Ch. 4 “Electron Configurations Quantum Mechanics Made Simple!

Ch. 4 Electron Configurations Quantum Mechanics Made Simple!

In chapter 3, we began our historical journey though the development of atomic theory. Rutherfords Nuclear Atom was more useful than Daltons or Thomsons models because it was able to explain the results of the alpha particle scattering experiment. As more evidence was accumulated, it, too, was replaced by a better model!

An atom consists of a nucleus nucleus (of protons and neutrons) (of protons and neutrons) electrons in space outside the nucleus. electrons in space outside the nucleus. What we know so far Nucleus Electron cloud

Much of our understanding of how electrons behave in atoms comes from studies of how light interacts with matter. As you know, light travels through space & is a form of radiant energy. This is what makes you feel warm as you stand in sunlight! How light travels through space has been a major source of debate for centuries!

1600s, Isaac Newton suggested that light was made of tiny particles. Newton used a glass prism to refract (bend) sunlight (white light) into a continuous spectrum. Continuous Spectrum a complete array of colors from red to violet. (a rainbow!) ROYGBIV (or VIBGYOR) This process is called Diffraction passing white light through a diffraction grating to produce a continuous spectrum.

1600s, Christian Huygens (Dutch) suggested that light consists of waves (rather than particles) - Wave Model of Light He thought light travels away from its source the way water waves travel away from a stone dropped in a pond. This Wave Model of Light Survived into the 1900s!

In the early 1900s scientists were still using cathode ray tubes to study light ! When they passed electricity through gases, the electrons in the gas atoms would absorb the extra energy. The atom is then said to be excited! However, the electrons dont keep this extra energy for long. They immediately give it back off in the form of Electromagnetic Radiation Energy that travels through space as waves.

12 Light (E-M Radiation) All types travel at light speed (c) 3.00x10 8 m/s All types have wave characteristics (wavelength, frequency) wavelength ( lambda) - distance between successive peaks (m) Frequency ( nu) - # cycles passing a given point each second (1/s or Hz One cycle (Frequency is # of cycles per second)

Electromagnetic Radiation covers a broad spectrum: Radio waves ~ 10 3 m Microwaves ~ 10 -3 m Infrared light ~ 10 -5 m Visible light ~ 10 -6 m Ultraviolet light ~ 10 -8 m X-rays ~ 10 -10 m Gamma rays ~ 10 -12 m red750 nm orange yellow green blue indigo Violet400 nm (decreasing (only long wave on list!) Types of Electromagnetic Radiation

Link to FCC Radio Frequency Chart

Because all EM radiation moves at the same speed, wavelength and frequency are inversely proportional: c What is the wavelength of radiation whose frequency is 6.24 x l0 14 sec -1 ? c 3 x 10 8 m/s c 3 x 10 8 m/s 6.24 x 10 14 s = = = 4.81 x 10 -7 m Speed of light! Is this visible light? If so, what color? 4.81 x 10-7m x 10 9 nm 1 m = 481 nmYES! Blue

2. what is the frequency of radiation whose wavelength is 2.20 x l0 -6 nm? (1 m = 10 9 nm) 2.20 x l0 -6 nm x 1 m 10 9 nm 10 9 nm c = 3 x 10 8 m/s 2.20 x l0 -15 m = 1.36 x 10 23 s -1 Is this visible light? If so, what color? No! Gamma or cosmic radiation = 2.20 x 10 -15 m

Remember - When heat or electricity is passed through a gas, the electrons in the gas atoms absorb the extra energy. The atom is then said to be excited! But, the electrons dont keep this extra energy for long. They immediately give it back off in the form of Electromagnetic Radiation (visible light) One way to demonstrate the emission of light from excited atoms is by using a Flame Test.

Flame Tests strontiumsodiumlithiumpotassiumcopper Many elements give off characteristic light which can be used to help identify them.

Flame Tests You heat a metallic salt & it burns with a colored flame! You heat a metallic salt & it burns with a colored flame! This is the characteristic glow of the excited metal ions! This is the characteristic glow of the excited metal ions!

Fireworks

Flame Emission Spectra methane gas wooden splintstrontium ioncopper ionsodium ion calcium ion

Neon signs Bent up cathode ray tube! NOT!!!

The Electric Pickle Excited atoms can emit light. Excited atoms can emit light. Here the solution in a pickle is excited electrically. The Na + ions in the pickle juice give off light characteristic of that element. Here the solution in a pickle is excited electrically. The Na + ions in the pickle juice give off light characteristic of that element.

