Electronic Structure and the Periodic Table
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Transcript of Electronic Structure and the Periodic Table
Electronic Structureand the Periodic Table
Unit 6 Honors Chemistry
Electromagnetic Waves:
Electromagnetic waves:progressive, repeating disturbances that come from the movement of electric charges
Electromagnetic Waves & Light
Wavelength and FrequencyWavelength and Frequency
Wavelength (, lambda): distance between any two points in a wave
measured in any distance unit
(mainly nm or m:
1 nm = 1x10-9 m)
Wavelength Can be Measured in One of Two Ways…
Wavelength and FrequencyWavelength and Frequency
Frequency (; pronounced nu): the number of cycles of the wave that pass through a
point in a unit of timeMeasured in sec-1 (/sec)
1 sec-1 = 1 Hertz (Hz)
Illustration of Frequency
Wavelength is indirectly proportional to frequency
As Wavelength increases, frequency _________________.
As Wavelength decreases, frequency _________________.
AmplitudeAmplitude
Note: height of wave is amplitude (intensity or brightness of wave)
Amplitude is INDEPENDENT of frequency or wavelength!
SpeedSpeed
Speed (c): The speed of light!
c = 3.00 x 108 m/s
(rounded to 3 sig figs)
Equation Equation
One equation relates speed, frequency and wavelength:
c =
Example The wavelength of the radiation which
produced the yellow color of sodium vapor light is 589.0 nm. What is the frequency of this radiation?
The electromagnetic spectrum complete range of wavelengths and frequencies mostly invisible
What is color?
TED Talk: What is color?
The visible/continuous spectrum
continuous spectrum: components of white light split into its colors, ROY G BIV from 390 nm (violet) to 760 nm (red)
Line Spectra
Pattern of lines produced by light emitted by excited atoms of an elementunique for every elementused to identify unknown elements
How do we see color?
TED Talk: How we see color
Max Planck
Light is generated as a stream of particles called PHOTONS
Equation: E (Energy of a photon)= h
(h =Plank’s constant= 6.626x10-34Js)
Low frequency, high Low frequency, high λλ, low E., low E.Low frequency, high Low frequency, high λλ, low E., low E.
High frequency, low High frequency, low λλ, high E., high E.High frequency, low High frequency, low λλ, high E., high E.
E = h • E = h • E = h • E = h •
Relationships in Planck’s Eqn.Relationships in Planck’s Eqn.Relationships in Planck’s Eqn.Relationships in Planck’s Eqn.
Photoelectric effect – Nobel Prize in Physics 1921 to Einstein
Occurs when light strikes the surface of a metal and electrons are ejected.
Practical uses:
Automatic
door openers
Photoelectric Effect: ConclusionPhotoelectric Effect: Conclusion
Light not only has Light not only has wave wave propertiesproperties but also has but also has particle particle
properties. These massless properties. These massless particles, called particles, called photonsphotons, , are are
packets of energy.packets of energy.
Example 6.2
Using the frequency calculated in the previous example, calculate the energy, in joules, of a photon emitted by an excited sodium atom. Calculate the energy, in kilojoules, of a mole of excited sodium atoms.
Bohr’s Hydrogen Atom: Bohr’s Hydrogen Atom: A Planetary Model A Planetary Model
Niels Bohr: ProposedNiels Bohr: Proposedplanetary model.planetary model.
Electrons “orbit” the nucleus like planets Electrons “orbit” the nucleus like planets around the sun.around the sun.
NOT current model of atom but used to NOT current model of atom but used to explain some features of atom.explain some features of atom.
Ground State vs. Excited State
ground state: all electrons in lowest possible energy levels
excited state: an electron that has absorbed energy and moved to a higher energy levelThis is a temporary state!!
Explanation of Line Spectra & Equation
Niels Bohr
Energy of an electron is quantized: can only have specific values.
Energy proportional to energy level.
Explanation of Line Spectra
Electron will drop from excited state to ground state and will emit energy as a photon.
Explanation of Line Spectra
Type of photon emitted by electron depends on energy difference of energy levels
Elevel = -RH 1 – 1 (nhi)2 (nlow)2
AND Elevel = h = hc/
(h: Planck’s constant, 6.626 x 10-34 J sec/photon)
Flaw in Bohr’s Model
Only works well for 1 electron species (H atom).
Does not explain fine structure of line spectra.
