Development of the New Atomic Model PDF

Development of the New Atomic Model Ms Obie – May the Force Be With You!  Light Early in the nineteenth century, Thomas Young: – Light passing through narrow, closely spaced slits produced interference patterns – Could not be explained in te...

Development of the New Atomic Model Ms Obie – May the Force Be With You!  Light Early in the nineteenth century, Thomas Young: – Light passing through narrow, closely spaced slits produced interference patterns – Could not be explained in terms of Newtonian particles but could be easily explained in terms of waves James Clerk Maxwell: – Light was the visible part of a vast spectrum of electromagnetic waves. Particles and waves are connected on a fundamental level called wave-particle duality. Wave: oscillation or periodic movement that can transport energy from one point in space to another. (Shaking a rope or dropping a pebble into pond)  Properties of Light Electromagnetic radiation – Form of light energy that exhibits wave-like behavior –Example: X-rays, UV, Infrared light, microwaves, radiowaves Together above form electromagnetic spectrum Moves at a constant speed – c = 3.0 x 108 m/s (speed of light)  Properties of Light Continued Wavelength () – lambda – Distance between corresponding points on adjacent waves – Frequency () Number of waves that pass a given point in a specific time – usually one second – Expressed in waves/second called a Hertz – Frequency and wavelength are mathematically related to each other: c =   Figure 6.2 One-dimensional sinusoidal waves show the relationship among wavelength, frequency, and speed. The wave with the shortest wavelength has the highest frequency. Amplitude is one-half the height of the wave from peak to trough.  Electromagnetic Spectrum The electromagnetic spectrum is the range of all types of electromagnetic radiation. Visible light makes up only a small portion of the electromagnetic spectrum. Each of the various colors of visible light has specific frequencies and wavelengths associated with them. Portions of the electromagnetic spectrum are shown in order of decreasing frequency and increasing wavelength. (credit “Cosmic ray": modification of work by NASA; credit “PET scan": modification of work by the National Institute of Health; credit “X-ray": modification of work by Dr. Jochen Lengerke; credit “Dental curing": modification of work by the Department of the Navy; credit “Night vision": modification of work by the Department of the Army; credit “Remote": modification of work by Emilian Robert Vicol; credit “Cell phone": modification of work by Brett Jordan; credit “Microwave oven": modification of work by Billy Mabray; credit “AM radio": modification of work by Dave Clausen)  Wave Mechanics Wavelength,  , the distance between peaks, measured in nm.  Wave Mechanics Frequency, , the number of peaks passing a point in a second (unit = Hz). 3 waves/sec = 3 Hz  Wave Mechanics The speed of the radiation,c , can be determined from the frequency and wavelength according to the equation: c  speed = wavelength x frequency = m x 1/sec = m/s Problems Using the Speed Equation Problem: Determine the frequency of light with a wavelength of 4.257 x 10-7cm. Note that the speed of light is 3.0 x 108 m/s. c  we are looking for , so we must rearrange the equation to isolate the variable.  Problems Using the Speed Equation Problem: Determine the frequency of light with a wavelength of 4.257 x 10-7cm. Note that the speed of light is 3.0 x 108 m/s. c    divide both sides by  and simplify Problems Using the Speed Equation Problem: Determine the frequency of light with a wavelength of 4.257 x 10-9cm. Note that the speed of light is 3.0 x 108 m/s. c now substitute so:   and calculate 3.0x10 8 m /s 16  9 7.0x10 Hz 4.257x10 m converted cm to m  Relationship of Energy and Frequency The relationship is as follows: E h E = energy in joules where: h = Planck’s constant  34 6.626x10 J  s = frequency in Hz  Relationship of Energy and Frequency - sample problem Determine the energy of a photon having a frequency of 4.50 x 1014 Hz: plan: E h  34 14 execute: (6.626x10 J  s)x(4.50x10 Hz)  19 calculate: 2.98x10 J  Planck, Max: developed a theoretical expressure for blackbody radiation—an ideal emitter that approximates the behavior of many materials when heated. – His constant is: h = 6 –.626 x 10-34 J·s Einstein’s equation: E = mc2 Planck’s equation: E = hv  Johann Balmer (Balmer Series): developed an empirical equation to determine the bright line spectrumbright line spectrum Rydberg developed an empirical formula that predicted all of hydrogen’s emission lines. Rydberg constant (1.097 x 107 m-1) DeBroglie – determined that wavelength is characteristic of particles, not electromagnetic radiation. Heisenberg Uncertainty Principle and Schrodinger’s equation is the birth of quantum mechanics.  