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**Chapter 4: Arrangement of Electrons in Atoms**

Chemistry

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**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

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**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

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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

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**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

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**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

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**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

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**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!**

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**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

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**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

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**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.

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**Orbits = The circular paths electrons followed in the Bohr model of the atom**

Spectroscopy Study of light emitted by excited atoms Bright line spectrum

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**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

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**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

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**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!!

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**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

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**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

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**Angular Momentum Quantum Number**

What is the shape of the orbital? Symbol – l l = s,p,d,f

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**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

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**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

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**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

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**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!

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**Electron Configuration**

Lags 1 behind Lags 2 behind

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**Representing Electron Configurations**

Three Notations Orbital Notation Electron Configuration Notation Electron Dot Notation

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

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Orbital Notation More examples

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**Electron Configuration Notation**

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

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**Electron Configuration Notation**

Short form examples – “noble gas configuration”

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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

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Electron Dot Notation More examples

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**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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