Chem 101 Chapter 2 Notes : Bethel
Chem 101 Chapter 2 Notes : Bethel CHEM 101
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Popular in Chemistry
This 10 page Class Notes was uploaded by AnnaBanana on Sunday September 18, 2016. The Class Notes belongs to CHEM 101 at Texas A&M University taught by Dr. Ryan Bethel in Fall 2016. Since its upload, it has received 30 views. For similar materials see Chemistry 101 in Chemistry at Texas A&M University.
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Date Created: 09/18/16
CHEM 101: LECTURE (CHAPTER 2) Highlight – Important Concepts Highlight – Important Principles Highlight – Key Terms 2.1ElectromagneticRadiation - James Maxwell stated that electromagnetic radiation was energy traveling through space. - It has wavelength and frequency - Made of oscillating electric and magnetic fields traveling through space. Electromagnetic radiation – oscillating waves and magnetic fields traveling through space. Wavelength – (λ) distance between 2 peaks and 2 troughs of a wave. Frequency – (v) points to # of waves (cycles) per second that pass a given point in space. Velocity – ( c ) all light has the same speed in a vacuum. (Speed of Light ) ???? = 2.998 × 10 8 Speed of Light Equations ???? = ???????? Electromagnetic Spectrum: Longer Wavelength Shorter Wavelength Lower Frequency Higher Frequency Lower Energy Higher Energy Radio Wave Microwave Infrared Visible Light Ultraviolet X- Rays Gamma Rays Wave Wave Visible light is only a small portion of the actual electromagnetic spectrum. *Dr. Bethel says : Know the order of the electromagnetic spectrum, but not the wavelength/frequency cutoffs. 2.2NatureOfMatter Planck – energy can be gained or lost in whole number multiples of hv. H = Planck’s Constant = 6.626 × 10 −34 ???? · ???? 1. Energy’s Movement a. Energy moves in little ‘packets’ called quanta. Atoms can only trade quantas of energy. 2. Photons a. Photons are streams of ‘particles’ of light. Lots of photons make up electromagnetic radiation i. When the frequency of light increases, the photon’s energy increases. ii. When the frequency of light decreases, the photon’s energy decreases. ???????? E = a single photon’s energy E = ???????? = h = 6.626 × 1034???? · ???? ???? ???? = frequency C = speed of light λ = wavelength ???? 1. Photoelectric Effect a. Light striking a metal surface causes electrons to be ejected from that surface. i. But the light must be above the threshold frequency,???????? b. When Electrons Emit: i. Never, when ???? < ???? 0 no electrons come off of the metal. ii. Yes, According to Brightness of light, when ???? > ???? 0 Also, the KE of the electrons ejected increases along with the frequency of the light. 2. Wave-Particle Duality a. According to situation, light can be described as a particle or a wave. i. Particle when light is giving energy to another object. 1. Photon – particle of light b. All properties of light can’t be encompassed by just waves and particles. 3. Binding Energy a. Energy that holds an electron to a metal; minimum energy required to remove an electron from the surface. i. ????????= ???????? ???? 4. Diffraction a. Light getting scattered from a regular array of points or lines. i. Destructive - waves coinciding with 1 peak & 1 trough. Dark patches. ii. Constructive – waves coinciding with 2 peaks or 2 troughs. Bright Spot. 5. Wave Particle Duality with Light a. Waves can act like particles and particles can act like waves. i. Well then the waves must have a wavelength! Wavelength of A Particle ** Wave equations are ALWAYS in Kg, so make sure to convert masses. ???? ???? = ???????? v = velocity m – mass of particle h – planck’s constant. 2.3AtomicSpectrumofHydrogen AtomicSpectra - Light coming from tube with2, put energy in it, then it gets energized and gives off light. o Gives off only 4 lines of light because of the Atomic Spectra. - Light Spectrum - # and type (color) of light given off by excited gases. - Prisms work opposite than the tube of light. Instead of creating lines of light, they block lines of light. o The lines of blocked light line up according to element. ** Each element has its own particular pattern of wavelengths it emits and absorbs. ** Atoms can only have certain energy levels. When light is emitted, the atom goes down in energy. This energy is those ‘packets’ of light particles, photons. 