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Chapter 4: Arrangement of Electrons in Atoms

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1 Chapter 4: Arrangement of Electrons in Atoms

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 nucleus Why the e- did not crash into the nucleus Why 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 Nature It was discovered that light and e- both have wave-like and particle-like properties

4 Wave Nature of Light Electromagnetic radiation – form of energy that exhibits wave-like behavior as it travels through space Electromagnetic spectrum All the forms of electromagnetic radiation Speed of light in a vacuum 3.0 x 108 m/s

5 Wave Nature of Light Wavelength Frequency
Distance between two corresponding points on adjacent waves λ nm Frequency Number of waves that pass a given point in a specified time -u Hz - Hertz

6 Wave Nature of Light Figure 4-1, page 92 Equation Spectroscope c=λu
Speed = wavelength * frequency Indirectly related! Spectroscope Device that separates light into a spectrum that can be seen

7 Particle Nature of Light
Quantum Minimum quantity of energy that can be lost or gained by an atom Equation E=hu Direct relationship between quanta (particle nature) and frequency (wave nature) Planck’s Constant (h) h=6.626 x Js

8 Particle Nature of Light
Photon Individual quantum of light; “packet” The Hydrogen Atom Line emission spectrum (Figure 4-5, page 95) Ground State Lowest energy state (closest to the nucleus) Excited State State of higher energy Each 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 state Atoms return to original energy, e- drops back down to ground state The energy is transferred out of the atom in a NEW FORM Continuous spectrum Emission of continuous range of frequencies Line Emission Spectrum Shows distinct lines

11 Bohr Model of the Hydrogen Atom
Described electrons as PARTICLES 1913 – Danish physicist – Niels Bohr Single e- circled around nucleus in allowed paths or orbits e- 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 in E 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
Spectroscopy Study of light emitted by excited atoms Bright line spectrum

14 The Quantum Model of the Atom
e- act as both waves and particles!! De Broglie 1924 – French physicist e- may have a wave-particle nature Would explain why e- only had certain orbits Diffraction Bending of wave as it passes by edge of object Interference Occurs when waves overlap

15 The Quantum Model of the Atom
Heisenberg Uncertainty Principle 1927 – German physicist It 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 Equation 1926 – Austrian physicist Applies to all atoms, treats e- as waves Nucleus is surrounded by orbitals Laid foundation for modern quantum theory Orbital – 3D region around nucleus in which an e- can be found Cannot pinpoint e- location!!

18 Quantum Numbers Quantum Numbers Solutions to Schrodinger’s wave eqn
Probability of finding an e- “address” of e- Four Quantum Numbers Principle Angular Momentum Magnetic Spin

19 Principle Quantum Number
Which main energy level? (“shell”) The distance from the nucleus Symbol- n n is normally 1-7 Greater n value means farther from the nucleus

20 Angular Momentum Quantum Number
What is the shape of the orbital? Symbol – l l = s,p,d,f

21 Magnetic Quantum Number
Orientation of orbital around nucleus Symbol – ml s – 1 p – 3 d – 5 f – 7 Every orientation can hold 2 e-!! A “subshell” is made of all of the orientations of a particular shape of orbital Figures 4-13, 4-14, 4-15 on page

22 Spin Quantum Number Each e- in one orbital must have opposite spins
Symbol – ms + ½ , - ½ Two “allowed” values and corresponds to direction of spin

23 Electron Configuration
Electron configurations – arrangements of e- in atoms Rules: Aufbau Principle – an e- occupies the lowest energy first Hund’s Rule – place one electron in each equal energy orbital before pairing Pauli Exclusion Principle – no 2 e- in the same atom can have the same set of QN 14:30-18:25

24 Electron Configuration
Representing electron configurations Use 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 behind Lags 2 behind

26 Representing Electron Configurations
Three Notations Orbital Notation Electron Configuration Notation Electron Dot Notation

27 Orbital Notation Uses a series of lines and arrows to represent electrons Examples

28 Orbital Notation More examples

29 Electron Configuration Notation
Long Form: Eliminates lines and arrows; adds superscripts to sublevels to represent electrons Long form examples

30 Electron Configuration Notation
Short form examples – “noble gas configuration”

31 Electron Dot Notation Outer shell e- - Outermost electrons; In highest principle quantum # Inner shell e- - not in the highest energy level Highest occupied energy level / highest principle quantum number Valence electrons – outermost e- Examples

32 Electron Dot Notation More 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?

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