Bright-Line Spectra Passing the light from excited atoms through a prism does something different - Passing the light from excited atoms through a prism does something different - The spectrum contains lines of only a few colors or wavelengths.

Bright-Line Emission Spectrum ground state excited state ENERGY IN PHOTON OUT 656 nm486 nm 410 nm 434 nm Wavelength (nm) Prism Slits

Each element has a unique bright-line spectrum. Each element has a unique bright-line spectrum. i.e. an elements fingerprint Helium This is how we know what stars are made of!

Spectrum of White Light

Spectrum of Excited Hydrogen Gas

Emission Spectrum of Hydrogen 1 nm = 1 x 10 -9 m = a billionth of a meter 410 nm434 nm486 nm656 nm 1 nm = 1 x 10-9 m = a billionth of a meter

Continuous and Line Spectra 4000 A o 5000 6000 7000 light Na H Ca Hg 400 450 500 550 600 650 700 750 nm Visible spectrum (nm)

At the beginning of the 20 th century the accepted theory of light was still the wave model. (light & other forms of electromagnetic radiation travel as waves) Scientists found that only a certain minimum energy could excite atoms & get them to emit light. So they knew that energy had to be related to the fundamental properties of frequency & wavelength

The temperature of a Pahoehoe lava flow can be estimated by observing its color. The result agrees well with the measured temperatures of lava flows at about 1,000 to 1,200 C.Pahoehoe

1900, Max Planck (Germany) accurately predicted how the spectrum of radiation emitted by an object changes with its temperature. Max Planck The color (wavelength) of light depends on the temperature Red hot objects are cooler than white hot objects

He named each small chunk of energy a quantum (meaning fixed amount) A quantum is the smallest unit of energy Although small, quanta are significant amounts of energy on the atomic level. Planck suggested that the energy absorbed or emitted by an object is restricted to pieces of particular size.

Planck said that the energy of a Planck said that the energy of a light is directly proportional to light is directly proportional to its frequency its frequency E = h E:energy (J, joules) h:Plancks constant (6.6262 10 -34 Js) :frequency (Hz) :frequency (Hz) really small! E = h c (so inversely proportional to wavelength!) c = so c/

Long Wavelength = Low Frequency = Low ENERGY Short Wavelength = High Frequency = High ENERGY Wavelength Table

Small wavelength Large frequency Large energy Large wavelength Small frequency Small energy

Quantum Theory GIVEN: E = ? = 4.57 10 14 1/s h = 6.6262 10 -34 J s WORK: E = h E = ( 6.6262 10 -34 J s ) ( 4.57 10 14 1/s ) E = 3.03 10 -19 J Example: Find the energy of a red photon with a frequency of 4.57 10 14 1/s. Example: Find the energy of a red photon with a frequency of 4.57 10 14 1/s.

Examples: 1. If a certain light has 7.18 x l0J of energy, what is the frequency of this light? 1. If a certain light has 7.18 x l0 -19 J of energy, what is the frequency of this light? A: 1.08X10 15 s -1 or Hz b. what is the wavelength of this light? A: 2.78X10 -7 m 2. If the frequency of a certain light is 3.8 x l0 14 Hz, what is the energy of this light? A: 2.5X10 -19 J 3. The energy of a certain light is 3.9 x l0 -19 J. What is the wavelength of this light? Is it visible? A: 510 nm Yes visible light.

What if the energy of a car was quantized? The car would only be able to move at certain speeds! Lets say a cars fundamental quantum of energy was equal to 10 mph. If it had 7 quanta, how fast would it be going? Yeppers! 70 mph If it had 3 quanta? 30 mph The car can gain or lose energy only in multiples of its fundamental quantum 10 No gradual acceleration or deceleration! It couldnt go 25 mph or 67 mph

So, why arent we aware of quantum effects in the world around us? Remember the size of Plancks Constant? It is very small (10 -34 ) To us, energy seem continuous because the quanta are too small to be noticed. However, for atoms, which are also very small, quanta are of tremendous significance!

Albert Einstein saw Plancks idea of quantized energy as a new way to think about light. In 1905, Einstein used Plancks equation to explain another puzzling phenomenon The Photoelectric Effect.

