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Quantum Model.

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Presentation on theme: "Quantum Model."— Presentation transcript:

1 Quantum Model

2 Modern View The atom is mostly empty space Two regions Nucleus
protons and neutrons Electron cloud region where you might find an electron

3 (positive and negative charges) + + Rutherford (1911) (the nucleus) +
Dalton (1803) Thomson (1904) (positive and negative charges) + + Rutherford (1911) (the nucleus) + + + + . Bohr (1913) (energy levels - orbits) . Schrödinger (1926) (electron cloud model – orbitals) From the time of Dalton to Schrödinger, our model of the atom has undergone many modifications.

4 Quantum Mechanical Model
Niels Bohr & Albert Einstein Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals)

5 Bohr Model Planetary model Niels Bohr After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus

6 Bohr’s Work: Quantized Energy
Bohr realized that the atom has various levels of energy (states of energy) The first level is the ground state Although below this is the nucleus The second and following states are “excited” How does it go from ground state to an excited state? BY GAINING ENERGY FROM THE UNIVERSE

7 Bohr Model Nucleus Ground State First Excited State
Second Excited State

8 Absorption of Photon by Atom
When a photon of light is absorbed by an atom, it causes an increase in the energy of the atom. The photon disappears. The energy of the atom increases by exactly the amount of energy contained in the photon. The photon can be absorbed ONLY if it can produce an “allowed” energy increase in the atom.

9 Absorption Spectrum When an atom absorbs photons, it removes the photons from the white light striking the atom, resulting in dark bands in the spectrum Therefore, a spectrum with dark bands in it is called an absorption spectrum

10 Atomic Spectra When an atom absorbs energy from a photon of light, it is absorbed in an electron and thus moves further from the nucleus. Since Bohr explained this principle, the opposite must be also be true and the electron will return to its prior potential energy state when it releases the energy (“loses it”) as light energy.

11 Emission of Photon by Atom
When a photon of light is emitted by an atom, it causes a decrease in the energy of the atom. A photon of light is created. The energy of the atom decreases by exactly the amount of energy contained in the photon that is emitted. The photon can be emitted ONLY if it can produce an “allowed” energy decrease in an excited atom.

12 Emission Spectrum When an atom emits photons, it glows! The photons cause bright lines of light in a spectrum. Therefore, a spectrum with bright bands in it is called an emission spectrum.

13 Bohr Model of Atom Increasing energy of orbits n = 3 n = 2 The Bohr model of the atom, like many ideas in the history of science, was at first prompted by and later partially disproved by experimentation. n = 1 e- A photon is emitted with energy E = hf

14 Bohr’s Model Although his model only worked correctly for the hydrogen atom, it further provided an explanation of some of the chemical properties of the elements The idea that atoms have electron arrangements unique to each element is the foundation of much of our chemical reactions/bonding knowledge Electron configurations!!!

15 Atoms and Their Lights When an atom absorbs energy they are temporarily energized As it loses this energy, it emits light Since each element has its own unique atomic structure, each element has its unique COLOR!

16 Bohr’s Contributions:
Electrons can occupy only certain regions of space, called orbits Orbits closer to the nucleus are more stable — they are at lower energy levels Electrons can move from one orbit to another by absorbing or emitting energy, giving rise to characteristic spectra Bohr’s model only worked for the hydrogen atom, so it needed work His concept of electrons moving in fixed orbits was later abandoned.

17 Bohr Model of the Atom The planetary model was suggested but this is not plausible because if the electrons orbited the nucleus, then it would collide after ~10-9 s Niels Bohr joined Rutherford’s team in and took the planetary model from Rutherford and combined it with Einstein’s Theory of Light and Planck’s quantized energy levels

18 Advances in Atomic Theory
Major advances didn’t begin until scientists applied the properties of waves and light to discovering the atom structure Visible light is a type of electromagnetic radiation

19 Electromagnetism All of these forms of electromagnetism will travel as waves and move at the speed of light 3.0 x 108 m/s Represented by the variable c The significant feature of these, besides their similar speed, is their repetitive nature They will travel in waves continuously repeating

