the most stable arrangement of the nucleus and the electrons

8/14/2019 The Most Stable Arrangement of the Nucleus and the Electrons

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The most stable arrangement of the nucleus and the electrons inan atom is one for which the total energy of the atom (kinetic energy and potentialenergy) is at a minimum. When an atom is exposed to heat, light, or when itcollides with another particle, it may absorb additional energy.

Electromagnetic radiation: is most simply defined as light and as you know, not alllight is visible to the human eye.

Relative size of the wavelengths for the EM spectrum

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As the EM radiation travels through space, it creates both oscillating electric andmagnetic fields along the way. The two fields are perpendicular to one another.

Radiation is more than just the invisible stuff that nuclear bombs leave behind.Ultraviolet, microwave, infrared, FM radio signals, visible light; these are allconsidered radiation. The EM spectrum (above) refers to all of the different types ofradiation that exist. EM radiation is naturally transmitted by stars (including oursun), travels at the speed of light, and can vary in wavelengths from0.000000000001 meter (Gamma rays) to 10,000 m (television & radio)! In theory,the possible wavelengths extend in size to infinitely large and small.

Wave Properties of EM radiation:

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Imagine a buoy floating on water. As a boat passes by, the waves produced willcause the buoy to bob up and down. This wave is a periodic disturbance oroscillation that passes through space. A wave consists of repeating units calledcycles. The vertical motion of the buoy is caused by the passage of successivecrests and troughs as the waves move through the water.

The wave properties of electromagnetic radiation are described by two independent

variables.

Frequency: ( ) the number of cycles that pass a given point each second. This is

the vertical motion. The buoy will bob up and down times per second.

Wavelength: ( ) the distance between two successive points on the wave.Simply speaking, peak to peak or trough to trough is typically defined as thewavelength.

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Types of radiation are arranged in order by their wavelengths. The wavelength of awave is the distance between 2 consecutive points of a wave. Typically we saypeak to peak or trough to trough as these points are very easy to pick out on thewave. But one could look at the middle of the peak to the middle of the next peakand call that the wavelength. Radiation in the EM spectrum is made up of wavesthat contain an Electric field and a Magnetic field (see figure above). The frequencyof a wave may be defined by how often the wave passes a fixed point in space in 1second. For example, 1 peak/second = 1 Hz; 10 peaks/second = 10 Hz.

Units for frequency are sec-1 (orsec

1) often called Hertz (Hz). Remember we are

talking about cycles, so the units could also be labeled assec

cycles

Units for wavelength are meters orcycle

meters

The speed of the wave is the distance traveled per some period of time. Thus,

sec

meters

We can get those units by multiplying frequency x wavelength

Speed of a wave: x (units aresec

meters)

In a vacuum, the speed of any type of EM radiation is the same and is defined as2.9979 x 108meters/sec. This is defined as the speed of light and is defined as c which is a constant.

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c = x

2.9979 x 108meters/sec = x

There is an inverse relationship between frequency and wavelength. If we

rearrange the equation for the speed of light for and then again for , therelationship becomes more apparent.

= c/ = c/

If you are having trouble seeing the relationship, we can divide up the equationeven further:

= c x 1/ = c x 1/

As you can see above, if the wavelength is long (red and orange lines for example)the frequency (how many wavelengths pass a point in a certain period of time) islow. Large wavelengths thus mean low frequency. If the wavelength is short (blueand purple lines) then the number wavelengths that pass a particular point in acertain period of time is high. For example, if the above picture is a snapshot of 1second, the red line has 1.5 wavelengths (or thereabouts) passing a point in asecond. The purple line has 14 wavelengths passing in 1 second.

Large wavelength = small frequencySmall wavelength = High frequency

We can use this relationship as an answer check when we do light calculations!

Given: What is the frequency of a radar wave with a = 1.0 cm?