Electron Configuration Ground State vs. Excited State When all electrons in an atom occupy the lowest available orbitals, it is said to be in the ground state. Electron(s) absorb energy, they jump to higher energy levels. This is excited state. (Absorption) As electrons fall back down to the ground state they release energy in the form of the visible light we see. (Emission)  Absorption When an electron “jumps” to a higher energy level it absorbs energy. The excited state is a temporary state. Excited State (i.e. energy level 2) e- Ground State (i.e. Energy level 1)  Emission The electron then falls back down to the ground state, emitting energy. The energy is in the form of light. Every element has its own unique spectrum (think about a prism. Energy Increases as you Move away from the nucleus  Light and Atomic Spectra (bright line spectra) Electromagnetic spectrum consists of light that exists as waves.  Atomic emission spectra produce narrow lines of color called bright line spectra. Each line corresponds to an exact wavelength.  Experiments – Flame Tests Flame Tests – demonstrates the emission spectrum of a substance. Completed by heating elements to high temperatures so they may enter excited state. Characteristic color will be emitted as excited electrons return to ground state. Used to determine metal ion presence in unknown substance.  Experiments – Spectroscopy Spectroscopy – used to view the bright line spectra for given gases. Completed by viewing a gas tube through which an electric current is passed. Use an instrument called a spectroscope, contains a prism to separate emitted light into line spectra.  Quantum Theory Developed to explain chemical behavior of atoms by the scientists discussed earlier. Quantum numbers describe the location of electrons in an atom Quantum – minimum quantity of energy that can be lost or gained by an atom  Electrons Found around the nucleus of the atom in orbitals – Orbital: three-dimensional region around the nucleus Indicates the probable location of an electron at a given point in time. Chemical properties of an element based on the number of outer energy levels of the atom.  Electrons - cont Electrons occupy the lowest sublevel possible Sublevels – orbitals of different shapes Example: # of electrons in sublevel 1s 1 Principle energy level Sublevel The Quantum Model of the Atom The electron location concept consists Study of emission spectra for various of three levels: elements led scientists to the 1. the principal quantum development of an electron level (n) - there are 7 of these configuration (arrangement) model. 2. the sublevel (s,p,d,f) 3. the orbital (s=1, p=3, d=5, f=7)  Rules of Occupancy 1. There are 7 main energy levels 2. number of sublevels in each main energy level = n where n = the main level # 3. number of orbitals per sublevel is:s=1, p=3, d=5, f=7 4. there can only be 2 electrons in each orbital 5. electrons per main energy level is 2n2 Which is the maximum electrons n=3 2(n)2 2(3)2 = 18 electrons Putting it As you move together away from nucleus shells hold more and more electrons AND HAVE “n” describes the size and principle energy GREATER level of the orbital ENERGY There are n2 orbitals per shell As “n” increase the size of the orbital increases and electrons are farther away from the nucleus Principle energy level # of e- 1s 2 Sublevel  Section 4-3: Electron Configurations The electrons accumulate in the orbital locations in the following manner: 1. they occupy the lowest energy level available first 2. each orbital can hold a maximum of 2 electrons 3. they “spread out” into a sublevel before they pair up into orbitals  Electron Configuration The arrangement of electrons in the orbitals of an atom is called the electron configuration of the atom. An electron configuration consists of symbols that contain three pieces of information: 1) The principal quantum shell, n. 2) The letter that designates the orbital type (l). 3) A superscript number that designates the number of electrons in that particular subshell.  