2 ElectronsRutherford described the dense center of the atom called the nucleus.But the Electrons spin around the outside of that nucleus.Provide the chemical properties of the atoms.Responsible for color and reactivity.
3 Energy Energy is transmitted from one place to another. Light carries this energy.Converted into heat.Light is called Electromagnetic Radiation.
4 Electromagnetic Spectrum RadioInfraredVisible LightROY G BIVUltravioletX RaysGamma Rays
5 Light Light travels as a wave. Wave Properties Wavelength (λ) = distance between two waves (m)Frequency (f) = number of peaks per second (Hz)Speed of Light (c) = how fast light moves.
6 Light c= ƒλ Light Equation Speed of light is a constant = 3 x 108 m/s Nothing travel faster than the speed of light!Maybe?!?!?!?!?!?!?!?!?
7 Light The Dual Nature of Light Light carries energy through space like a wave.Light also behaves like a particle?!?A beam of light is made of tiny packets of energy called PHOTONS!Which travel in waves!?!
8 Light E = hƒ The Energy of a photon depends on its frequency. So is the color of light!!!E = hƒELECTRONS are like photons!Act as waves and particles.Orbit the nucleus in a wave-like motion.
9 Blackbody RadiationRutherford could never explain why objects change colors when they are heated.As the object heats, it must give off electrons of certain frequencies and energies.
10 Photoelectric EffectSimilarly, light on a metal object can knock off electrons.Shine different colors on a metal.Measure the number of electrons knocked off.Found that no electrons were knocked off below a certain frequency.
11 The Bohr ModelProposed the electrons orbit the nucleus with fixed energies.Called Energy LevelsMuch like the rungs of a ladder.Quantum describes the amount of energy required to move an electron from one level to another.
12 The Bohr Model Ground State Excited State Lowest possible energy of an electron.Normal locationExcited StateIf electron absorbs energy, it moves up an energy level (absorption)If an electron gives off energy, it moves down an energy level (emission).
14 Atomic Spectra Hydrogen Atom Line Emission Spectrum Expected continuous spectrum of lightBut only specific frequencies were given off.Red (656.6 nm)Blue-green (486.1 nm)Violet (434.1 nm)Violet (419.2 nm)
15 Atomic Spectra Shine a light on an Atom Atomic Spectra When atoms absorb energy, electrons move to higher energy levels.When atoms release the energy, electrons return to the lower energy level.Atomic SpectraFrequencies of light emitted by a certain element.No two elements have the same spectrum.
16 Flame TestsBecause no two atoms produce the same spectrum, elements can be identified by the colors they emit.Spectral Analysis uses this properties to identify elements.
17 E = hf Quantum Mechanics Max Planck (1900) Albert Einstein (1905) Founder of Quantum MechanicsE = hfAlbert Einstein (1905)Wave-Particle DualityElectrons are small particles that move like waves.
18 mv/λ = h Quantum Mechanics Neils Bohr (1922) Louis de Brogelie (1923) Electrons orbit in distinct energy levels.Louis de Brogelie (1923)Wave Mechanics says that ALL MATTER behaves like waves.mv/λ = h
19 Quantum Mechanics Werner Heisenberg (1927) Erwin Schrödinger (1930) Principle of IndeterminacyYou can’t know both the position and the velocity of an electron.Erwin Schrödinger (1930)Used wave mechanics to show the PROBABLE location of an electron.Electrons exist in 3D clouds of probability!!!
21 Quantum Mechanical Model Uses Schrodinger’s equation to predict the probable location of an electron.Determines the energies an electron is allowed to have.Determines how likely it is to find the electron in various locations around the nucleus.
22 Quantum Numbers Describes the location and behavior of an electron Like an electron’s addressNo two electrons can have the same quantum numbers.Four Numbers
23 Quantum Numbers Principle (1st) Quantum Number (n) The Energy Level Describes the size of the cloud and the distance of the cloud from the nucleus.Shows the number of electronsn = 1 = 2 electronsn = 2 = 8 e-n = 3 = 18 e-n = 4 = 32 e-
25 Quantum Numbers 2nd Quantum Number (l) s = spherical p = peanut-shaped Each energy level has sublevels.The number of sublevels equals n.Sublevels are called:s = sphericalp = peanut-shapedd = daisy-shapedf = unknown?
27 Quantum Numbers s = 1 orbital p = 3 orbitals d = 5 orbitals 3rd Quantum Number (ml)Divides sublevels into orbitals.Tells the shape the electron moves in.Number of orbitals = n2Exampless = 1 orbitalp = 3 orbitalsd = 5 orbitalsf = 7 orbitals
28 Quantum Numbers 4th Quantum Number (ms) Describes the electron’s spin. Only two electrons fit in an orbital.Their charges repel causing them to spin in opposite directions (+½ or –½)Use up and down arrows.
30 Quantum Numbers Pauli Exclusion Principle Hund’s Rule No two electrons can have the same set of 4 quantum numbers.The electrons repel each other.Hund’s RuleEvery orbital must get one electron before doubling up.
31 Quantum Numbers 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f Diagonal RuleElectrons fill orbitals in predictable patternsSome People Do ForgetElectrons dill the lowest energy level possible.2s 2p3s 3p 3d4s 4p 4d 4f5s 5p 5d 5f
32 Orbital NotationDraw out the locations of each electron in an atom with arrows.
33 Electron Configuration Write out the configurations of electrons using superscripts.Examples:H = 1s1He = 1s2
34 Electron Configurations Noble Gas ShorthandWrite the Noble Gas just before the element.Add the remainder of the configuration.
35 Lewis Dot DiagramsA way to show the number and position of the valence electrons.Outermost energy levelLook at the column number to get this number.Use the chemical symbol and number of valence electrons.All four sides must have a dot before you double up.p orbitalss orbitalXp1sp2p3