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Another term to be aware of is amplitude, which simply refers to the height of thewave. Waves with higher amplitude have higher intensities, and waves with loweramplitudes have weaker intensities but they represent the same species (e.g. thesame color of light) as their wavelength and frequency are the same.

It is important to know the visible spectrum: ROYGBIV (red, orange, yellow, green,blue, indigo, violet), and their relative wavelengths and frequencies. Red has thelongest wavelength and thus the lower frequency while violet has the shortestwavelength and thus the highest frequency. If asked, be sure to know the order ofthe colors and the relative frequencies!!

Light of a given wavelength travels at different speeds depending on the medium inwhich it is located: e.g. air, a vacuum, water, etc . .. When light passes from onemedia to another, the speed of the wave changes. The speed of light throughmedia such as water or glass is somewhat less than 2.9979 x 108m/sec. The changein speed is also accompanied by a change in direction thus it gets bent at someangle which depends on the media from which it came and the media which it nowentered. This is a phenomenon known as refraction. It does not happen to particles(e.g. think of a stone thrown into water, it follows a curved path, not a new bent linepath even though its speed did change).

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Refraction happens when a wave passes from one material to another of differentdensities. As the wave crosses the boundary, it bends. This kind of bending is calledrefraction. Waves travel at different speeds in different materials. The change inspeed is what causes the wave to bend.Waves do not refract if they cross the barrier at a perpendicular angle. Examine apencil in a glass straight up and down it appears no different but slanted we seethe refraction). But if they cross the barrier at any other angle, they will bend asthey go through it. One edge of the wave will slow down or speed up before theother edge does. That is why it appears broken.We see refraction if we look at something through a glass or water. Eye glasses andcontacts have lenses which refract light to correct vision problems.

Insight: How many have noticed that when you put an ice cube in water it developsswirls around it? If you stare at the ice cube in the water you will see these swirls.We know from looking in the CRC Handbook of Chemistry and Physics that waterhas different densities at different temperatures. Although the change might besmall, it is sufficient enough to change the path of the light that travels through ourice water. When you are driving down the highway in the summer and you stare atthe road ahead and see the oasis in front of you on the asphalt, the same thing ishappening. It is still air but it is air at a different temperature which has adifferent density. As the light travels through the air it gets refracted bent andwe see it in front of us as a mirage of liquid on the road!

When a wave is forced to go through a smaller opening, it actually bends around theopening in a process known as diffraction. This forms a semicircular wave on theother side of the opening. Particles behave very differently. When forced throughan smaller opening, some of the particles will hit the barricade and the others willpass through as if undeflected just like cars merging on an interstate.

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If waves travel through adjacent slits, the resulting circular waves interact with oneanother. Think about when two boats pass one another going in opposite directions.Both boats traveling through the water are making waves, as those waves come incontact with one another they can combine to form a larger wave (calledconstructive interference) or they can smash into each other and neutralize eachother (if they were the same size wave) called destructive interference.

At this point you may think that it is pretty obvious that light behaves like a wave.But where is the proof that light is really composed of particles called photons? Theproof comes from an experiment that is called the photoelectric effect.

Max Planck: proposed that energy emitted is not done so in a continuous mannerbut is given off in small packets which he called quanta. He determined that anatom can emit only certain amounts of energy and therefore they must containcertain quantities of energy and that those are fixed. Thus, the energy of an atom

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is quantized. The change in the atoms energy results from the gain or loss of oneor more packets of energy. Planck derived an equation to explain this quantizedform of energy (as opposed to the idea that energy emitted was continuous)

Eatom = hwhere h = Plancks constant = 6.626 x 10-34 Js

= frequency (as above)

Despite the fact that Planck thought that energy was quantized, physicistscontinued to think of energy as traveling in waves. Energy as waves, however,could not explain the photoelectric effect.

The electron is emitted from the metal with a specific kinetic energy (i.e. a specificspeed).