The Photoelectric Effect refers to the emission of electrons from a metal when light shines on the metal. The wave theory of light (early 1900) could not explain this phenomenon. For a given metal, no electrons were emitted if the lights frequency was below a certain minimum regardless of how long the light was shone. Light was known to be a form of energy, capable of knocking loose an electron from a metal. But the wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. Scientists couldnt explain why the light had to be of a minimum FREQUENCY in order for the photoelectric effect to occur. The wave theory of light (early 1900) could not explain this phenomenon. For a given metal, no electrons were emitted if the lights frequency was below a certain minimum regardless of how long the light was shone. Light was known to be a form of energy, capable of knocking loose an electron from a metal. But the wave theory of light predicted that light of any frequency could supply enough energy to eject an electron. Scientists couldnt explain why the light had to be of a minimum FREQUENCY in order for the photoelectric effect to occur.

Solar Calculator Solar Panel

Albert Einstein expanded on Plancks theory by explaining that electromagnetic radiation has a dual wave-particle nature. While light exhibits many wavelike properties, it can also be thought of as a stream of particles. Albert Einstein expanded on Plancks theory by explaining that electromagnetic radiation has a dual wave-particle nature. While light exhibits many wavelike properties, it can also be thought of as a stream of particles. Each particle of light carries a quantum of energy directly proportional to the frequency. Einstein called these particles photons. A PHOTON is a packet of light carrying a quantum of energy.

Photon Is light a wave or a particle? Macroscopically it behaves as a wave! On the atomic level, we observe particle properties! Seen on the door to a light-wave lab: "Do not look into laser with remaining good eye."

Dual Nature of Light light exhibits wave properties & particle properties

Einstein explained the photoelectric effect by proposing that electromagnetic radiation is absorbed by matter only in whole numbers of photons. In order for an electron to be ejected from a metal surface, the electron must be struck by a single photon possessing at least the minimum energy (E photon = hv) required to knock the electron loose, this minimum energy corresponds to a minimum frequency. If a photons frequency is below the minimum, then the electron remains bound to the metal surface. Electrons in different metals are bound more or less tightly, so different metals require different minimum frequencies to exhibit the photoelectric effect.

Photoelectric Effect No electrons are emitted Electrons are emitted Metal plate Bright red light infrared rays or Dim blue light ultraviolet rays or

57 Quantized Energy and Photons Phenomena not explained by wave nature of light: 1) Black-body radiation light coming from a heated object (Planck) 2) Photoelectric effect electrons emitted from light illuminated surface (Einstein) 3) Emission Spectra light from electronically excited gas atoms emission spectrum (top), absorption spectrum (bottom)

1913 Niels Bohr studied under Rutherford at Victoria University in Manchester. Bohr refined Rutherford's idea by adding that the electrons were in orbits. Rather like planets orbiting the sun. With each orbit only able to contain a set number of electrons. Neils Bohr

Bohrs Model of Hydrogen Neils Bohr incorporated Plancks quantum theory to explain bright-line spectra. Bohr said the absorptions and emissions of light by hydrogen corresponded to energy changes within the atom. The fact that only certain frequencies are absorbed or emitted by an atom tells us that only certain energy changes are possible in an atom. Niels Bohr (1885-1962) (1913)

Bohrs Bohrs Planetary Model of the Atom electrons exist only in orbits with specific amounts of energy called energy levels Therefore electrons can only gain or lose certain amounts of energy (quanta) The orbit closest to The nucleus is the most stable & lowest In energy.

Bohrs Planetary Model of the Atom Nucleus Electron Orbit Energy Levels Niels Bohr &Albert Einstein

The lowest energy state of an atom is its. The lowest energy state of an atom is its ground state. A state in which an atom has a higher amount of energy is an excited state. When an excited atom returns to its ground state, it gives off photons of energy (light!)

e-e- e-e- Ground state Excited state Electrons can only be at specific energy levels, NOT between levels. Electrons can jump to a higher energy level when the atom absorbs energy. When the electron drops back down to a lower level, it gives the extra energy off as light. Electrons cant stop between energy levels so the jumps involve definite amounts of energy. (amount of energy ~ light color!)

An excited lithium atom emitting a photon of red light as it drops to a lower energy state. Photon of red light emitted Li atom in lower energy state Excited Li atom Energy

Electron Energy Levels nucleus 1 st energy level 2 nd energy level 3 rd energy level Energy absorbed Energy lost

Bohr Model of Atom Increasing energy of orbits n = 1 n = 2 n = 3 A photon is emitted e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e- e-e-

Bohr Model 1 2 3 4 5 6 Energy of photon depends on the difference in energy levels Energy of photon depends on the difference in energy levels Bohrs calculated energies matched the bright-line spectrum for the H atom Bohrs calculated energies matched the bright-line spectrum for the H atom nucleus

70 Bohr Model Limitations Unfortunately, this model only works for Hydrogen! The success of Bohrs model of the hydrogen atom is explaining observed spectral lines led many scientist to conclude that a similar model could be applied to all atoms. It was soon recognized, however, that Bohrs approach did not explain the spectra of atoms with more than one electron. Nor did Bohrs theory explain the chemical behavior of atoms.

