely small o The Mass of an Electron is even smaller (9.11*10 to the -28) o Even the largest atom’s mass is miniscule o Masses can be determined however using a Mass Spectrometer BUT o The Mass Spectrometer gives small numbers that are a pain to work with SO <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt; tab-stops: 55.5pt; tabstops: 55.5pt;">The Atomic Mass Unit o Instead of using the Mass Spectrometer the amu was invented by using a carbon 12 atom as a reference isotope o 1/12 of the mass of a Carbon 12 atom is the definition of how much one amu is   o One Carbon 12 atom has 6 protons and six neutrons meaning that one of each is equal to 1amu <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt; tab-stops: 55.5pt; tabstops: 55.5pt;">In Nature o In Nature the isotopes usually occur as a mixture o Because of this, the masses are not usually whole numbers <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt; tab-stops: 55.5pt; tabstops: 55.5pt;">The Atomic Mass o To find the atomic mass you must know the natural percent of abundance o If you average two isotopes you might not get the actual atomic mass o The Atomic mass of an element is a weighted average mass of the atoms in a naturally occurring sample o This way the masses are more correct because of how they reflect the mass and abundance of isotopes <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt; tab-stops: 55.5pt; tabstops: 55.5pt;">Key Concepts o In order to find the atomic mass of an element, multiply the mass of an isotope by it’s natural abundance (which is a decimal) and then add the products

o Use [] to watch a video on figuring out the Atomic Mass o Check out [] for an explanation on this topic as well as sample problems for practicing.

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 * <span style="display: block; padding-bottom: 4.35pt; padding-left: 7.95pt; padding-right: 7.95pt; padding-top: 4.35pt;"> By: Kevin McAllister ||

Pg 127-129. 5.1 Models of the Atom By: Joe Hatch The Development of Atomic Models (p127) <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Rutherford’s model could not explain the chemical properties of elements <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt 0.5in; text-indent: -0.25in;">- Explaining what leads to these chemical properties requires a better model that better describes the behavior of electrons The Bohr Model (p128-129) <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Niels Bohr was a student of Rutherford <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Bohr’s model included the electrons orbiting around the nucleus <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Bohr was the first to include energy levels, or the fixed energies an electron can have <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- To jump from one energy level to another, the electron must gain or lose the right amount of energy <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- Quantum- is the amount of energy required to move an electron from one energy level to the next <span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 0pt 0.5in; text-indent: -0.25in;">- The amount of energy an electron gains or loses in an atom is not always the same (the higher energy levels are closer together) - The higher the energy level occupied by an electron, the less energy it takes to move from that energy level to the next higher energy level

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By Joe Hatch  ||
 * <span style="display: block; padding-bottom: 4.35pt; padding-left: 7.95pt; padding-right: 7.95pt; padding-top: 4.35pt;"> Bohr model of Iodine

<span style="font-family: 'Calibri','sans-serif'; margin: 0in 0in 0pt; tab-stops: 9.0pt; tabstops: 9.0pt;">Pgs. 130-132 <span style="font-family: 'Calibri','sans-serif'; margin: 0in 0in 0pt; tab-stops: 9.0pt; tabstops: 9.0pt;">by: Chris Hughes

<span style="font-family: 'Calibri','sans-serif'; margin: 0in 0in 0pt; tab-stops: 9.0pt; tabstops: 9.0pt;">The Quantum Mechanical Model • <span style="font-family: 'Calibri','sans-serif';">In 1926, the Austrian physicist Erwin Schrödinger devised and solved a mathematical equation describing the behavior of the electron in a hydrogen atom. • <span style="font-family: 'Calibri','sans-serif';">The Quantum Mechanical Model is the modern description of the electrons in atoms and comes from the mathematical solutions to the Schrödinger equation. • <span style="font-family: 'Calibri','sans-serif';">Unlike the Bohr model, the quantum mechanical model does not involve the exact path the electron takes around the nucleus. • <span style="font-family: 'Calibri','sans-serif';">The quantum mechanical model determines the allowed energies an electron can have and how likely it is to find the electron in various locations around the nucleus. • <span style="display: block; height: 198pt; left: 0px; margin-left: 331.9pt; margin-top: 49.4pt; position: absolute; text-align: left; visibility: visible; width: 175.5pt; z-index: 4;"> <span style="display: block; height: 187.5pt; left: 0px; margin-left: 4.95pt; margin-top: 2.85pt; position: absolute; text-align: left; visibility: visible; width: 118.5pt; z-index: 3;"> <span style="font-family: 'Calibri','sans-serif';">The locations of electrons are described through probability, and are similar to the motion of a rotating propeller blade. In the picture, we know the blade is there, but the image does not tell us the exact location of the blade. • <span style="font-family: 'Calibri','sans-serif';">The quantum mechanical model can be represented as a fuzzy cloud. • <span style="font-family: 'Calibri','sans-serif';">Even though we cannot determine the exact location of the electron using the model, the cloud is more dense where the probability of finding the electron is high • <span style="font-family: 'Calibri','sans-serif';">To visualize an electron probability cloud, imagine you could mold a sack around the cloud so that the electron was inside the sack 90% of the time. • <span style="font-family: 'Calibri','sans-serif';">Solving the Schrödinger equation gives the energies an electron can have, which are energy levels. Each energy level, through the equation, leads to a mathematical expression called an atomic orbital.