The energy associated with a wave is related to its amplitude or intensity. Forexample, at the ocean the bigger the wave, the higher the energy associated withthe wave. It is not the small waves that knock you over it's the big waves! Wavetheory associates the energy of the light with its amplitude, not its color. Soeveryone who thought light is just a wave was really confused when the intensity(amplitude) of the light was increased (brighter light) and the kinetic energy of theemitted electron did not change. What happens is that as you make the lightbrighter more electrons are emitted but all have the same kinetic energy. The wave

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theory thus predicts that an electron would break free when it absorbs enoughenergy from light of any color.

It was already known that light of sufficient energy emitted electrons from a metalsurface.

Well, they thought the kinetic energy of the emitted electron must depend on

something. So they varied the frequency of the light and this changed the kineticenergy of the emitted electron. Varying the frequency of the light changes its color!

This is the idea behind the threshold frequency. Light shining on the metal objectmust be of a sufficient frequency in order to eject and electron from the metal.Different metals have different minimum frequencies. According to the photontheory presented by Einstein, a beam of light consists of an enormous amount ofphotons. Einstein viewed light as being particulate in nature. Light intensity(amplitude) is related to the number of photons striking the surface per unit of time.Therefore, a photon of a certain minimum energy must be absorbed for the electronto be ejected.

The absence of a time lag, a current is detected the minute the light hits the metalplate, regardless of how intense the light is. This violates the wave theory in thatdimmer light would have to shine on the plate longer in order to eject the electron.Basically it was determined that the metal and thus the electrons cannot save upor bank their energy until they store enough for the electron to be emitted. Theelectron will break free the moment a photon of enough energy hits the metal. Thecurrent was weaker in dim light than in brighter light (amplitude again) becausethere are fewer photons per unit of time, but those photons had the correct energyin order to emit the electron.

This result is not consistent with the picture of light as a wave. An explanation thatis consistent with this picture is that light comes in discrete packages, calledphotons, and each photon must have enough energy to eject a single electron.Otherwise, nothing happens. So, the energy of a single photon is:

Eatom= Ephoton = hWhen this was first understood, it was a very startling result. It was Albert Einsteinwho first explained the photoelectric effect and he received the Nobel Prize inPhysics for this work.

So, in summary-light is a particle, but has some wave-like behavior.

Given: Calculate the energies of a photon from the UV region ( = 1 x 10-8 m),

visible ( = 5 x 10-7 m), and infrared ( = 1 x 10-4 m)

E = h

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A spectrum obtained from a glowing source is called an emission spectrum. Whenwhite light is passed through a prism we see a myriad of colors specifically whatwe term to be a rainbow. This dispersion of white light demonstrates that whitelight contains all the wavelengths of color and is thus considered to be continuous.Each color blends into the next with no discontinuity.

When elements are vaporized and then thermally excited, they emit light, however,this light was not in the form of a continuous spectrum as was observed with whitelight. Instead, a discrete line spectrum was seen when the light was passed througha narrow slit. A series of fine lines of different colors separated by large blackspaces was observed. The wavelengths of those lines are characteristic of theelement producing them thus, elements can be identified based on the spectralline data that they produce.

Typically, we can examine the visible line spectra produced by an element in lab using electricity, tubes filled with elements in the gaseous state and a spectroscopeor diffraction grating which separates the light emitted by the gas into itscomponents.

Spectroscopists studied the emission spectrum of hydrogen and identified lines indifferent regions of the EM spectrum. All hydrogen emits these same linesreproducibly. Using a particular equation, the location or wavelength of emissionlines could be predicted.

1= R

2

2

2

1

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nn

where R = 1.096776 x 107 m-1

n2> n1

The observation of the line spectra did not correlate with classical theory for theelectron spinning around the nucleus. It was believed that the electron spinningaround the nucleus should emit radiation and slowly spiral inwards until it collidedwith the nucleus. As the electron spirals inwards, it would do so smoothly and thusshould emit a continuous array of frequencies but that is not so line spectra fromelements are not continuous.

