Presentation on theme: "Chapter 4: Arrangement of Electrons in Atoms"— Presentation transcript:
1 Chapter 4: Arrangement of Electrons in Atoms Chemistry
2 Development of a New Atomic Model There were some problems with the Rutherford model…It did not answer:Where the e- were located in the space outside the nucleusWhy the e- did not crash into the nucleusWhy atoms produce spectra (colors) at specific wavelengths when energy is added
3 Properties of Light Wave-Particle Nature of Light – early 1900’s A Dual NatureIt was discovered that light and e- both have wave-like and particle-like properties
4 Wave Nature of LightElectromagnetic radiation – form of energy that exhibits wave-like behavior as it travels through spaceElectromagnetic spectrumAll the forms of electromagnetic radiationSpeed of light in a vacuum3.0 x 108 m/s
5 Wave Nature of Light Wavelength Frequency Distance between two corresponding points on adjacent wavesλnmFrequencyNumber of waves that pass a given point in a specified time-uHz - Hertz
6 Wave Nature of Light Figure 4-1, page 92 Equation Spectroscope c=λu Speed = wavelength * frequencyIndirectly related!SpectroscopeDevice that separateslight into a spectrum thatcan be seen
7 Particle Nature of Light QuantumMinimum quantity of energy that can be lost or gained by an atomEquationE=huDirect relationship between quanta (particle nature) and frequency (wave nature)Planck’s Constant (h)h=6.626 x Js
8 Particle Nature of Light PhotonIndividual quantum of light; “packet”The Hydrogen AtomLine emission spectrum (Figure 4-5, page 95)Ground StateLowest energy state (closest to the nucleus)Excited StateState of higher energyEach element has a characteristic bright-line spectrum – much like a fingerprint!**
10 Particle Nature of Light Why does an emission spectrum occur?Atoms get extra energy – ex. voltage – and the e- jumps from ground state to excited stateAtoms return to original energy, e- drops back down to ground stateThe energy is transferred out of the atom in a NEW FORMContinuous spectrumEmission of continuous range of frequenciesLine Emission SpectrumShows distinct lines
11 Bohr Model of the Hydrogen Atom Described electrons as PARTICLES1913 – Danish physicist – Niels BohrSingle e- circled around nucleus in allowed paths or orbitse- has fixed E when in this orbit (lowest E closest to nucleus)Lot of empty space between nucleus and e- in which e- cannot be inE increases as e- moves to farther orbits
12 Bohr Model (cont) ONLY explained atoms with one e- Therefore – only worked with hydrogen!!The principles of his work is applied to the models of other atoms, but the models do not perfectly fit the experimental data.
13 Orbits = The circular paths electrons followed in the Bohr model of the atom SpectroscopyStudy of light emitted by excited atomsBright line spectrum
14 The Quantum Model of the Atom e- act as both waves and particles!!De Broglie1924 – French physiciste- may have a wave-particle natureWould explain why e- only had certain orbitsDiffractionBending of wave as it passes by edge of objectInterferenceOccurs when waves overlap
15 The Quantum Model of the Atom Heisenberg Uncertainty Principle1927 – German physicistIt is impossible to determine simultaneously both the position and velocity of an e-12:28-14:28
17 The Quantum Model of the Atom Schrodinger Wave Equation1926 – Austrian physicistApplies to all atoms, treats e- as wavesNucleus is surrounded by orbitalsLaid foundation for modern quantum theoryOrbital – 3D region around nucleus in which an e- can be foundCannot pinpoint e- location!!
18 Quantum Numbers Quantum Numbers Solutions to Schrodinger’s wave eqn Probability of finding an e-“address” of e-Four Quantum NumbersPrincipleAngular MomentumMagneticSpin
19 Principle Quantum Number Which main energy level? (“shell”)The distance from the nucleusSymbol- nn is normally 1-7Greater n value means farther from the nucleus
20 Angular Momentum Quantum Number What is the shape of the orbital?Symbol – ll = s,p,d,f
21 Magnetic Quantum Number Orientation of orbital around nucleusSymbol – mls – 1p – 3d – 5f – 7Every orientation can hold 2 e-!!A “subshell” is made of all of the orientations of a particular shape of orbitalFigures 4-13, 4-14, 4-15 on page
22 Spin Quantum Number Each e- in one orbital must have opposite spins Symbol – ms+ ½ , - ½Two “allowed” values and corresponds to direction of spin
23 Electron Configuration Electron configurations – arrangements of e- in atomsRules:Aufbau Principle – an e- occupies the lowest energy firstHund’s Rule – place one electron in each equal energy orbital before pairingPauli Exclusion Principle – no 2 e- in the same atom can have the same set of QN14:30-18:25
24 Electron Configuration Representing electron configurationsUse the periodic table to write!Know the s,p,d,f block and then let your fingers do the walking!
25 Electron Configuration Lags 1 behindLags 2 behind
26 Representing Electron Configurations Three NotationsOrbital NotationElectron Configuration NotationElectron Dot Notation
27 Orbital NotationUses a series of lines and arrows to represent electronsExamples
29 Electron Configuration Notation Long Form: Eliminates lines and arrows; adds superscripts to sublevels to represent electronsLong form examples
30 Electron Configuration Notation Short form examples – “noble gas configuration”
31 Electron Dot NotationOuter shell e- - Outermost electrons; In highest principle quantum #Inner shell e- - not in the highest energy levelHighest occupied energy level / highest principle quantum numberValence electrons – outermost e-Examples
33 Summary Questions How many orbitals are in a d subshell? How many individual orbitals are found in Principle Quantum #3 (the third main energy level)How many orbital shapes are found in Principle Quantum #2?How many electrons can be found in the fourth energy level?A single 4s orbital can hold how many electrons?