Transcript Chapter 4

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 wavelike 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
-u
– Hz - Hertz
Wave Nature of Light
• Figure 4-1, page 92
• Equation
– c=λu
– 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=hu
• 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 ejumps 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/B
OHRQD.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 Questions
1. 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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