Niels Bohr was working in Rutherfords lab and suggested a model for the H atomthat predicted the existence of line spectra. Bohr used Planck and Einsteins ideasabout quantization of energy and proposed three postulates:

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1.) The H atom has only certain allowable energy levels. These were termedstationary states and can be thought of as a fixed circular orbital that theelectron travels in around the nucleus.

2.) The atom does NOT radiate energy when an electron is in one of its stationarystates. Thus, this violated the ideas of classical physics as Bohr postulatedthat the atom does NOT change its energy while the electron moves in orbit.

3.) The atom can change to another stationary state by the electron moving to

another orbit, only by absorbing or emitting a photon whose energy equalsthe difference between the two stationary states. Thus:

Ephoton = Estate A Estate B = hwhere the energy of state A is greater than the energy of state B

The spectral line results when a photon of a specific energy (and thus specificfrequency and wavelength) is emitted as the electron moves from a higher energystate to a lower energy state. Bohrs model explained that the reason that a linespectrum is not continuous because the atom has only certain discrete levels whichthe electron can travel between.

Think of the discrete levels like steps on a ladder, or lily pads on a pond. A frog (theelectron) can only jump on the lily pads just like a person climbing a ladder canonly climb up the ladder by standing ON the steps. It is very tough to climb a ladderwhen you are not standing on the steps! In fact, I would be that you cant climb aladder that way!

In Bohrs atoms, the principal quantum number, n, is associated with the orbitallocation (the radius of the orbit from the nucleus) . The lower the n value, the closerthe electron is to the nucleus. When the electron for H is in the first energy level itis said to be in the ground state. When energy is imparted to the atom, the electronwill take that energy and jump to a new level, perhaps on n=2 or 3. This is theexcited or high energy state. Maintaining the high energy state requires too muchenergy (think of water at the top of the waterfall how difficult it would be for thatwater to stay at the top). Eventually, the electron falls back down to its groundstate and releases the energy it had absorbed as a photon. Remember that thereare 6.022 x 1023 atoms of H in 1 mole of H which means that the 1 single electroncan have different percentages of electrons in different excited states dropping todifferent levels. The electron can drop from 5 to 2, 4 to 2, 3 to 2 etc . . . Whenelectrons drop from an excited state to the third level (Paschen series), infraredenergy is emitted. When electrons drop from an excited state to the second level,visible energy is emitted (Balmer series). When electrons drop from an excited

state to the first energy level (Lyman series), ultraviolet energy is emitted.

Unfortunately, Bohrs theory only worked for Hydrogen, or Hydrogen like elements(e.g. other 1 electron species, such as ions formed from He, Li, Be, B, C, N, and O).The reason for this is simply, multi-electron systems have

1.)electron electron repulsion2.)electron nucleus attraction

3.) because electrons arent really in fixed orbitals

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Bohrs work did generate an equation which can be used to determine energy levelsthat the electrons are jumping between and also the energy associated with themovement of electrons between energy states:

En = -2.179 x 10-18Jn2

where n = level the electron occupies

It can be further expanded to examine the changes between two energy states suchthat:

Ephoton = - 2.179 x 10-18J

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LHnn

where nH = higher energy levelNL = lower energy level

This energy can then be used to calculate or

Why the negative value for E (equation above). It is due to an arbitrary assignmentof the zero point energy. The zero point energy is defined as when the electron iscompletely removed from the nucleus. Thus, all values for E are negative. It is anarbitrary assignment try not to think about it too hard it does boggle the mind abit! Just remember that our frequencies and wavelengths are not negative. Referback to the diagram of EM radiation none of the numbers are negative!!

Light seems to be able to behave as if it is a wave, and also a particle known asthe wave-particle duality. The wave nature is evident when light is shined through aprism, the particle nature is evident when examining the photoelectric effect. So, ifenergy is particle like, then maybe matter is wave-like said Louis de Broglie.

