Virtual Museum

1. Virtual Museum Quantum Mechanics and The Periodic Table

2. Main Floor Electromagnetic Radiation 2nd Floor

3. Waves and Electromagnetic Radiation Low Frequency (n) Long Wavelength (l) Low Energy • Electromagnetic Radiation: form of energy that exhibits wavelike behavior as it travels through space. • Wavelength: the shortest distance b/w 2 corresponding points on a wave (l) • Amplitude: the height of a wave from its origin to its crest. • Frequency: the number of waves that pass a given point per second (n) • Frequency and wavelength are inversely related therefore the longer the wavelength the lower the frequency. • Frequency and energy are directly related therefore the higher the frequency the higher the energy. High Energy High Frequency (n) Short Wavelength (l) Main Floor Photo taken from forddog.glogster.com

4. A Practical Look At The Electromagnetic Spectrum Electromagnetic Spectrum: a spectrum representing all forms of electromagnetic radiation based on wavelength and frequency. Main Floor Image taken from lbl.gov

5. Electromagnetic Spectrum Radio waves have long wavelengths, low frequencies, and low energy. Gamma rays have short wavelengths, high frequencies, and high energy Notice that the visible spectrum (ROYGBIV) is only a small portion of the electromagnetic spectrum. Image taken from kollewin.com Main Floor

6. Atomic Line Spectrum • Planck discovered that matter can only lose or gain energy in small specific amounts that he called quanta. • Einstein proposed that light could have both wavelike behavior and particle-like behavior after performing an experiment that involved the photoelectric effect. • When light acts as a particle, it is called a packet of light named a photon. Photons have no mass and carry a quantum of energy. • The atomic line spectrum shows that elements emit their energy as photons (particles of light) hence the line instead of the continuous spectrum. Image taken from faculty.virginia.edu Main Floor

7. 2nd Floor—Bohr Models 3rd Floor Main Floor

8. Bohr Model of the Atom • Why are elements’ emission spectrum line spectrums instead of continuous spectrums? • Niels Bohr proposed that atoms only have certain allowable energy states. • The lowest allowable energy state of an atom is called the ground state (closest to the nucleus) • Electrons that move to higher energy levels are said to be in an excited state. • When electrons move from an excited state to a lower energy state they emit energy in the form of photons, hence the line spectrum instead of a continuous spectrum. Image taken from reich-chemistry.wikispaces.com 2nd Floor Main Floor

9. Niels Bohr • Danish Physicist who worked in Rutherford’s Laboratory. • Bohr studied the Hydrogen atom to try to create an explanation for hydrogen’s line emission instead of a continuous spectrum. • Bohr’s model correctly explains and predicts information for the hydrogen atom but fails when multi-electronic atoms are explored. Image taken from counterbalance.org Main Floor 2nd Floor

10. Bohr Models of Atoms You may have drawn Bohr Models of atoms in Middle School. Do you remember the 2, 8, 18, 32 rule for numbers of electrons in energy levels? We will explain where that rule comes from and how to properly illustrate the placement of electrons in the atom. Image taken from library.thinkquest.org 2nd Floor Main Floor

11. 3rd Floor—Influential Scientists 4th Floor Main Floor

12. De Broglie Louis de Broglie was the 1st to propose that particles (ie electrons) could have wavelike behavior. He felt that if light could act like a particle then particles could act like light. This phenomenon is called the wave particle duality of nature. He developed an equation to describe the wave characteristic of a particle l=h/mv l-wavelength h-Planck’s constant m-mass v-velocity Image taken from nobelprize.org Main Floor 3rd Floor

13. Heisenberg • Heisenberg was a German Physicist who determined that it is impossible to make a measurement on an object without disturbing the object. • Heisenberg’s Uncertainty Principle states that it is impossible to know both the position and the velocity of a particle at the same time. • Consider a helium balloon in a dark room. You could reach out and determine the position of the balloon with your hand but the touch of your hand would alter the position of the balloon. • Conversely, if a billiard ball is moving you can determine the velocity, but not the position. In order to determine the actual position, you would have to alter the velocity of the billiard. Image taken from davebruns.com Main Floor 3rd Floor