Wave-Particle Duality
Light has properties of both WAVES and PARTICLES.
most matter has undetectable wavelengths (1000 kg car at 100 km/hr has = 2.39 x 10-38 m)
This work led to the development of the electron microscope
Quantum Mechanics
Quantum mechanics:
atomic structure based on wave-like properties of the electron
Schrödinger: wave equation that describes hydrogen atom
Heisenberg Uncertainty Principle
The exact location of an electron cannot be determined (if we try to observe it, we interfere with the particle)
You can know either the location or the velocity but not both
Electrons exist in electron clouds
and not on specific rings or orbits
Quantum Numbers Four quantum numbers are a mathematical
way to represent the most probable location of an electron in an atom
analogy...state = energy level, n
city = sublevel, l
address = orbital, ml
house number = spin, ms
Principal Quantum Number: n
Always a positive integer (1,2, 3…7)
Indicates size of orbital, or how far electron is from nucleus
Similar to Bohr’s energy levels or shells
Larger n value = larger orbital or distance from nucleus
The Periodic Table and The Periodic Table and ShellsShells
The Periodic Table and The Periodic Table and ShellsShells
n = 1n = 1
n = 2n = 2
n = 3n = 3
n = 4n = 4
n = 5n = 5n = 6n = 6n = 7n = 7
n = row number on periodic table for a given element
Angular Momentum Quantum Number: l
positive integer from zero to n-1 Sublevel within an energy level; indicates
shape of orbital 0 = s 1 = p 2 = d 3 = f
Types of Types of SublevelsSublevels
s s p p d d
Magnetic Quantum Numbers: ml
integer from -l to +l Indicates orientation of orbital in space
Orbital = electron containing area
Spin Quantum Number: ms
Two values only: + ½ or -½ 2 electrons max. allowed in each orbital (Pauli Exclusion Principle)
Indicates spin of electron; spins of each electron must be opposite
n ---> level1, 2, 3, 4, ...
l ---> sublevel 0, 1, 2, ... n - 1
ml ---> orbital -l ... 0 ... +l
ms ---> electron spin +½ and -½
n ---> level1, 2, 3, 4, ...
l ---> sublevel 0, 1, 2, ... n - 1
ml ---> orbital -l ... 0 ... +l
ms ---> electron spin +½ and -½
REVIEW: REVIEW: QUANTUM NUMBERSQUANTUM NUMBERSEvery Electron has four!Every Electron has four!
OrbitalsOrbitalsNo more than 2 e- assigned to an orbitalNo more than 2 e- assigned to an orbitalOrbitals grouped in s, p, d (and f) Orbitals grouped in s, p, d (and f)
subshellssubshells
s orbitalss orbitals
d orbitalsd orbitals
p orbitalsp orbitals
Capacities of levels, sublevels, and orbitals—see packet
Example
Example 6.6 Give the n and l values for the following orbitals:
a. 3p
b. 4s
Example
Example 6.8 What are the possible ml values for the following orbitals:
a. 3p
b. 4f
Shapes of Atomic Orbitals
Shapes of Atomic Orbitals
s = spherical
p = peanut
d = dumbbell (clover)
f = flower
Multielectron Atoms
In the hydrogen atom the subshells (sublevels) of a principal energy level or shell are at the same energy level.
Previous Equation: En = –RH /n2
Multielectron Atoms
In a multielectron atom, only the orbitals are at the same energy level: the sublevels are at different energy levels!
The increasing energy order of sublevels is generally:
s < p < d < f
Overlapping subshells
At higher energy levels, sublevels overlap.
Note:
4s vs. 3d!
Introduction to Electron Configuration
Definition: describes the distribution of electrons among the various orbitals in the atom
Represents the most probable location of the electron!
EOS
Electron Configurations
The system of numbers and letters that designates the location of the electrons
3 major methods: Full electron configurations Abbreviated/Noble Gas configurations Orbital diagram configurations
Full or Complete Electron Configuration (uses spdf)
Uses numbers to designate a principal energy level and the letters to identify a sublevel; a superscript number indicates the number of electrons in a designated sublevel.
EOS
Rules for Electron Configurations
The Aufbau principle:Electrons fill from the lowest energy level to the highest (they don’t skip around)
1s22s22p63s23p64s23d10etc.
Pauli Exclusion Principle
No two electrons in the same atom can have the same set of 4 quantum numbers.
That is, each electron has a unique “address”
In other words, the maximum # of electrons an orbital can hold is 2 e- (one with ms = +1/2 and one
with ms = -1/2)
HUND’S RULE
Orbitals of equal energy in a sublevel must all have 1 electron before the
electrons start pairing up
a.k.a “creepy person on the bus rule”
*** also electrons in half-filled orbitals have same spin
Why are these incorrect?
Why are these incorrect?
Why are these incorrect?
Full Electron Configuration
Example Notation: 1s2 2s1 (Pronounced “one-s-two, two-s-one”)
A. What does the coefficient mean?Principle energy level
B. What does the letter mean?Type of orbital (sublevel)
C. What does the exponent mean?# of electrons in that orbital
Steps to Writing Full Electron Configurations
1. Determine the total number of electrons the atom has (for neutral atoms it is equal to the atomic number for the element).