The waves associated with moving particles are called matter waves. It was provenby J.J. Thomsons son when he detected electron (particles and thus matter) waves

by passing streams of electrons through thin metal foils and onto photographicplates. What he observed were interference patterns similar to those observedwhen light waves passed through a double slit. Thomsons work could only meanthat electrons behave like waves . Similar patterns were soon obtained by beamingneutrons and protons through various crystals. Further work confirmed de Brogliesequation:

=mv

h

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where once again we see wavelength, velocity, and Plancks constantm stands for the mass of the object

One can see that if the mass of the object is very large, the resulting wavelengthwill be very small. In fact, calculating the wavelength for a car moving at 100 mphresults in a wavelength that is far shorter than anything on our EM radiation figure.

It is virtually undetectable. However, for smaller particles, the wavelengths can beobserved. An electron moving at a speed of 100 mph has a wavelength of about 10-5 m almost 100,000 times the size of its atomic radius!

Both matter and energy exhibit wave-like and particle-like properties. This isknown as the wave-particle duality.

If an electron is a moving particle, the we should be able to determine a few thingsabout it namely its speed, and its location in the atom. Heisenberg came along in1927 and said we could not determine both simultaneously. That by determiningthe electrons momentum we would change its location in the orbital, and bydetermining its exact location, we would alter its momentum.

If we measure the position of an electron we must bombard it with photons thisinterferes with the electrons original momentum as well as its location rememberthat when we bombard an electron in the H atom with photons we actually excitethe atom and move the electron from some ground state energy level to an excited

state energy level.

Ultimately this means that we cannot assign fixed paths that the electrons travel in,such as the orbits proposed by Bohr. We thus, can only determine the probability offinding an electron within some region of the space contained in the atom.

Thus, Bohrs atomic model with fixed orbital was abandoned for a model that wasless precise and based on probabilities known as the quantum-mechanical modelof the atom.

Quantum Mechanics examines the wavelike properties of matter on an atomic scale.Erwin Schrdinger came up with his own theory about the structure of the atom andhe called it the quantum-mechanical model of the atom. Based on the work ofHeisenberg, he abandoned the idea of set energy levels described by Bohr andinstead focused on the wave motion of the electron and the probability of theelectron being located in some general space.

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Remember that space is 3-dimensional! We all have volume! Electrons, therefore,move in 3-dimensional space as they travel along their path around the nucleus.Schrdinger came up with an equation to describe the motion of the electrons.Regardless of what the scary equation looks like:

(dont panic, we wont talk about it or do anything with it!!) the answer to thesolved equation results in a given wave function called an atomic orbital. Thisorbital has nothing to do with Bohrs description of electron orbit (like planetsaround the sun).

Heisenberg showed us that we cannot possibly know the location of the electronand its momentum. Schrdinger showed us that we can get some idea or theprobability of the location of the electron e.g. where it is most likely to be found ormost likely to spend most of its time. Using Schrdingers equation, we can identifythe probable location of the electron using an electron density diagram. Thesedensity diagrams are then transposed into pretty pictures in textbooks and aregiven a less scary name and called electron cloud diagrams. Just know that someartist did not make up those pictures, they are based on Schrdingers complexequation (want more on the equation?? Take Physical Chemistry or upper levelphysics classes offered at any of your local or distant universities!!). Schrdingers

equation also verifies that the electron does not reside in the nucleus but outside(confirms previous theories!!) and shows that as the distance away from thenucleus increases, the likelihood of finding the electron there decreases.Unfortunately, we do not know the location to an accuracy of 100%. These electroncloud diagrams are given for the 90% probability of finding an electron in thatlocation! Where it goes the other 10% of the time???!! Maybe nowhere but allthey can say for these probabilities is that the electron spends 90% of its time there. . .

We have already talked about the periodic table have the answers right on it well,here is another example of the periodic table giving you the answer. It is importantto realize that the electron probability diagrams have given way to 4 main types oforbital/electron cloud diagrams. They are known as the s orbital, the p orbital, the dorbital and the f orbital. Remember we are talking about space and space is 3dimensional. So these orbitals much account for space in the x, y, and z directions.