2.4BohrModel QuantumModelforHydrogenAtom - Electron moves around nucleus in a specific order. o By moving down a level, closer to the nucleus, it releases energy. (lightbulb’s glow) o By moving up a level, farther from the nucleus, it takes energy. (prism restricting types of light from shining). ????2 E = -2.178 x 10 -18J (????2 N – energy level Z – charge of nucleus (usually 1) - Energy Levels in the nucleus resemble an arrangement of highways around a city. o The amount of rings change according to the type of element. - All rings aren’t visible. Any element has an infinite number of rings, but there aren’t any electrons to orbit on them, so we don’t count them. o Electrons can skip when moving electron levels. (ex. N= 1 to N=3) CalculatingChangeinEnergy(????????) ???????? = E –fE i 1 1 ???????? = −2.178 ∗ 10 −18 [ − ] ????????2 ???? ???? Is the same as 1 1 1 = ???? ℎ − ] ???? ???? ???? ???? ???? ???? RydbergConstant = 1.094* 101m -1 ℎ= −???? ???? ???? ???? Energy of Specific “n” level. ????2 BohrModel - Finds the correct energy levels for hydrogen only. - A free electron is at infinite distance from the nucleus. o Ground state – lowest possible energy level o Excited State – any energy level other than ground. - When an electron gets closer to the nucleus, it gives off more energy, making it ‘more negative’ than the free electron. 2.5Uncertainty&QuantumMechanicalModel HeisenbergPrinciple - We can never know the location and momentum of a particle at the same time. o If we know how fast it is, we don’t know where it is. o If we know where it is, we don’t know how fast it is. ℎ Δ???? ∗ Δ ???????? ≥ ) 4???? SchroedingerEquation - Matter can be described as both a wave and a particle. - The term ‘wave function’ has no real meaning in that it is simply the probability of finding an electron in a specific region of space. It has no substance. ???????? = ???????? 2.6QuantumNumbers Principle quantum # (n) – integral values and is related to the size and energy of the orbital. Angular momentum quantum # - (l) integral values from 0 to (n-1). Determines the shape of the atomic orbital. Magnetic Quantum # (m) has integral values between l and -l, including 0. Related to the position of the orbital in space. N L value M l 1 0 0 2 0,1 -1,0,1 3 0,1,2 -2,-1,0,1,2 4 0,1,2,3 -3,-2,-1,0,1,2,3 - Each l within an “n-level” is called a subshell. (Because it’s inside a shell, aka layer of atom). - Each l subshell is divided into orbitals. o Orbital – pair of atoms that orbit the nucleus together. L # Type of Orbital # of Orbital 0 “s” – spherical 1 (2 atoms) 1 “p” – principal 3 (6 atoms) 2 “d” 5 (10 atoms) 3 “f” 7 (14 atoms) 2.7OrbitalShapesandEnergies Nodes - Electrons won’t be in regions of space called nodes. o There is zero probability of finding them here. - To find the number of nodes in an n-level, use “N – 1= # total nodes” formula. - 2 types of nodes o Angular (planar) nodes – “slice” of angled 3D plane. This has a measurement and is not round. It is a straight line that cuts. o Radial (Spherical) nodes – If atom with a nucleus is the origin of an xy plane, then the node would be a circle centered on the graph where the electrons cannot exist. To Find : n – l – 1 - Nodes give orbitals shapes (s, p, d, f!) SOrbitals - Only one s orbital that extends from nucleus in a radial manner, forming a spherical shape. - Only one S orbital for each energy level. - No nodes because (n – L – 1 = 1 – 0 – 1 = 0 , so 0 nodes) o L – 0 o M = L POrbitals - Peanut/dumbbell shape - Two lobes separated by a node at the nucleus. - Labeled like the Cartesian plane (xyz) - Always have 1 angular node. # of radial nodes depends on “n”. (n-L-1) o L = 1 o M = L1,0,1 DOrbitals - N is always at least 1 more than L - Shaped like four leaf clovers or (in one case) a doughnut. o L = 2 o M = L2, -1, 0, 1, 2 2.8ElectronSpin(4 Quantum#) Electron spin quantum # (m ) s can be either +1/2 or -1/2. Electron spin is just as important as charge, and it can spin either “up” or “down”. - Called spin b/c in research, atoms either move away from an external magnet (+1/2) or toward it (-1/2) Pauli Exclusion Principle – In a given atom, no two electrons can have the same spin numbers. 