<span style="font-family: 'Calibri','sans-serif'; margin: 0in 0in 0pt; tab-stops: 9.0pt; tabstops: 9.0pt;">Atomic Orbitals • <span style="font-family: 'Calibri','sans-serif';">An atomic orbital is a mathematical expression describing the probability of finding an electron at various locations; usually represented by the region of space around the nucleus where there is a high probability of finding an electron. • <span style="font-family: 'Calibri','sans-serif';">There are principal energy levels which are labeled principle quantum numbers, and assigned values of n= 1,2,3,4 and so forth. • <span style="font-family: 'Calibri','sans-serif';">For each principal level there may be several orbits with different shapes and sizes and at different energy levels. These variations form sublevels.

• <span style="font-family: 'Calibri','sans-serif';">Each energy sublevel corresponds to an orbital of a different shape, which describes where the electron is likely to be found. • <span style="font-family: 'Calibri','sans-serif';">The colorful image describes the s, p, d, and f orbitals. The only ones shown in the section are the s and p orbitals. • <span style="font-family: 'Calibri','sans-serif';">Each one of the principal levels, as shown in the chart, contains each of the orbitals shapes before it, just at the higher energy level. • <span style="display: block; height: 26.25pt; left: 0px; margin-left: 44pt; margin-top: 339.75pt; position: absolute; text-align: left; width: 398.6pt; z-index: 6;"> <span style="font-family: 'Calibri','sans-serif';">The maximum number of electrons an orbital can hold is found by the formula 2n2, where n is the principal quantum number (level number). So the first energy level can contain maximum of 2 electrons, the second 8 electrons, the third 18 electrons, and the fourth 32 electrons.
 * // http://upload.wikimedia.org/wikipedia/commons/4/4a/Single_electron_orbitals.jpg // ||

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Group 3: Gabriel Hannawi (co-editor) -pg. 133-134     Tess Murphy –pg. 135-136     Cassie Naimie –pg. 138-140     Sahana Nazeer –pg 141-143 Caity Vogt –pg 144-145

Gabriel Hannawi Pg 133-134

Electron Configurations: · The orbital and spin arrangement of an atom’s electrons, specifying the quantum numbers of the atom’s electrons in a given state · The Aufbau Principle states the way that electrons fill out the orbitals - According to the principle, electrons fill orbitals starting at the lowest available energy states before filling higher states (e.g. 1s before 2s).

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· The number of electrons that can fill up an orbital is limited by the Pauli exclusion principle - The Pauli Exclusion Principle states that no two electrons in an atom can be in the same state or configuration at the same time, meaning that at most, there are at most two electrons per orbital in an atom · Hund’s rule states that electrons occupy orbitals of the same energy in a way that makes the number of electrons with the same spin direction as large as possible. - The greater total spin state usually makes the resulting atom more stable - If two or more orbitals of equal energy are available, electrons will occupy them singly before filling them in pairs.