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Each orbital can only hold two electrons we will get into that more when we doelectron configurations right now take my word for it!

The s orbitals: spherical in shape like a basketball

Notice that there is a chance (10%) that the electron will be outsidethis sphere but for the most part, the electron density is centered.

_____________________________________________________________________________________

The p orbitals: elliptical in shape like a dumbbell

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The electron density of the p orbital is shown above. Notice the p orbital is split intoan x component, a y component and a z component. Putting the 3 p orbitals

together results again in a spherical motif of electron density and contribute to thespherical shape of the atom:

___________________________________________________________________________________

The d orbitals: elliptical in shape like a double-dumbbell

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Again: these shapes indicate the electron clouds indicating the probable location for

finding an electron in the d sublevel. Examining the total composite of each of the dorbitals shows a spherical overall shape:

________________________________________________________________________________

The f orbitals: elliptical in shape like a triple-dumbbell

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S orbital: s subshell: spherical in shape with the nucleus in the center. As theprincipal shell level increases, the size of the s orbital increases. Thus a 2s subshellis bigger than a 1s, and a 5s is bigger than a 4, 3, 2, or 1s subshell. All elementshave the s subshell.

P orbital: p subshell: dumbbell shapes with the two regions, or lobes, indicating the

high probability of finding the electron on either side of the nucleus. Neither lobe isfavored. The nucleus lies at a nodal plane (meaning that the probability of finding

the electron at that location is between slim and none and slim is out of town Unlike the s orbital, the p orbital is directional meaning that there is one p orbitalin the x direction, one in the y, and another in the z. Each orbital can hold twoelectrons. One p orbital consists of BOTH lobes. Again, as the principal shellnumber increases, the size of the p orbital increases, such that a 5p is bigger than a4, 3, and 2p orbital. The joining of the p orbitals in a group showing all at thesame time gives one the overall impression of a spherical shape lending credenceto our belief that atoms are spherical in nature. The minimum principal quantumnumber needed to see the p subshell is n=2.

D orbital: d subshell: double dumbbell shapes with four regions, or lobes, indicatingthe high probability of finding the electron on sides of the nucleus in this doublefigure 8 pattern (except for the dz2 orbital). Again there is a node at the nucleus.Again, as the principal shell number increases, the size of the d orbital increases,such that a 6d is bigger than a 5, 4, and 3d orbital. The joining of the d orbitals in agroup showing all at the same time gives one the overall impression of a sphericalshape lending credence to our belief that atoms are spherical in nature. Theminimum principal quantum number needed to see the d subshell is n=3.

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Back to the periodic table giving us the answers! The table IS organized to tell usthe subshells for particular elements. These regions on the periodic table areknown as s, p, d, and f blocks. The block is a pretty obvious term since theyform squares or rectangles on the periodic tables.

The rows on the sides of the periodic table tell us the principal level that we are infor a particular element. This number in turn, becomes our principal quantumnumber. Row 1 has a principal quantum number of n=1. Row 2 has an n=2, row 3has an n=3, row 4 has an n=4, row 5 has an n=5, row 6 has an n=6. You get theidea!! The exception to the rule (because remember, there always seems to be inchemistry!!) is the d block. They do not follow the row number as being theprincipal quantum number, instead, that d block is a principal quantum numberbehind. Thus, when you are in Row 4, the d block principal quantum number = 3.

When you are in row 5, the d block n= 4. When you are in row 6, n =5.

Notice the atomic numbers of the f block elements. They slide into the periodictable in row 6. Their principal quantum numbers are 2 behind their row number.This means that when in row 6, the f block n=4, and when in row 7, the f block n=5.

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1. The principal quantum number = nn is a positive integer and its value is indicated by the row number (with someexceptions shown above!). It indicates the relative size of the atoms andwhat energy level the electron is located in. When n=1, the electron is in thefirst energy level, when n=5, the electron is in the fifth energy level.