2.9PolyelectronicAtoms Polyelectronic atoms – atoms with more than one electron. (That is all elements except hydrogen). - Since electron patterns are unknown, electron repulsions can’t be accurately calculated. Degenerate Orbital – An orbital that has the same energy as another orbital. ElectronAssignments 1. Electrons in lower (n + L) values are more stable because they’re lower in energy and ‘glued’ through magnetic attraction to the nucleus. 2. If (n+l) for both electron levels equal each other, the level with the lower “n” value is more stable because it is still closer to the nucleus. Electron Level (N & Type) N L N + L 2s 2 0 2 2p 2 1 3 3s 3 0 3 3p 3 1 4 3d 3 2 5 Lowest E : 2s , 2p, 3s, 3p, 3d : Highest E PenetrationofShielding - Inner electrons shield outer electrons from the nucleus. o S and some d can penetrate the defense of the closer electron orbitals. o More radial nodes means that the inner electrons are closer to the nucleus. 2.10PeriodicTable - Let’s organize electrons into groups of elements because this shows similarities between electrons. Mendeleev - Repeating trend to table of elements - Estimate mass and property of unknown elements - USED ATOMIC MASS! (wrong!!) (We use atomic #!!!) 2.11AufbauPrinciple Aufbau Principle – Lower energy Orbitals fill first. Pauli Exclusion Principle – Individual Orbitals hold only two electrons, which should have different spin. Hund’s Rule – all degenerate orbitals are filled with electrons until half full before pairing of electrons happens. ElectronConfigurationandOrbitalBoxNotation - Electrons fill orbitals from lowest to highest energy - Electron configuration of an atom is the total sum of the electrons from lowest to highest shell. ElectronConfigurations - System indicates electrons within atom. 1. First by # (1, 2, 3, 4) a. Period # (n) 2. Lower Case Letter (s, p, d, f) a. Orbital Type (L) 3. Superscript over letter showing # of electrons in orbital. ValenceElectrons - Electrons on the outermost quantum level of the atom. o These are the unpaired electrons who don’t make complete orbitals. - Atoms of elements in the same group have the same number of valence electrons. (Key to similar chemical behavior!) ** We can tell the number of valence electrons (aka Electron shells not filled to capacity) by the location on the periodic table. Clicker Question NobleGasNotation - Noble gases are very inert. They have no unfilled orbitals, and all of the electrons are paired. - Core orbitals of elements in the periodic table are Noble Gases. - We abbreviate the core orbitals of elements with Noble Gas symbols in order to save space and avoid stating what we already know. - TransitionElements - They are written in the order that they are filled, which isn’t in the numerical and letter order of the shells. So, in these cases, 4s comes before 3d because it has less energy, and thus, is filled previously. ElectronConfigurationofIons - Write an electron configuration of the neutral atom. o Then, remove or add electrons from the orbital with the HIGHEST N! (Even if 3d is the last entity in the notation, you subtract/add electrons to the 4s first!) - Cations – subtracting electrons. Less electrons. We subtract electrons from the orbital with the highest “n”. - Anions - adding electrons. We just add the electrons to the electron count. ExceptionstotheRuleofAddinginOrder Howdy, Y’all! Thanks for using my notes! I hope they help you study for CHEM 101. Please note that these are my notes and you shouldn’t be passing them off as yours. Find me on StudySoup! (There, you can buy all my notes & study guides for the whole semester!) Please let me know if you have any questions! Thanks, Anna
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