Chapter 4:Arrangement of Electrons in

Atoms

Chemistry

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

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

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

Wave Nature of Light

• Wavelength– Distance between two corresponding points on

adjacent waves– λ– nm

• Frequency– Number of waves that pass a given point in a

specified time– Hz - Hertz

Wave Nature of Light

• Figure 4-1, page 92• Equation

– c=λ– Speed = wavelength * frequency– Indirectly related!

• Spectroscope– Device that separates

light into a spectrum that

can be seen

Particle Nature of Light

• Quantum– Minimum quantity of energy that can be lost

or gained by an atom

• Equation– E=h

• Direct relationship between quanta (particle nature) and frequency (wave nature)

• Planck’s Constant (h)– h=6.626 x 10-34 Js

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!**

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

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– http://

chemmovies.unl.edu/ChemAnime/BOHRQD/BOHRQD.html

• 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.

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

• Spectroscopy– Study of light emitted by excited atoms– Bright line spectrum

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

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

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!!

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

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

Angular Momentum Quantum Number

• What is the shape of the orbital?

• Symbol – l• l = s,p,d,f

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 102-103

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

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

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!

Electron Configuration

Lags 1 behind

Lags 2 behind

Representing Electron Configurations

• Three Notations– Orbital Notation– Electron Configuration Notation– Electron Dot Notation

Orbital Notation

• Uses a series of lines and arrows to represent electrons

• Examples

Orbital Notation

• More examples

Electron Configuration Notation

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

• Long form examples

Electron Configuration Notation

• Short form examples – “noble gas configuration”

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

Electron Dot Notation

• More examples

Summary Questions1. How many orbitals are in a d subshell?2. How many individual orbitals are found in

Principle Quantum #3 (the third main energy level)

3. How many orbital shapes are found in Principle Quantum #2?

4. How many electrons can be found in the fourth energy level?

5. A single 4s orbital can hold how many electrons?

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