20 Anatomy of a Wave

21 Wavelength Measurements
Wavelength is measured as the distance between corresponding points of adjacent waves e.g. crest to crest, trough to trough, etc. Wavelength is equal to speed divided by frequency l = v/f v is velocity/speed thus it is actually c (speed of light) 𝑐=𝜆𝑓 where c always equals 3.0∙ 𝑚/𝑠

22 Photon Energy If all light moves a wave at the speed of light, then only its wavelength and frequency can vary. Therefore, the differences in these values are what separate the 7 EM waves on the spectrum The frequency (albeit the wavelength factors in too) will determine the amount of energy found in the wave Energy of an EM wave can be calculated by taking the frequency f (in Hertz) and multiplying by a universal constant known as Planck’s Constant h 𝐸=ℎ∙𝑓 or sometimes written as 𝐸=ℎ∙𝑣

23 I. Waves and Particles de Broglie’s Hypothesis
Particles have wave characteristics Waves have particle characteristics λ = h/mv Wave-Particle Duality of Nature

24 Electrons as Waves QUANTIZED WAVELENGTHS Louis de Broglie (1924)
~1924 Louis de Broglie (1924) Applied wave-particle theory to electrons electrons exhibit wave properties QUANTIZED WAVELENGTHS Fundamental mode Second Harmonic or First Overtone Standing Wave 200 150 100 50 - 50 -100 -150 -200 200 150 100 50 - 50 -100 -150 -200 200 150 100 50 - 50 -100 -150 -200

25 Dual Nature of Light Three ways to tell a wave from a particle…
Waves can bend around small obstacles… …and fan out from pinholes. Particles effuse from pinholes Three ways to tell a wave from a particle… wave behavior particle behavior waves interfere particle collide waves diffract particles effuse waves are delocalized particles are localized

26 Werner Heisenberg – The Uncertainty Principle
Think of this… a free electron moves into the focus of a hypothetical microscope and is struck by a photon of light; the photon transfers momentum to the electron The reflected photon is seen in the microscope, but the electron has moved out of focus The electron is not where it appears to be Think of this like seeing stars in space that may not be there anymore because the light is still being seen due to great distances A wave is a disturbance that travels in space and has no fixed position

27 Quantum Mechanics Heisenberg Uncertainty Principle g
Werner Heisenberg ~1926 Heisenberg Uncertainty Principle Impossible to know both velocity and position of an electron at the same time g Microscope Electron

28 New Atomic Model The Heisenberg Uncertainty Principle states in simple terms that it is impossible to describe precisely both the location and the speed of particles that exhibit wavelike behavior The upshot of all this is that we cannot state with absolute certainty the velocity and position of electrons and so we must replace the Bohr model with another which considers the probability of an electron being at a certain point In effect, all we can say is that we are pretty certain that the electron is within a particular region for some (most) of the time Of course outside that time, it could be anywhere

29 II. The Electron as a Wave
Schrödinger’s Wave Equation Used to determine the probability of finding the single electron in a hydrogen atom at any given distance from the nucleus Electron best described as a cloud Effectively covers all points at the same time (fan blades)

30 Quantum Mechanics Schrödinger Wave Equation (1926)
Erwin Schrödinger ~1926 Schrödinger Wave Equation (1926) finite # of solutions  quantized energy levels defines probability of finding an electron

31

32 Quantum Mechanics Orbital (“electron cloud”)
Region in space where there is 90% probability of finding an electron 90% probability of finding the electron Orbital Electron Probability vs. Distance 40 30 Electron Probability (%) 20 10 50 100 150 200 250 Distance from the Nucleus (pm)

33 Quantum Numbers Four Quantum Numbers:
Specify the “address” of each electron in an atom UPPER LEVEL

34 Writing & Reading Quantum Numbers
Since this is the specific “address” for each an every electron found in an atom, This will be done shortly after learning about electron configurations

35 Quantum Model of the Atom
Electron Configuration & Orbitals 1s22s22p63s23p64s23d104p65s24d104p65s24d105p66s24f145d106p6…

36 Ground-State Electron Configurations
The arrangement of electrons in an atom is called the electron configuration Because atoms of different elements have different number of electrons, each atom has its own unique configuration Since low-energy systems are more stable than high-energy systems, electrons in atoms tend to assume arrangements that give the lowest possible energy This is known as Aufbau’s Principle Helium with 2 valence e- in the s orbital