14. Schrodinger • Austrian Physicist Schrodinger was the 1st to create an equation that treated the hydrogen atom’s electron as a wave. Up until this point all equations treated the electron as a particle. • EY=HY • E=energy Y=wave fcn • H=Hamiltonian Factor • Schrodinger’s Equation led to the idea of areas of high probability of finding electrons (proposed by Max Born) that were called atomic orbitals. Image taken from reich-chemistry.wikispaces.com 3rd Floor Main Floor

15. 4th Floor Orbitals and Quantum Numbers Examples Quantum Numbers Main Floor 5th Floor

16. Atomic Orbitals • S orbitals—all energy levels have an “s” orbital • The “s” orbital is spherical in shape as shown red in the image to the left. • P orbitals—energy levels 2 and up have “p” orbitals • The “p” orbital is dumbbell in shape as shown in yellow in the image to the left. • D orbitals—energy levels 3 and up have “d” orbitals • The “d” orbital has a rosette shape as shown in blue in the image to the left. • F orbitals—energy levels 4 and up contain “f” orbitals • The “f” orbital has a very complex shape as shown in green in the image to the left. Image taken from chemwiki.ucdavis.edu Main Floor 4th Floor

17. Principle Quantum Number (n ranges from 1infinity) describes the energy level in which the electron resides Azimuthal Quantum Number (l ranges from 0n-1 ): describes the sublevel that the electron resides in (s,p,d,or f) Angular Momentum Quantum Number (m ranges from –l+l): describes the orientation of the orbital in space (px,py,pz) Spin Quantum Number (s can be +1/2 or -1/2): describes the spin of the electron. +1/2 means a clockwise spin and -1/2 means a counterclockwise spin Quantum Numbers Main Floor 4th Floor

18. Given n = 2 determine the l, m, and s values for all electrons found in n=2. n=2 l=0n-1 so l=0,1 m=-l+1 so m=-1,0,1 s=+1/2 or -1/2 This information tells that the electrons are in the 2nd energy level (n=2), there are s and p sublevels (l=0 is s, l=1 is p, l=2 is d, and l=3 is f), there is a pxpyand a pz orbital (m shows 3 numbers so 3 orbitals of maximum energy) and the electrons have both clockwise and counterclockwise spins Give a set of the 4 quantum numbers that describes an electron in the 3rd energy level, the d sublevel with a counterclockwise spin. n =3 because in 3rd energy level l =2 because the electron is in a d sublevel m could equal -2,-1, 0, 1, 2 s =-1/2 because the spin is counterclockwise Possible answers: (3, 2, -2, -1/2) or (3, 2, -1, -1/2) (3, 2, 0, -1/2) or (3, 2, 1, -1/2) or (3, 2, 2, -1/2) Notice that there are 5 answers. That is because there are 5 “d” orbitals each having its own set of quantum numbers. I asked for one of the electrons so all answers are correct. Examples Main Floor 4th Floor

19. Orbital Notations 5th Floor Configurations Noble Gas Configurations Electron Configurations Valencee- Main Floor 6th Floor

20. Orbital Notations: notations that show e- in their individual orbitals • The electrons fill the lowest orbital 1st—nature is lazy • Example (Al has 13 e-): Image taken from askmehelpdesk.com Image taken from mrbigler.com Aufbau’s Principle: each e- occupies the lowest energy orbital possible Notice that the 1s orbital is filled first, then the 2s orbital, then the 3-2p orbitals, then the 3s orbital, and lastly one of the 3p orbitals needs to be used to house all 13 e-. The order follows the chart to the left. Pauli’s Exclusion Principle: no 2 e- can have the same 4 quantum numbers (they must have opposite spins) Hund’s Rule: e- with the same spin occupy occupy equal energy levels separately and then are paired. Main Floor 5th Floor

21. Electron Configurations: Configurations that show the energy level, the orbital shape, and the number of electrons in the orbitals. • Example C—has 6 e- Example Ne—10 e- 1s22s22p2 1s22s22p6 Number of e- Notice that the exponents, the number of e-, add up to be 6 e-. Notice that the exponents, the number of e-, add up to be 10 e-. Energy level Sublevel (orbital shape) Notice that the e- configurations obey Aufbau’s Principle. MainFloor 5th Floor

22. Noble Gas Configuration: shortcut configurations that use the noble gases as a reference point • Example: Cl • Find the noble gas prior to the element and start there • Finish the e- configuration from the prior noble gas [Ne]3s23p5 Image taken from hibbing.edu The periodic table is arranged similar e- configurations and therefore can be used to determine both the e- configuration and the noble gas configuration of an element. Main Floor 5th Floor

