chapter 7 notes atomic structure and periodicity west valley high school ap chemistry mr. mata
TRANSCRIPT
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Chapter 7 Notes Atomic Structure
and Periodicity
West Valley High SchoolAP Chemistry
Mr. Mata
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The Puzzle of the Atom Protons and electrons are attracted to each
other because of opposite charges
Electrically charged particles moving in a curved path give off energy
Despite these facts, atoms don’t collapse
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Wave-Particle DualityJJ Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron
is a particle!
The electron is an energy
wave!
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Confused??? You’ve Got Company!
“No familiar conceptions can be woven around the
electron; something unknown is doing we don’t
know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
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The Wave-like Electron
Louis deBroglie
The electron propagates through space as an
energy wave. To understand the atom, one must understand
the behavior of electromagnetic waves.
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c = C = speed of light, a constant (3.00 x 108 m/s)
= frequency, in units of hertz (hz, sec-1)
= wavelength, in meters
Electromagnetic radiation propagates through space as a wave moving at the speed of light.
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Types of electromagnetic radiation:
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E = hE = Energy, in units of Joules (kg·m2/s2)
h = Planck’s constant (6.626 x 10-34 J·s)
= frequency, in units of hertz (hz, sec-1)
The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation.
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Long Wavelength
=Low Frequency
=Low ENERGY
Short Wavelength
=High Frequency
=High ENERGY
Wavelength Table
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Relating Frequency, Wavelength and Energy
c hE
hc
E
Common re-arrangements:
E
hc
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…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the visible spectrum…
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…produces a “bright line” spectrum
Spectroscopic analysis of the hydrogen spectrum…
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This produces bandsof light with definitewavelengths.
Electron transitionsinvolve jumps of definite amounts ofenergy.
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Electron Energy in Hydrogen
2
21810178.2
n
ZJxEelectron
***Equation works only for atoms or ions with 1 electron (H, He+, Li2+, etc).
Z = nuclear charge (atomic number)
n = energy level
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Calculating Energy Change, E, for
Electron Transitions
2
2
2
21810178.2
initialfinal n
Z
n
ZJxE
Energy must be absorbed from a photon (+E) to move an electron away from the nucleus
Energy (a photon) must be given off (-E) when an electron moves toward the nucleus
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Quantum Numbers
Each electron in an atom has a unique set of 4 quantum numbers which describe it.
Principle quantum number Angular momentum quantum number
Magnetic quantum number Spin quantum number
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Pauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
Wolfgang Pauli
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Principal Quantum NumberGenerally symbolized by n, it denotes the shell (energy level) in which the electron is located.
Number of electrons that can fit in a shell:
2n2
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Angular Momentum Quantum Number
The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.
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Magnetic Quantum NumberThe magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.
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Assigning the Numbers The three quantum numbers (n, l, and
m) are integers. The principal quantum number (n)
cannot be zero. n must be 1, 2, 3, etc. The angular momentum quantum
number (l ) can be any integer between 0 and n - 1.
For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can
be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1,
+2.
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Principle, angular momentum, and magnetic quantum numbers: n, l, and ml
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Spin Quantum NumberSpin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:
1
2
1
2
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Orbital shapes are defined as the surface that contains 90% of the total electron probability.
An orbital is a region within an atom where thereis a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level…
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Schrodinger Wave Equation
22
2 2
8dh EV
m dx
Equation for probability of a single electron being found along a single axis (x-axis)Erwin Schrodinger
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Heisenberg Uncertainty Principle
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
“One cannot simultaneously determine both the position and momentum of an electron.”
WernerHeisenberg
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Orbitals of the same shape (s, for instance) grow larger as n increases…
Nodes are regions of low probability within an orbital.
Sizes of s orbitals
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Orbitals in outer energy levels DO penetrate intolower energy levels.
This is a probabilityDistribution for a 3s orbital.
What parts of thediagram correspondto “nodes” – regionsof zero probability?
Penetration #1
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Which of the orbital types in the 3rd energy levelDoes not seem to have a “node”?
WHY NOT?
Penetration #2
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The s orbital has a spherical shape centered aroundthe origin of the three axes in space.
s orbital shape
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There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.
P orbital shape
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Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of:
…and a “dumbell with a donut”!
d orbital shapes
“double dumbells”
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Shape of f orbitals
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Orbital filling table
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Electron configuration of the elements of the first three
series
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Element Configuration notation
Orbital notation Noble gas notation
Lithium 1s22s1 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s1
Beryllium 1s22s2 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2
Boron 1s22s2p1 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p1
Carbon 1s22s2p2 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p2
Nitrogen 1s22s2p3 ____ ____ ____ ____ ____
1s 2s 2p
[He]2s2p3
Oxygen 1s22s2p4 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p4
Fluorine 1s22s2p5 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p5
Neon 1s22s2p6 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p6
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Irregular confirmations of Cr and Cu
Chromium steals a 4s electron to halffill its 3d sublevel
Copper steals a 4s electron to FILL its 3d sublevel
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Half of the distance between nucli in covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period Increased effective nuclear charge dueto decreased shielding
Radius increases down a group Addition of principal quantum levels
Determination of Atomic Radius
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Table of Atomic Radii
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Increases for successive electrons taken from the same atom
Tends to increase across a period
Electrons in the same quantum level do not shield as effectively as electrons in inner levels
Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group
Outer electrons are farther from thenucleus
Ionization Energy: the energy required to remove an electron from an atom
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Mg + 738 kJ Mg+ + e-
Mg+ + 1451 kJ Mg2+ + e-
Mg2+ + 7733 kJ Mg3+ + e-
Ionization of Magnesium
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Table of 1st Ionization Energies
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Affinity tends to increase across a period
Affinity tends to decrease as you go down in a period
Electrons farther from the nucleusexperience less nuclear attraction
Some irregularities due to repulsive forces in the relatively small p orbitals
Electron Affinity - the energy change associated with the addition of an electron
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Table of Electron Affinities
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A measure of the ability of an atom in a chemicalcompound to attract electrons
Electronegativities tend to increase across a period
Electronegativities tend to decrease down a group or remain the same
Electronegativity
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Periodic Table of Electronegativities
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Cations Positively charged ions
Smaller than the corresponding
atomAnions Negatively charged ions Larger than the corresponding
atom
Ionic Radii
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Table of Ion Sizes
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Summary of Periodic Trends