Remember the # of electrons equals the atomic number in a neutral atom Aufbau principle – electrons fill one at a time in the lowest possible energy level Pauli Exclusion principle: – (1) no more than two electrons can occupy the same orbital and – (2) two electrons in the same orbital must have opposite spins – Video on how to write electron configuration: – https://www.youtube.com/watch?v=NIwcDnFjj 98  The Filling Order 1s filling happens 2s, 2p from lowest 3s, 3p, 3d energy (1s) to 4s, 4p, 4d, 4f highest (7f) with 5s, 5p, 5d, 5f overlaps after 6s, 6p, 6d, 6f every filled “s” 7s, 7p, 7d, 7f  Configuration of hydrogen 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Hydrogen – 1 electron – 1s1 Member of period # 1 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 1  Configuration of helium 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Helium – 2 electrons – 1s2 Member of period # 1 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2  Configuration of lithium 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Lithium – 3 electrons – 1s22s1 Member of period # 2 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Quantum notation: 2-1  Configuration of beryllium 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Beryllium – 4 electrons – 1s22s2 Member of period # 2 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-2  Configuration of boron 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Boron – 5 electrons – Orbital Notation: 1s 22s22p1 Member of period # 2 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Quantum notation: 2-3  Configuration of carbon 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Carbon – 6 electrons – Orbital Notation: 1s 22s22p2 Member of period # 2 Note: Hund’s rule Arrow notation in effect. __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Quantum notation: 2-4  Configuration of nitrogen 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Nitrogen – 7 electrons – 1s22s22p3 Member of period # 2 Note: Hund’s rule in effect. (The simple rule is that electrons should be placed Arrow notation into separate orbitals before going into the same orbital. __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-5  Configuration of oxygen 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Oxygen – 8 electrons – 1s22s22p4 Member of period # 2 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-6  Configuration of fluorine 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Fluorine – 9 electrons – Orbital Notation: 1s 22s22p5 Member of period # 2 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Quantum notation: 2-7  Configuration of neon 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Neon – 10 electrons – 1s22s22p6 Member of period # 2 This completes the 2nd Arrow notation energy level. It now has __ __ __ __ __ __ __ __a __ stable __ octet! __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8  Configuration of sodium 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Sodium – 11 electrons – Quantum Notation: 1s22s22p63s1 Member of period # 3 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-1  Configuration of magnesium 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Magnesium – 12 electrons – 1s22s22p63s2 Member of period # 3 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-2  Configuration of aluminum 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Aluminum – 13 electrons – 1s22s22p63s23p1 Member of period # 3 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-3  Configuration of silicon 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Silicon – 14 electrons – 1s22s22p63s23p2 Member of period # 3 Arrow notation Hund’s Rule Again __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-4  Configuration of phosphorus 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Phosphorus – 15 electrons – 1s22s22p63s23p3 Member of period # 3 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-5  Configuration of sulfur 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Sulfur – 16 electrons – 1s22s22p63s23p4 Member of period # 3 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-6  Configuration of chlorine 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Chlorine – 17 electrons – 1s22s22p63s23p5 Member of period # 3 Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-7  Configuration of argon 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Argon – 18 electrons – 1s22s22p63s23p6 Member of period # 3 Note: the Arrow notation completion of a __ __ __ __ __ __ __ __ __ __stable __ __octet __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-8  Configuration of iron 1s2, 2s2, 2p6, 3s2 3p6, 4s2, 3d10 Iron – 26 electrons – 1s22s22p63s23p64s23d6 Member of period # 4 Note the energy overlap Arrow notation __ __ __ __ __ __ __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p 4s 3d Well diagram Orbital notation: 2-8-14-2  Quantum mechanics “l” describes the shape of the orbital and has values of: L=0s L=1p L=2d L=3f M- orientation of the electron cloud – angular momentum of orbitals - values from +l to –l Spin “s” +/- ½ Clockwise or counter clockwise arrow up (+ ½) arrow down (- ½) +-  Table 6.1 Quantum Numbers, Their Properties, and Significance Name Symbol Allowed Values Physical Meaning principal quantum n 1, 2, 3, 4, …. shell, the general number region for the value of energy for an electron in the orbital angular momentum ℓ 0 ≤ ℓ ≤ n –1 subshell, the quantum number shape of the orbital magnetic quantum ml –ℓ ≤ ml ≤ ℓ orientation of the number orbital spin quantum ms +½ , –½ direction of the number intrinsic quantum “spinning” of the electron  Values are from 0 to n-1 n (principle energy Sublevel l level) n=1 s 0 n=2 s, p 0, 1 n=3 s, p, d 0, 1, 2 n=4 s, p, d, f 0, 1, 2, 3 Example If you want to determine the principle energy level (n),l and m l you Do the following: n is the principle energy level so for 2p5 the principle energy Level is 2. the sublevel is p and the number of electrons are 5. the p sublevel Three orbitals if you draw the electron configuration because The 5th electron is on the second line(remember Hund’s rule), the spin is -1/2. https://www.youtube.com/watch?v=22vOPpAoxzA Sublevel “l” m S (2e-) 0 0 P (6e-) 1 +1, 0, -1 d (10e-) 2 +2, +1, 0 -1, -2 F (14e-) 3 3,2,1, 0 1,2,3  Electron Configurations and the Periodic Table The periodic table arranges atoms so that elements with the same chemical and physical properties are in the same group. Elements in the same group have similar valence electron configurations. Valence electrons play the most important role in chemical reactions. This describes why elements in the same group have similar chemical reactivity.  Electron Configurations and the Periodic Table Main group elements or representative elements are those in which the last electron added enters an s or a p orbital in the outermost shell. The valence electrons for main group elements are those with the highest n level. Example: Gallium (Ga): [Ar]4s23d104p1 – Ga has three valence electrons (4s2 and 4p1). – The completely filled 3d orbitals count as core, not valence electrons.  Electron Configurations and the Periodic Table Transition elements or transition metals are metallic elements in which the last electron added enters a d orbital. The valence electrons (those added after the last noble gas configuration) in these elements include the ns and (n – 1)d electrons. Example: Vanadium (V): [Ar]4s23d3 – V has five valence electrons (4s2 and 3d3).  Electron Configurations and the Periodic Table Inner transition elements are metallic elements in which the last electron added occupies an f orbital. The valence electrons (those added after the last noble gas configuration) in these elements include the ns, (n – 2)f, and if present the (n – 1)d subshells. Example: Promethium (Pm): [Ar]6s24f5 – Pm has seven valence electrons (6s2 and 4f5).  Electron Configurations of Ions Recall, ions are formed when atoms gain or lose electrons. A cation (positively charged ion) forms when one or more electrons are removed from an atom. – For main group elements, the electrons that were added last are the first electrons removed. – For transition metals and inner transition metals, the highest ns electrons are lost first, and then the (n – 1)d or (n – 2)f electrons are removed.  Electron Configurations of Ions An anion (negatively charged ion) forms when one or more electrons are added to a parent atom. – The electrons are added in the order predicted by the Aufbau principle.  Example 6.11 Predicting Electron Configurations of Ions What is the electron configuration of: (a) Na+ (b) P3– (c) Al2+ (d) Fe2+ (e) Sm3+ First, write out the electron configuration for the parent atom, and then add or remove the appropriate number of electrons from the correct orbital(s).  Example 6.11 (a) Na: 1s22s22p63s1. Sodium loses one electron. Na+: 1s22s22p6 (b) P: 1s22s22p63s23p3. Phosphorus gains three electrons. P3−: 1s22s22p63s23p6 (c) Al: 1s22s22p63s23p1. Aluminum loses two electrons. Al2+: 1s22s22p63s1 (d) Fe: 1s22s22p63s23p64s23d6. Iron loses two electrons and, since it is a transition metal, they are removed from the 4s orbital. Fe2+: 1s22s22p63s23p63d6 (e) Sm: 1s22s22p63s23p64s23d104p65s24d105p66s24f6. Samarium loses three electrons. The first two will be lost from the 6s orbital, and the final one is removed from the 4f orbital. Sm3+: 1s22s22p63s23p64s23d104p65s24d105p64f5  Summary All of the above explains the periodicity on the periodic table. Nuclear charge increases Shielding increases Atomic radius increases Ionic size increases Ionization energy decreases Electronegativity decreases Shielding is constant Nuclear charge increases Atomic Radius decreases Electronegativity increases Ionization energy increases Summary