With more sophisticated equipment, spectral lines were found to consist of closely spaced lines called Doublets Hydrogen (pretty simple!) Helium (not so basic!) doublets So there had to be more to Bohrs energy levels (orbits) than he realized. MORE SCIENTIFIC ADVANCEMENTS!

De Broglies Hypothesis - De Broglies Hypothesis - Duality of Matter Since waves have particle characteristics (Dual Nature of Light) Since waves have particle characteristics (Dual Nature of Light) Moving particles have wave characteristics Moving particles have wave characteristics Louis de Broglie ~1924 In 1924, Louis DeBroglie suggested that every moving particle has a wave nature just like light!

The Wave-like Electron Louis deBroglie The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves.

Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

According to Isaac Newton, we can determine both the position & momentum of a large body. (like an airplane) However, we CANNOT accurately predict where an electron will be at some future time! Heisenberg Uncertainty Principle (1926) says that it is impossible to know both the location and the momentum of an electron simultaneously. Werner Heisenberg 1901-1976

78 Heisenberg Uncertainty Principle You can find out where the electron is, but not where it is going. OR You can find out where the electron is going, but not where it is! One cannot simultaneously determine both the position and momentum of an electron. Werner Heisenberg

Microscope Electron In order to observe an electron, one would need to hit it with photons having a very short wavelength. Short wavelength photons would have a high frequency and a great deal of energy. If one were to hit an electron, it would cause the motion and the speed of the electron to change. Heisenberg Uncertainty Principle

In the Bohr Model of the atom, the electron is at a fixed distance from the nucleus. He assumed we knew both the position & the momentum of electrons. The Uncertainty Principle disproves this!

83 A New Model! (The Last One!) So De Broglie and Heisenbergs contributions lead us to a new atomic model. It will recognize the wave nature of the electron and describe it in terms appropriate to waves. The resulting model will precisely describe the ENERGY of the electron, while describing its location as a probability.

The Quantum Mechanical Model of the Atom Erwin Schrodinger 1887-1961

Schrdingers Quantum Mechanical Model Used to determine the PROBABILITY of finding an electron at any given distance from the nucleus Describes the electron as a 3-dimensional wave surrounding the nucleus. (fan blades) (fan blades) Schrodinger applied DeBroglies idea of electrons behaving as waves to the problem of electrons in atoms.

Today we say that the Today we say that the electrons are located in a region of space outside the nucleus called the. electron cloud. Quantum Mechanical Model The Quantum Mechanical Model of the atom describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals). The Quantum Mechanical Model of the atom describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals). 1926

Quantum Mechanics Orbital - densist, darkest region of the electron cloud) Orbital - densist, darkest region of the electron cloud) Region in space where there is a high (90%) probability of finding an electron Region in space where there is a high (90%) probability of finding an electron Electron Probability vs. Distance Electron Probability (%) Distance from the Nucleus (pm) 100150200250500 0 10 20 30 40 90% probability of finding the electron Electron cloud

LOHS AP Chemistry Fall 2007Dr. Schrempp89 Erwin Schrodinger (1887-1961) Won Nobel Prize in 1933 for his equation. Came up with a paradoxical thought experiment to show problems in observing isolated systems (Schrodingers Cat) Experiment: A cat is placed in a sealed box containing a device that has a 50% chance of killing the cat.

94 Unfortunately, Schrodingers cat could not cope with a life of uncertainty

Modern View The atom is mostly empty space The atom is mostly empty space Two regions Two regions Nucleus Nucleus protons and neutrons protons and neutrons Electron cloud Electron cloud region where you are likely to find an electron region where you are likely to find an electron I don't like it, and I'm sorry I ever had anything to do with it. - Erwin Schrodinger talking about Quantum Physics

Feeling overwhelmed? Just a little more! "Teacher, may I be excused? My brain is full." Chemistry

Quantum Numbers Four Quantum Numbers: Specify the address of each electron in an atom 4 4 4 4 4 4 4 4 4 4 The QM model makes it possible to describe the location of an electron Using four quantum numbers.