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Tess Murphy Pg. 135-136

Oxygen atom- ·  contains eight electrons ·  the orbital of lowest energy (1s) has one electron, then a second electron with the opposite spin ·  the next orbital (2s) has one electron, then a second electron with the opposite spin too ·  then, one electron occupies each of the three 2p orbitals of energy ·  the electron left pairs with an electron that occupies one of the 2p orbitals · the other two 2p orbitals are only half full, with only one electron each <span style="font-family: 'Tahoma','sans-serif'; font-size: 12pt;">￼ Easy method for showing electron configuration of an atom-

·  write the energy level and the symbol for every sublevel occupied by an electron ·  show the number of electrons occupying that sublevel with a superscript -  Hydrogen has one electron in a 1s orbital = 1s1 -  Helium has two electrons in a 1s orbital = 1s2 -  Oxygen has two electrons in a 1s orbital, two electrons in 2s orbital, and four   electrons in 2p orbitals = 1s22s22p4 -  the sum of the superscripts = the number of electrons in the atom -  the sublevels within the same principal energy level are written together -  not always the same order has the aufbau diagram - ex. The 3d sublevel is written before the 4s sublevel, even though the aufbau diagram has the 2s sublevel to have lower energy Video for help with electron configuration: []

<span style="font-family: Calibri; font-size: 12pt; height: 101.25pt; visibility: visible; width: 150.75pt;"> Phosphorus is an element used in matches. Its electron configuration is 1s22s22p63s23p3.

Practice problems: Write the electron configurations for each atom- a. carbon b. argon c. nickel

Write the electron configurations for each atom. How many unpaired electrons does each atom have? a. boron b. silicon

Exceptional Electron Configurations

·  Copper’s electron configuration is an exception to the aufbau principle ·  correct electron configurations for the elements up to vanadium can be found on the Aufbau diagram for orbital filling ·  continuing past that would give you incorrect configurations for chromium and copper ·  filled energy sublevels are more stable than partially filled sublevels ·  some actual electron configurations are different from the ones assigned using the Aufbau principle, because half-filled sublevels are less stable than filled sublevels, but more stable than other configurations · exceptions happen because of subtle electron- electron interactions in orbitals that have really similar energies at higher principle quantum numbers, some sublevels have even smaller energy differences than in chromium and copper there are other exceptions to the Aufbau principle ·  it’s good to know that exceptions occur · it’s more important to know the basic rules for figuring out electron configurations in the many cases where the Aufbau principle does apply

Cassie Naimie Pg 138-140

-The quantum mechanical theory grew out of a study of light -Isaac Newton tried to explain that light was made out of particles -By 1900, scientists figured out that light was actually made from waves -A complete cycle of a wave starts at zero, increases to highest value, passes through zero to its lowest value, and finally ends back at zero.

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-The amplitude of a wave is the wave’s height from zero to the crest -The wavelength (represented by Greek symbol λ) is the distance between the crests -The frequency (represented by Greek symbol v) is the number of wave cycles to pass a given point per unit of time -The product of frequency and wavelength always equals a constant (c) the speed of light

c= λv

-The wavelength and frequency of light are inversely proportionate to each other; as wavelength increases the frequency decreases

-The wave model shows light consisting of electromagnetic waves -Electromagnetic radiation includes radio waves, microwaves, infrared waves, visible light, ultraviolet waves, X-rays, and gamma waves -All electromagnetic waves travel at a speed of 2.998 x 108 m/s -When sunlight passes through a prism, the different frequencies separate into a spectrum of colors -A rainbow each droplet of water acts as a prism to produce a spectrum



Calculating the Wavelength of Light

Problem: calculate the wavelength of the yellow light emitted by the sodium lamp if the frequency of radiation is 5.10 x 1014 Hz (5.10 x 1014 /s).


 * 1) <span style="font-family: Calibri; font-size: 12pt; margin: 0in 0in 0pt; tabstops: list .5in; tabstops: list .5in;">Analyze- List the known and unknown

Known: -Frequency (v) = 5.10 x 1014 /s -Constant (c) = 2.998 x 108 m/s

Unknown: -wavelength (λ) = ?m

<span style="font-family: Calibri; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt; text-indent: 0.5in;">Solve the equation c= λv for (λ)
 * 1) <span style="font-family: Calibri; font-size: 12pt; margin: 0in 0in 0pt; tabstops: list .5in; tabstops: list .5in;">Calculate-

λ = c/v

Substitute and solve

λ = c/v = 2.998 x 108 m/s = 5.88 x 10-7 m 5.10 x 1014 /s

Yes, the magnitude of the frequency is much larger than the numerical value of the speed of light, so the answer should be much less than 1.
 * 1) <span style="font-family: Calibri; font-size: 12pt; margin: 0in 0in 0pt; tabstops: list .5in; tabstops: list .5in;">Evaluate: Does the answer make sense?