Example: F atomic # = # of p+ = # of e- =
2. Fill orbitals in order of increasing energy (see Aufbau Chart).
3. Make sure the total number of electrons in the electron configuration equals the atomic number.
Aufbau Chart (Order of Energy Levels)
When writing electron configurations:
d sublevels are n – 1 from the row they appear in
f sublevels are n – 2 from the row they appear in
Writing Electron Configurations
Writing Electron Configurations
Nitrogen:
Helium:
Phosphorous:
Rhodium:
Bromine:
Cerium:
Abbreviated/Noble Gas Configuration
i. Where are the noble gases on the periodic table?
ii. Why are the noble gases special?
iii. How can we use noble gases to shorten regular electron configurations?
Abbreviated/Noble Gas Configuration
Example: Barium
1.Look at the periodic table and find the noble gas in the row above where the element is.
2.Start the configuration with the symbol for that noble gas in brackets, followed by the rest of the electron configuration.
Abbreviated/Noble Gas Configuration
Practice! Write Noble Gas Configurations for the following elements:
Rubidium:
Bismuth:
Arsenic:
Zirconium:
Writing Electron Configurations
Writing Electron Configurations
Another way of writing
configurations is called an orbital
diagram.
(also called orbital notation or orbital
diagrams)
Another way of writing
configurations is called an orbital
diagram.
(also called orbital notation or orbital
diagrams)
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
Arrowsdepictelectronspin
ORBITAL BOX NOTATIONfor He, atomic number = 2
1s
21 s
One electron has n = 1, One electron has n = 1, ll = 0, m = 0, mll = 0, = 0, mmss = + = + ½½
Other electron has n = 1, Other electron has n = 1, ll = 0, m = 0, mll = 0, = 0, mmss = - = - ½½
Orbital Diagrams
Orbital diagrams use boxes (sometimes circles) to represent energy levels and orbitals. Arrows
are used to represent the electrons.
= orbital
sublevels
Orbital Diagrams
Don’t forget - orbitals have a capacity of two electrons!! Two electrons in the same orbital must have opposite
spin so draw the arrows pointing in opposite directions.
Example: oxygen 1s22s22p4
1s
2s
2p
Incr
easi
ng E
nerg
y
Drawing Orbital Diagrams
1. First, determine the electron configuration for the element.
2. Next draw boxes for each of the orbitals present in the electron configuration.
Boxes should be drawn in order of increasing energy (see the Aufbau chart).
3. Arrows are drawn in the boxes starting from the lowest energy sublevel and working up. This is known as the Aufbau principle.
Add electrons one at a time to each orbital in a sublevel before pairing them up (Hund’s rule)
The first arrow in an orbital should point up; the second arrow should point down (Pauli exclusion principle)
4. Double check your work to make sure the number of arrows in your diagram is equal to the total number of electrons in the atom.
# of electrons = atomic number for an atom
Electron Configurations for Nitrogen
Electron Configurations for Nitrogen
Electron Configurations for Nickel
Electron Configurations for Nickel
LithiumLithium
Group 1A
Atomic number = 3
1s22s1 ---> 3 total electrons
1s
2s
3s3p
2p
BerylliumBeryllium
Group 2A
Atomic number = 4
1s22s2 ---> 4 total electrons
1s
2s
3s3p
2p
BoronBoron
Group 3A
Atomic number = 5
1s2 2s2 2p1 --->
5 total electrons
1s
2s
3s3p
2p
CarbonCarbon
Group 4A
Atomic number = 6
1s2 2s2 2p2 --->
6 total electrons
Here we see for the first time Here we see for the first time
HUND’S RULEHUND’S RULE. .
1s
2s
3s3p
2p
NitrogenNitrogen
Group 5A
Atomic number = 7
1s2 2s2 2p3 --->
7 total electrons
1s
2s
3s3p
2p
OxygenOxygen
Group 6A
Atomic number = 8
1s2 2s2 2p4 --->
8 total electrons
1s
2s
3s3p
2p
FluorineFluorine
Group 7A
Atomic number = 9
1s2 2s2 2p5 --->
9 total electrons
1s
2s
3s3p
2p
NeonNeon
Group 8A
Atomic number = 10
1s2 2s2 2p6 --->
10 total electrons
1s
2s
3s3p
2p
Note that we Note that we have reached the have reached the end of the 2nd end of the 2nd period, and the period, and the 2nd shell is full!2nd shell is full!
Exceptions to the Filling Order Rule (Cr, Cu)—these will not be on test!