23. Valence e-: Electrons found in the outermost energy level—responsible for chemical bonding • Example • Na 1s22s22p63s1 • P 1s22s22p63s23p3 Outermost energy level is 3 so count electrons in 3rd energy level only 3s1—exponent is 1 so 1 valence electron 3rd energy level outermost—count all e- in 3rd energy level 3s23p3—add exponents to get valence e-: 3 + 2 = 5 valence e- Main Floor 5th Floor

24. 6th Floor—The Periodic Table Periodic Law Families And Series PT and s, p, d, and f orbitals 7th Floor Main Floor

25. The Modern Periodic Law: elements have a periodic repetition of chemical and physical properties when arranged in order of increasing atomic number. • Periodic Table has rows and columns • Rows are called periods or series • Columns are called groups or families • Elements in the same column share similar chemical properties due to having similar e- configurations • Elements in the same column have the same number of valence electrons 6th Floor Main Floor

26. PT and s, p, d, and f orbitals The 1st and 2nd column elements end in an “s” configuration. The 13th-18th column elements represents the “p” block elements The 3rd-12th column elements represent the “d” block elements The bottom 2 rows of the periodic table represent the “f” elements Image taken from chemwiki.ucdavis.edu We will practice using the periodic table to write configurations. Main Floor 6th Floor

27. The elements that border both the metals and nonmetals are called semi-metals or metalliods. They exhibit characteristics of both metals and nonmetals. Families and Series Most of the atoms on the PT are metals. Metals are defined as atoms that readily lose electrons The right side elements are nonmetals. They readily gain electrons Notice that certain columns have special names as do the 2 rows at the bottom of the PT. Main Floor 6th Floor

28. 7th Floor—Periodic Trends Electronegativity Ionic Radii Ionization Energy Electron Affinity Atomic Radii Main Floor

29. Atomic Radii: basically tells the size of the atom • The atomic radii increases going down a group and decreases across a period • The atomic radii increases down a group due to the addition of energy levels in the atoms • The atomic radii decreases across a period due to the increased nuclear charge pulling the electrons towards the nucleus • Example: Which of the following has the largest atomic radii? Na or Cl—Na because it is farther left on the PT • Example: Which of the following has the smallest atomic radii? H or Fr—H because it has 1 energy level whereas Fr has 7 energy levels (H is highest up in the group so it is smaller. 7th Floor Main Floor

30. Ionic Radii: The size of the ion (charged atom due to loss or gain of e-) • Ionic radii compare the atom to the ion. • Cations: positive ions are smaller than their corresponding atoms due to a loss of electrons to become charged • Anions: negative ions are larger than their corresponding atoms due to a gain of electrons to become charged • Example: Which is larger F or F-? Mg or Mg2+? • F- because it has more electrons than protons so the electrons spread out • Mg because it has more electrons than protons so the electrons spread out Main Floor 7th Floor

31. Ionization Energy: The energy required to remove an electron • Ionization energy decreases down a group and increasesacross a period. • The decrease down a group is due to the increased distance from the nucleus for the electron being removed. • The increase across a period is due to the increased nuclear pull in an energy level • Example: Who has the higher ionization energy He or Cs? • He because it is smaller (farther right on the PT) so it has a greater nuclear pull and because He has only 1 energy level as opposed to Cs which has 6 energy levels (the increased number of energy levels prohibits the nucleus from holding on to the outer electrons which are being removed to create an ion) Main Floor 7th Floor

32. Electronegativity: The “want” or “pull” on electrons by an atom when in a chemical bond. • F is the most electronegative element on the PT (this is due to its small size and greater nuclear pull) • The closer an atom to F, the more electronegative the atom is • Example: Which is more electronegative, O or Na? • O is closest to F and therefore is more electronegative than Na (O is smaller and has a greater nuclear pull; O is also a nonmetal so it wants electrons whereas Na is a metal and it readily gives away its electrons) Main Floor 7th Floor

33. Electron Affinity: The happiness (decrease in energy) that an atom achieves when it gains an electron. • Electron affinity is greatest for halogens because they will achieve an octet (8 valence electrons in the outermost energy level) if they gain just one electron. • Electron affinity is least for a metal because they wish to lose electrons not gain them. Main Floor 7th Floor