37 Order of Electron Subshell Filling: It does not go “in order” of reading a book
2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 7s2 7p6 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6

38 s p d f 1 2 3 4 5 6 7 1A 2A 3B 4B 5B 6B 7B 8B 1B 2B 3A 4A 5A 6A 7A 8A
Group Number = # valence e- s 1 2 3 4 5 6 7 p Row = # shells d f

39 Electron Configurations
Orbital Energy Diagram Reading the Diagram The energy levels are approximate Meaning that sometimes the higher sub-orbital will fill despite the lower orbital not completely full Beginning in the third energy level (third shell), the sublevels begin to overlap This is because it’s easier to put two electrons together in 4s than 10 in 3d

40 Electron Configurations
There are many different ways to show an atom’s electron configuration: Orbital Diagrams (Arrows) Electron Configuration Notation (1s22s2…) A shortcut has been used to make notations faster too Noble Gas Configurations Lewis Dot There are some exceptions, but we will ignore those rules for now n = # of orbit (or layer of electron cloud) n2 = # of orbitals per orbit (boxes that hold 2 electrons each) 2n2 = # of electrons per orbit (layer)

41 Pauli and Hund Pauli Exclusion Principle Hund’s Rule
Electrons in an atom have a specific spin, just like a top spins on its axis The electron can only spin two ways, these are represented by an up arrow (↑) and a down arrow (↓) Only two electrons can fill a single atomic orbital, but only if they are spinning in opposite directions Hund’s Rule ***Electrons are negative and therefore, must repel one another Hund’s Rule states that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can be paired

42 Electron Configuration
1s1 # valence e- possibilities are: s: 1 or 2 p: 1-6 d: 1-10 f: 1-14 Total e- should equal Atomic # (not for ions) row # shell # possibilities are 1-7 7 rows subshell possibilities are s, p, d, or f 4 subshells What element has an electron configuration of 1s1?

43 Practice Write the standard electron configuration for a neutral phosphorus-32 atom 1s22s22p63s23p3

44 1s22s22p63s23p64s23d104p65s24d7 [Kr]5s24d7 Practice
Write the standard electron configuration for a neutral rhodium-102 atom Well that takes too long, what about the shortcut… Look at the row or period that it belongs to and count down from the start This atom is in the 5th period, so start the electron configuration on the s-orbital of that period with the most recent Noble Gas 1s22s22p63s23p64s23d104p65s24d7 [Kr]5s24d7

45 1s22s22p63s23p64s23d1 [Ar]4s23d1 Practice
Write the standard electron configuration and the Noble Gas electron configuration for Scandium-45 Noble Gas configurations will always be [Noble Gas] followed by row # “s” 1s22s22p63s23p64s23d1 [Ar]4s23d1

46 Practice Write the standard electron configuration for the anion O 2− 1s22s22p6

47 Valence Shells The electrons of the outermost occupied shell in any atom are directly exposed to the external environment and are the first to interact with other atoms. The electrons in the outermost shell, therefore, are quite important They are called valence electrons (valent – combining power of an atom)

48 Valence Electrons Again
Since protons determine the type of atom/matter and electrons determine the reactivity of the atoms… Valence electrons determine the chemical properties of an element (e- in outermost valence shell) As we will cover in the not too distant future, these valence electrons determine the chemical bonds that will form between atoms

49 Valence e- – Electron-Dot Structures
The element’s symbol represents the nucleus and inner valence shell electrons The dots are the valence electrons and must be placed one at a time on the four sides (any sequence) and then paired You only represent the s & p orbital electrons! This method was designed by G.N. Lewis ( ) and thus is also called Lewis Dot Structures

50 Apparent Valence Electrons
What if the atom is from the d or f orbital? As noted earlier, it is sometimes easier to fill shells partially before completely for equal stability You do not need to memorize this an any way Sc Ti V Cr Mn Fe Co Ni Cu Zn Outer Configuration 4s23d1 4s23d2 4s23d3 4s13d5 4s23d5 4s23d6 4s23d7 4s23d8 3d104s1 3d104s2 Apparent Valence Electrons 3 2-4 2-5 2-6 2-7 2 or 3 1 or 2 2

51 Feeling Overwhelmed? Read the Book!
Chemistry "Teacher, may I be excused? My brain is full."


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