Quantum Numbers Principal Quantum Number ( n ) Angular Momentum Quantum # ( l ) Magnetic Quantum Number ( m ) Spin Quantum Number ( s )

Quantum Numbers 1. Principal Quantum Number ( n ) Energy level Energy level Size of the orbital cloud Size of the orbital cloud n = 1, 2, 3, 4 n = 1, 2, 3, 4 n 2 = # of orbitals in the energy level n 2 = # of orbitals in the energy level 2n 2 = # of electrons per energy level 2n 2 = # of electrons per energy level 1s1s 2s2s 3s3s

Electron Energy Level (Shell) Principle Quantum number Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. n is also known as the Principle Quantum number

Quantum Numbers s p d f 2. Angular Momentum Quantum # ( l ) Corresponds to: Energy sublevel Shape of the orbital l = s, p, d, f (in order of increasing energy) s cloud is spherical, p cloud is dumb-bell shaped

Sublevel names

Quantum Numbers 3. Magnetic Quantum Number ( m ) Orientation (direction in space) of orbital Specifies the exact orbital within each sublevel that the electron occupies.

Quantum Numbers pxpxpxpx pzpzpzpz pypypypy x y z x y z x y z A p sublevel has 3 possible orbitals oriented along the x, y, & z axes.

d-orbitals A d sublevel has 5 possible orbital clouds

f Orbitals An f sublevel has 7 possible orbitals

Quantum Numbers 4. 4. Spin Quantum Number ( s ) Electron spin + or - Electron spin + or - An orbital can hold 2 electrons that spin in opposite directions clockwise or counter- clockwise. An orbital can hold 2 electrons that spin in opposite directions clockwise or counter- clockwise. +-

3s3s 3p3p 3d3d 2s2s 2p2p A Cross Section of an Atom 1s1s n0p+n0p+ The first energy level has only one sublevel (1s). The second energy level has two sublevels (2s and 2p). The third energy level has three sublevels (3s, 3p, & 3d). *Although the diagram suggests that electrons travel in circular orbits, this is a simplification and is not actually the case. Rings of Saturn # of sublevels per energy level = n

Quantum Numbers n = 3 n = 2n = 1 Principallevel Sublevel Orbital ssp sp d pxpx pypy pzpz d xy d xz d yz dz2dz2 d x 2 - y 2 pxpx pypy pzpz

Maximum Electron Capacities of Subshells and Principal Shells n 1 2 3 4...n l 0 0 1 0 1 2 0 1 2 3 Sublevel designation designation s s p s p d s p d f Orbitals in sublevel sublevel 1 1 3 1 3 5 1 3 5 7 Sublevel capacity capacity 2 2 6 2 6 10 2 6 10 14 Principal energy Level capacity Level capacity 2 8 18 32...2n 2

Quantum Numbers 1. Principal # n 2. Ang. Mom. # l 3. Magnetic # m 4. Spin # s energy level sublevel (s,p,d,f) orbitalelectron Each electron has a unique address: Each electron has a unique address: 2p x +1/2 Energy level sublevel orbital Electron spin

ATOMIC STRUCTURE There are 3 ways to represent the electron arrangement of an atom: 1.Electronic Configuration 3.Lewis Dot Diagrams 2. Orbital Filling Diagrams

3 ways to represent electron arrangements in atoms: Orbital Notation (orbital filling diagrams): An orbital is represented by a line or box. An orbital is represented by a line or box. The lines are labeled with the principal quantum number and the sublevel letter. The lines are labeled with the principal quantum number and the sublevel letter. Arrows represent the electrons. Arrows represent the electrons. An orbital containing one electron is written as , an orbital with two electrons is written as . An orbital containing one electron is written as , an orbital with two electrons is written as .

Electron Configurations 2p42p42p42p4 Energy Level Sublevel Number of electrons in the sublevel A list of all the electrons in an atom (or ion) 1s 2 2s 2 2p 4

Filling Rules for Electron Orbitals Pauli Exclusion Principle No two electrons in an atom can have the same 4 quantum numbers. No two electrons in an atom can have the same 4 quantum numbers. Each orbital can only hold TWO electrons with opposite spins. Each orbital can only hold TWO electrons with opposite spins. Wolfgang Pauli