2. The angular momentum quantum number = ll tells us the shape of the orbital or subshell where the electron is located.

The s orbital has been assigned the l value =0, p=1, d=2, f=3 and so on

subshell l values 0p 1d 2f 3g 4

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The largest l value is ALWAYS n-1!!! Thus, for a principal quantum number (n)= 0, the only l value possible 0, which corresponds to the s subshell. Whenn=2, we can have l values equal to 1 and 0. This corresponds to the s and psubshells. Notice that the number of possible l values always equals theprincipal quantum number value (e.g. when n=2 we have 2l values). Noticethe correlation between the principal quantum number and the appearance of

orbitals

Given: For an n=3 we can have l values of ?????Which orbitals do these numbers correspond to?Does this make sense off the periodic table and our blocks?

Given: For an n=4, we can have l values of ?????Which orbitals do these numbers correspond to?Does this make sense off the periodic table and our blocks?

3. The magnetic quantum number ml:ml tells us the orientation of the orbital in space. Remember that we can havepx, py, and pz orbitals? Each one of the orbitals is assigned a number toidentify it from its identical twins. The number assignment is arbitrary, buteach orbital gets its own ID number. If there are 3 p orbital types, then there

must be 3 numbers to identify them with. The l value for the p subshell is 1.So we take the + value, the value and every integer in-between to assign theml values. Thus, px = -1, py = 0, and pz = 1. AGAIN, the number assignmentsare arbitrary! Notice that the ml value = -l . . . +1 increasing by integers.

Given: l=2 : what subshell is this? What are the possible ml values?

Ifl=2 then we have values of -2 . . . +2 by integers, that means we have -2,-1, 0, 1, and 2 as possible ml values. By golly how many d orbitals do wehave?? Why there are 5!! Each number above corresponds to one of those

orbitals!!

Quantum numbers can be used to determine what level the electron is in, whatorbital the electron is in, and even which specific orbital the electron is in!

Blackbody radiation:

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A black body is a theoretical object that absorbs 100% of the radiation that hits it.Therefore it reflects no radiation and appears perfectly black.

In practice no material has been found to absorb all incoming radiation, but carbonin its graphite form absorbs all but about 3%. It is also a perfect emitter of radiation.At a particular temperature the black body would emit the maximum amount ofenergy possible for that temperature. This value is known as the black body

radiation. It would emit at every wavelength of light as it must be able to absorbevery wavelength to be sure of absorbing all incoming radiation. The maximumwavelength emitted by a black body radiator is infinite. It also emits a definiteamount of energy at each wavelength for a particular temperature, so standardblack body radiation curves can be drawn for each temperature, showing the energyradiated at each wavelength. All objects emit radiation above absolute zero.

Some Examples:Objects at around room temperature emit mainly infra-red radiation (l 10mm)which is invisible. The sun emits most of its radiation at visible wavelengths,particularly yellow (l 0.5mm). A simple example of a black body radiator is the

furnace. If there is a small hole in the door of the furnace heat energy can enterfrom the outside. Inside the furnace this is absorbed by the inside walls. The wallsare very hot and are also emitting thermal radiation. This may be absorbed byanother part of the furnace wall or it may escape through the whole in the door. This radiation that escapes may contain any wavelength. The furnace is inequilibrium as when it absorbs some radiation it emits some to make up for this andeventually a small amount of this emitted radiation may escape to compensate forthe radiation that entered through the hole. Stars are also approximate black bodyradiators. Most of the light directed at a star is absorbed. It is therefore capable ofabsorbing all wavelengths of electromagnetic radiation, so is also capable ofemitting all wavelengths of electromagnetic radiation. Most approximateblackbodies are solids but stars are an exception because the gas particles in themare so dense they are capable of absorbing the majority of the radiant energy.

the most stable arrangement of the nucleus and the electrons

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