For a Frequency and Wavelength Calculator check out this site: []

Section 5.3 – Physics and the Quantum Mechanical Model Sahana Nazeer – p. 141-143 Atomic Spectra Key: When atoms absorb energy, electrons move into higher energy levels. Those electrons lose energy by emitting light when they return to lower energy levels. Hence passing an electric current through gas in a neon tube causes the electrons to emit light. A prism separates light into the colors it contains. Light that is emitted by atoms consists of a mixture of only specific frequencies; each specific frequency of visible light emitted corresponds to a particular color.

Atomic Emission Spectrum – the pattern that is formed when light passes through a prism or diffraction grating to separate it into the different frequencies of light it contains frequencies of light emitted by an element separate into discrete lines on order to give the atomic emission spectrum of the element An emission spectrum of each element is like a person's fingerprint. Therefore, no two elements can have the same emission spectrum.

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An Explanation of Atomic Spectra Bohr's model explains: (1) why the emission spectrum of hydrogen consists of specific frequencies; (2) predicts specific values of these frequencies that agreed with experiment **Bohr's model has a lone electron in the hydrogen atom that can only have certain specific energies.

ground state – the lowest possible energy of the electron; the electron's principal quantum number (n) is 1. electronic transition - emission occurs in a single abrupt step

Key: The light emitted by an electron moving from a higher to a lower energy level has a frequency directly proportional to the energy change of the electron. Each transition produces a line of a specific frequency in the spectrum.

Lyman Series – lines at the ultraviolet end of hydrogen spectrum; “excited states'; n=1 Balmer Series – lines in the visible spectrum; n=2 Paschen Series- lines in the infared range; n=3 Spectral lines in each group become more closely spaced at increased values of n because energy levels become closer together.

Caitlyn Vogt Pages 114-115 Einstein- patent examiner in Bern Switzerland In 1905 he used Newton’s developed ideas and discovered that light could be described as quanta of energy Quanta behave as if they were particles Photons- light quanta Louis de Broglie- French graduate student In 1924 he examined the idea that because light behaves as waves and particles could the particles of matter behave like waves Yes: the wave like behavior of particles was matter waves This leads to the creation of a mathematical expression for the wavelength of a moving particle Clinton Davisson and Lester Germer- confirmed Louis de Broglie’s hypothesis within the next three years They scientists were studying the bombardment of metals with beams of electrons There were patterns created when the electrons were reflected from the metal Specifically the electrons were reflected like waves Davisson and De Broglie both received Nobel prizes for their work in the wave nature study The wavelike properties of beams of electrons are used when trying to enlarge objects Example: the electrons in an electron microscope have smaller wavelengths than light revealing a clearer image From De Broglie’s discoveries there was an equation and two theories created The equation predicts that all objects which move have wavelength and that the wavelength can only be observed if the mass of the object is small CLASSICAL MECHANICS: older theory which describes the motions of bodies much larger than atoms QUANTAM MECHANICS: newer theory which describes the motions of subatomic particles and atoms as waves Heisenberg uncertainty Principle- developed by German physicist Werner Heisenberg It states that it is impossible to know exactly the velocity and position of a particle at the same time For example: if you need to find a pair of keys in a dark room you use a flashlight and see the keys- it’s a done deal: therefore the principle doesn’t apply to the keys <span style="font-family: 'Calibri','sans-serif'; font-size: 12pt; height: 94.5pt; line-height: 115%; margin: 0in 0in 10pt 0.5in; visibility: visible; width: 185.25pt;"> The principle doesn’t apply to ordinary sized objects On the other hand though, if you used a flashlight to find an electron, the photon of light would affect the motion of the electron because of the electron’s small mass. Therefore the location of the electron would change which would affect its velocity and so on. A never ending cycle of not knowing velocity or location of an electron will continue for eternity; hence the Heisenberg Uncertainty Principle. “Quantum Mechanics” <span style="color: windowtext; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt 0.5in; text-decoration: none; textunderline: none;">[] This leads to the discovery of the concept of electron orbital and configuration.