General Rules Aufbau Principle Electrons fill the lowest energy orbitals first. Electrons fill the lowest energy orbitals first. Lazy Tenant Rule Lazy Tenant Rule 2s2s 3s3s 4s4s 5s5s 6s6s 7s7s 1s1s 2p2p 3p3p 4p4p 5p5p 6p6p 3d3d 4d4d 5d5d 6d6d 4f4f 5f5f 1s1s 2s2s 2p2p 3s3s 3p3p 4s4s 4p4p 3d3d 4d4d 5s5s 5p5p 6s6s 7s7s 6p6p 6d6d 4f4f 5f5f 5d5d Energy * Aufbau is German for building up

Examples: (Remember that you must place one electron into each orbital of a sublevel before a second electron in placed into an orbital.) Hydrogen 1s 1sHelium Lithium 1s 2s 1s 2s Carbon 1s 2s 2p x 2p y 2p z 1s 1 1s 2 1s 2 2s 1 Be 1s 2s 1s 2 2s 2 Boron 1s 2s 2p x 2p y 2p z 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2

1s 2s 2p x 2p y 2p z Carbon 1s 2 2s 2 2p 2 Hunds Rule Within a sublevel, place one electron per orbital before pairing them. Empty Bus Seat Rule

Hunds Rule Also called The Monopoly Rule Also called The Monopoly Rule In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties has at least 1 house. In Monopoly, you have to build houses EVENLY. You can not put 2 houses on a property until all the properties has at least 1 house.

O 8e - Orbital Diagram Orbital Diagram Electron Configuration Electron Configuration 1s 2 2s 2 2p 4 Orbital Notation 1s 2s 2p O 15.9994 8

Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration Orbital Filling Element 1s 2s 2p x 2p y 2p z 3s Configuration Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2 Orbital Notations Electron H He Li C N O F Ne Na 1s 1 1s 2 2s 2 2p 6 3s 1 1s 2 2s 2 2p 6 1s 2 2s 2 2p 5 1s 2 2s 2 2p 4 1s 2 2s 2 2p 3 1s 2 2s 2 2p 2 1s 2 2s 1 1s 2

4f4f 4d4d 4p4p 4s4s n = 4 3d3d 3p3p 3s3s n = 3 2p2p 2s2s n = 2 1s1s n = 1 Energy Filling order becomes irregular after Ar Because of overlapping of electron clouds (orbitals) in larger atoms Write the electron config. for K Sc ?

Diagonal Rule Must be able to write it for the test! Without it, you may not get correct answers ! Must be able to write it for the test! Without it, you may not get correct answers ! The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy The diagonal rule is a memory device that helps you remember the order of the filling of the orbitals from lowest energy to highest energy Order in which orbitals are filled with electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2

1s22s23s24s25s26s27s21s22s23s24s25s26s27s2 2p63p64p65p66p67p6 2p63p64p65p66p67p6 3d 10 4d 10 5d 10 6d 10 7d 10 4f 14 5f 14 6f 14 7f 14 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2

Electron Configurations Ws. #3 Symbol# e - Orbital Diagram and Longhand Electron Configuration Mg 12 1s 2s 2p 3s P 15 1s 2s 2p 3s 3p V 23 1s 2s 2p 3s 3p 3d 4s Ge 32 1s 2s 2p 3s 3p 3d 4s ___ 4p Kr 36 1s 2s 2p 3s 3p 3d 4s 4p O 8 1s 2s 2p

Part B Rules of Electron Configurations Which of the following rules is being violated in each electron configuration below? Explain your answer for each. Hunds Rule, Pauli Exclusion Principle, Aufbau Principle __ __ Hund s Rule should be 1 electron in each of 1s 2s 2p the 1st 2 orbitals (not doubled up) ___ _ _ Aufbau Principle need 1s 2s 2p 3s 3p to fill 3s before 3p _ Pauli Exclusion Principle 1s 2s 2p 3s 3p 1 electron in 3s needs to point down (opposite spins) 1s 2s 2p 3s 3p 3d Aufbau must fill 4s before 3d

Write out the complete electron configuration for the following: 1) An atom of nitrogen 2) An atom of silver Fill in the orbital boxes for an atom of nickel (Ni) 2s2s2p2p 3s3s 3p3p 4s4s 3d3d1s1s 1s 2 2s 2 2p 3 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 9

Shorthand Notation A way of abbreviating long electron configurations A way of abbreviating long electron configurations Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration Since we are only concerned about the outermost electrons, we can skip to places we know are completely full (noble gases), and then finish the configuration

Shorthand Notation Step 1: Its the Showcase Showdown! Step 1: Its the Showcase Showdown! Find the closest noble gas to the atom (or ion), WITHOUT GOING OVER the number of electrons in the atom (or ion). Write the noble gas in brackets [ ]. Step 2: Find where to resume by finding the next energy level. Step 2: Find where to resume by finding the next energy level. Step 3: Resume the configuration until its finished. Step 3: Resume the configuration until its finished.