Valence electrons
Importance and definition:
Definition: Electrons in the outermost energy levels; they determine the chemical properties of an element.
Write the noble gas configuration...the valence electrons are the ones beyond the core.
Example: Sulfur
Valence Electrons and Core Configuration (Shorthand)
What is the shorthand notation for S?
EOS
Sulfur has six valence electrons
Configurations of Ions
Cations: Formed when metals lose e– in highest principal energy level.
Example:
(Z = 11) Na
EOS
(Z = 11) Na+
Configurations of Ions
Anions: Formed when non-metals gain e– to complete the p sublevel.
Example:
EOS
-Z= 18 Cl-
Transition Metals
Transition metals (and p block metals) lose e– from the highest principal energy level
(n) FIRST, then lose their d electrons!
EOS
Zr: [Kr] 5s24d2
Zr+2 : [Kr] 4d2
Isoelectronic Species
Definition: Ions or atoms that have the same number of electrons
Example: Neon, O2-, F-, Na+, Mg2+, Al3+
all have the same configuration (1s22s22p6) and are isoelectronic
Electron Spin and Magnetism
•DiamagneticDiamagnetic: NOT : NOT
attracted to a magnetic attracted to a magnetic
fieldfield
•ParamagneticParamagnetic: :
substance is attracted to substance is attracted to
a magnetic field. a magnetic field. •Substances with Substances with
unpaired electronsunpaired electrons are are
paramagnetic.paramagnetic.
Examples
Mg
Cl
Write orbital notation: if it has an unpaired e- it is paramagnetic.
Periodic Properties & Trends
Electronegativity Ability of an atom to pull e- towards itself
Increases going up and to the right Across a period more protons in nucleus = more
positive charge to pull electrons closer Down a group more electrons to hold onto =
element can’t pull e- as closely
Periodic Properties & Trends
Electronegativity Ability of an atom to pull e- towards itself Across a period more protons in nucleus = more
positive charge to pull electrons closer Down a group more electrons to hold onto = protons
in nucleus can’t pull e- as closely
Atomic Radius
Definition:
½ experimental distance between centers of two bonded atoms
Atomic Radius
Trend in a family:
Size increases
down a group.
(More principal
energy levels)
Atomic Radius
Trend in a period:
Size decreases across a period, e- more strongly attracted to nucleus.
Atomic Radius
Transition metals:
Size stays relatively constant across a period; e-
added to inner energy level.
Memory DeviceMemory Device
LLLL: Lower Left, Larger AtomsLLLL: Lower Left, Larger Atoms
Sizes of IonsSizes of Ions
CATIONS are SMALLER than the atoms from which they are formed.
Size decreases due to increasing he electron/proton attraction.
Li,152 pm3e and 3p
Li+, 78 pm2e and 3 p
+
Sizes of IonsSizes of Ions
ANIONS are LARGER than the atoms from which they are formed.
Size increases due to more electrons in shell.
F, 71 pm9e and 9p
F-, 133 pm10 e and 9 p
-
Trends in Ion SizesTrends in Ion Sizes
Active Figure 8.15Active Figure 8.15
Trends in ion sizes are the same Trends in ion sizes are the same as atom sizes.as atom sizes.
First Ionization EnergyFirst Ionization Energy
Definition: energy required to remove an electron from an atom in the gas phase.
Mg (g) + 738 kJ ---> MgMg (g) + 738 kJ ---> Mg++ (g) + e- (g) + e-
First Ionization Energies
EOS
Trend in a group:
Decreases going down a group (e- further away; easier to remove)
Trend in a period:
Increases going across a period (e- held more tightly).
Memory DeviceMemory Device
LLLL: Lower Left, Larger Atoms;Looser
electrons
LLLL: Lower Left, Larger Atoms;Looser
electrons
Mg (g) + 738 kJ ---> Mg+ (g) + e-
MgMg+ + (g) + 1451 kJ ---> Mg(g) + 1451 kJ ---> Mg2+2+ (g) + e- (g) + e-
Definition: energy required to remove 2nd electron from an atom in the gas phase. Takes more energy because e- is removed from increasingly positive ion.
Second Ionization EnergySecond Ionization Energy
Electron Affinity
Some elements GAIN electrons to form anions.
Electron affinity is the energy involved when an atom gains an electron to form an anion.
A(g) + e- ---> A-(g) E.A. = ∆E
Trends in Electron Affinity
Trend in a group:
Affinity for e- decreases
going down a group
Trend in a series or period:
Affinity for e- increases going across a period
Electron Affinity
Note that the trend for E.A.
is the SAME as for I.E.!
Trends in Metallic Properties
Most metallic means easiest loss of electrons!Metals are on left, nonmetals on right of p.t.
A Summary of Periodic Trends
Remember LLLL!!