neon's electron configuration (1s 2 2s 2 2p 6 ) Shorthand Configuration for Na [Ne] 3s 1 third energy level one electron in the s orbital orbital shape Na = [1s 2 2s 2 2p 6 ] 3s 1 electron configuration A B C D

Part B Shorthand Electron Configuration Use the Noble gas that comes before an element on the periodic table to represent all inner electrons. Put the symbol of the Noble gas in parentheses Sy mb ol # e - Shorthand Electron Configuration Ca 20[Ar] 4s 2 Pb 82[Xe] 6s 2 4f 14 5d 10 6p 2 F 9[He] 2s 2 2p 5 U 92[Rn] 7s 2 5f 4

Shorthand Notation Chlorine Chlorine Longhand is 1s 2 2s 2 2p 6 3s 2 3p 5 Longhand is 1s 2 2s 2 2p 6 3s 2 3p 5 You can abbreviate the first 10 electrons with a noble gas, Neon. [Ne] replaces 1s 2 2s 2 2p 6 The next energy level after Neon is 3 is 3 So you start at level 3 on the diagonal rule (all levels start with s) and finish the configuration by adding 7 more electrons to bring the total to 17 [Ne] 3s 2 3p 5

Boron is 1s 2 2s 2 p 1 Boron is 1s 2 2s 2 2p 1 The noble gas preceding Boron is He, so the short way is [He]2s 2 p. The noble gas preceding Boron is He, so the short way is [He]2s 2 2p 1. Sulfur is 1s 2 2s 2 Sulfur is 1s 2 2s 2 2p 6 3s 2 3p 4 Short way: [Ne]3s3p Short way: [Ne]3s 2 3p 4 Example: Titanium [Ar]4s 2 3d 2 [Ar]4s 2 3d 2

Practice Shorthand Notation Write the shorthand notation for each of the following atoms: Write the shorthand notation for each of the following atoms:KCaIBi

Shorthand Configuration [Ar] 4s 2 Electron configurationElement symbol [Ar] 4s 2 3d 3 [Rn] 7s 2 5f 14 6d 4 [He] 2s 2 2p 5 [Kr] 5s 2 4d 9 [Kr] 5s 2 4d 10 5p 5 [Kr] 5s 2 4d 10 5p 6 Ca V Sg F Ag I Xe Fe [Ar] 4s 2 3d 6

Valence Electrons Electrons are divided between core and Electrons are divided between core and valence electrons electrons in the highest energy level of an atom. These are the electrons that take part in reactions (so most important!) Never more than 8 valence electrons in an atom (2 s & 6 p)

Lewis (Electron) Dot Structures Shorter than configs. & only show valence electrons Element symbol surrounded by # dots = to number of valence electrons. No more than 2 electrons per side on symbol. Start at top of symbol, add dots clockwise, one per side before doubling them up.

B 1s 2 2s 2 2p 1, Core = [He], valence = 2s 2 2p 1 Br [Ar] 4s 2 3d 10 4p 5 Core = [Ar] 3d 10 valence = 4s 2 4p 5 Br is in 7A, so 7 valence electrons No. of valence electrons = Group number (for A groups) B B is in Group 3A, so 3 valence electrons (dots) Br

It is very important to define stable here. STABLE means: STABLE means: (in order of stability) 1. all equal energy orbitals are FULL 2. 2. all orbitals are half-full 3. 3. all orbitals are totally empty.

1. The highest energy electron is the LAST one you write in the electron configuration. 1s 2 2s 2 2p 6 3s 2 3p 5 -- the 3p 5 electron is the last written. *Remember Aufbaus Principle, electrons fill from the lowest to the highest energy. 2. 2. The outermost electron is the one with the LARGEST principle quantum number. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2. The 4 p 2 is the farthest from the nucleus. OR (2) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10. Here, it is the 4s 2 electron, because it has the largest principle q.n. Some More Stuff!!

Irregular Electron configurations sometimes the electron configuration is NOT what we would predict it to be. Sometimes electrons are moved because (1) it will result in greater stability for that atom or (2) for some unknown reason??

Examples Predict the electron configuration for Cr #24: [Ar]4s 2 3d 4 However, the real E. C. is [Ar]4s 1 3d 5. The 4s 2 electron has been moved to 3d to achieve greater stability.

153 Exceptions, Cont. Cr: [Ar] 4s 1 3d 5 NOT Cr: [Ar] 4s 2 3d 4 Because lower energy results from half- filling 6 orbitals with spins aligned instead of causing repulsion in one of the 3d orbitals Cu: [Ar] 4s 1 3d 10 NOT [Ar] 4s 2 3d 9 Because easier to add the electron to a sublevel with four electrons that already have the same spin than causing repulsion in a different orbital

Electron Orbitals: Electron orbitals Equivalent electron shells (Bohr) Neon Ne-10: 1s, 2s and 2p

Relative Sizes 1s and 2s 1s 2s

s, p, and d-orbitals A s orbitals: Hold 2 electrons (outer orbitals of Groups 1 and 2) B p orbitals: Each of 3 pairs of lobes holds 2 electrons = 6 electrons (outer orbitals of Groups 13 to 18) C d orbitals: Each of 5 sets of lobes holds 2 electrons = 10 electrons (found in elements with atomic no. of 21 and higher)

Principal Energy Levels 1 and 2

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N HH He Li C N Al Ar F Fe LaHeLiCNAlArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N H = 1s 1 Hydrogen H He Li C N Al Ar F Fe LaHeLiCNAlArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N He = 1s 2 Helium HH He Li C N Al Ar F Fe LaLiCNAlArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N Li = 1s 2 2s 1 Lithium HH He Li C N Al Ar F Fe LaHeCNAlArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N C = 1s 2 2s 2 2p 2 Carbon HH He Li C N Al Ar F Fe LaHeLiNAlArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N N = 1s 2 2s 2 2p 3 Bohr Model Nitrogen Hunds Rule maximum number of unpaired orbitals. HH He Li C N Al Ar F Fe LaHeLiCAlArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N F = 1s 2 2s 2 2p 5 Fluorine HH He Li C N Al Ar F Fe LaHeLiCNAlArFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Bohr Model Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N Al = 1s 2 2s 2 2p 6 3s 2 3p 1 Aluminum HH He Li C N Al Ar F Fe LaHeLiCNArFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS Electron Configuration CLICK ON ELEMENT TO FILL IN CHARTS N Ar = 1s 2 2s 2 2p 6 3s 2 3p 6 Bohr Model Argon HH He Li C N Al Ar F Fe LaHeLiCNAlFFeLa

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS CLICK ON ELEMENT TO FILL IN CHARTS Fe = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 N HH He Li C N Al Ar F Fe LaHeLiCNAlArFLa Bohr Model Iron Electron Configuration

Aufbau Diagram Arbitrary Energy Scale 1s 2s 2p 3s 3p 4s 4p 3d 5s 5p 4d 6s 6p 5d 4f NUCLEUS CLICK ON ELEMENT TO FILL IN CHARTS La = 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 5d 1 N HH He Li C N Al Ar F Fe LaHeLiCNAlArFFe Bohr Model Lanthanum Electron Configuration

So, where are the electrons of an atom located? Various Models of the Atom Daltons Model Thompsons Plum Pudding Model Rutherfords Model Bohrs Planetary Model electrons rotate around the nucleus Quantum Mechanics Model modern description of the electron in atoms, derived from a mathematical equation (Schrodingers wave equation)

Development of Atomic Models Rutherford model In the early twentieth century, Rutherford showed that most of an atom's mass is concentrated in a small, positively charged region called the nucleus. Bohr model After Rutherford's discovery, Bohr proposed that electrons travel in definite orbits around the nucleus. Thomson model In the nineteenth century, Thomson described the atom as a ball of positive charge containing a number of electrons. Quantum mechanical model Modern atomic theory described the electronic structure of the atom as the probability of finding electrons within certain regions of space.

Diagonal Rule s s 3p 3d s 2p s 4p 4d 4f s 5p 5d 5f s 6p 6d 6f s 7p 7d 7f 1234567 Steps: 1.Write the energy levels top to bottom. 2.Write the orbitals in s, p, d, f order. Write the same number of orbitals as the energy level. 3.Draw diagonal lines from the top right to the bottom left. follow the arrows! 4.To get the correct order, follow the arrows! By this point, we are past the current periodic table so we can stop.

Atoms like to either empty or fill their outermost level. Since the outer level contains two s electrons and six p electrons (d & f are always in lower levels), the optimum number of electrons is eight. This is called the octet rule.

Documents

Ch. 4 “Electron Configurations Quantum Mechanics Made Simple!