unit 2: chapters 6, 7, 2.6, 8 · notice that the speed of light unit and the wavelength unit does...
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ELECTRONIC STRUCTURE OF ATOMS(I.E., QUANTUM MECHANICS)
Unit 2: chapters 6, 7, 2.6, 8.3
1
UNIT 2: REVIEW
Chapters: Chapters 6.1-6.3
2
WAVE-PARTICLE DUALITY
JJ 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!
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.
c = c = speed of light, a constant (3.00 x 108 m/s)
= frequency, in units of hertz (hz, sec-1) = wavelength, in meters
sometimes waves are measure in nanometers: 1 x109 nm = 1 m
ELECTROMAGNETIC RADIATION PROPAGATES THROUGH SPACE AS A WAVE MOVING AT THE SPEED OF LIGHT.
E = h
E = 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.
Long Wavelength
=Low Frequency
=Low ENERGY
Short Wavelength
=High
Frequency=
High ENERGY
8
How much energy, in kJ, does a wave have, if it has a wavelength of 4.0 x 103 nm?
Find the equation needed on the data equation sheet:
c =
E = h
Watch your units, make sure they match
Use the first equation to find the frequency of the wave(speed of light can be found on the data equation sheet)
3.0 x108 m/s = (4.0 x 103nm) x frequency notice that the speed of light unit and the wavelength unit does not match so convert the wavelength to meter
and then divide speed of light by wavelength to get the frequency
4.0 x 103 nm x 1 m = 4.0 x 10-6 m 3.0x108 / 4.0 x10-6 = 7.5 x 1013 s-1 this is the frequency1x109 nm
Use the 2nd equation to find the energy of the wave (planks constant can be found on the data equation sheet)
E = 6.626 x10-34 Js x 7.5 x 1013 s-1 = 5.0 x 10-20 J Notice the answer is in Joules because of planks constant so convert to kJ 5.0 x 10 -20 J x 1 kJ = 5.0 x 10-23 kJ 1000J
EXAMPLE PROBLEM
(ultra red) ROY G BIV (ultra violet)
Large wavelengths to short wavelengths
ELECTROMAGNETIC SPECTRUM
10
LIGHT IS A PARTICLE (QUANTUM THEORY)
ch h E
Max Planck
(1858-1947)
• Planck:
Energy can be released or absorbed in packets or
fixed amounts of a standard size he called quanta
(singular: quantum). (photon)
Joule (J) is used to express energy.
Particle theory can be compared to a piano where the wave
theory is compared to the violin.
High frequency waves have high energy.
Planck’s constant (h) = 6.63 x 10-34 J-s
Electrons orbit the nucleus in orbits, like a solar system.
NIELS BOHR’S ATOM
Planetary
Model
Electrons cannot
exist between orbits
(energy is quantized)
Electrons closest to the nucleus are lowest in energy.
Ground state- electrons are in the lowest energy level possible
If energy is put into the atom, the electrons will jump up in energy- move away from the nucleus (excited state).
BOHR’S ATOM
Excited electrons naturally go back to ground state. In order to do this, energy must leave the atom. Because energy is quantized in an atom, the amount of energy that leaves is the difference in energy between orbits If this energy is in the visible light range, we will see certain colors (line emission spectrums)
BOHR’S ATOM
This produces bandsof light with definite wavelengths.
ELECTRON TRANSITIONSINVOLVE JUMPS OF DEFINITE AMOUNTS OFENERGY.
…produces a “bright line” spectrum
SPECTROSCOPIC ANALYSIS OF THE HYDROGEN SPECTRUM…
17
Atomic emission spectra:
Most sources produce light that contains many wavelengths at once.
However, light emitted from pure substances may contain only a few specific wavelengths of light called a line spectrum (as opposed to a continuous spectrum).
Atomic emission spectra are inverses of atomic absorption spectra.
BOHR’S MODEL OF THE H ATOM (AND ONLY H!)
Hydrogen: contains 1 red, 1 blue and 1 violet.
Carbon:
THIS IS A COMPARISON OF A CONTINUOUS SPECTRUM AND A LINE SPECTRUM
CALCULATING ENERGY CHANGE, E, FOR ELECTRON TRANSITIONS
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
UNIT 2.1
Chapter 6.4-6.9
Crash course: chapter 3
20
21
1. Green light has a wavelength of 550 nm. The energy of a photon of green light is:A) 3.64 x 10-38 J B) 2.17 x 105 J C) 3.61 x 10-19
D) 1.09 x 10-27 J E) 5.45 x 1012 J
2. In Bohr’s atomic theory, when an atom moves from one energy level to another energy level more distant from the nucleus:
A) energy is emitted B) energy is absorbedC) no change in energy occurs D) light is emitted
3. You dissolve 0.4500 g of impure potassium chloride, KCl, in water and add an excess of silver nitrate, AgNO3. You get 0.8402 g of insoluble silver chloride, AgCl. Calculate the percent by mass of KCl in the original sample.
PRACTICE
HEISENBERG UNCERTAINTY PRINCIPLE
The more certain you are about where the electron is, the less certain you can be about where it is going.
The more certain you are about where the electron is going, the less certain you can be about where it is.
“One cannot simultaneously determine both the position and momentum of an electron.”
WernerHeisenberg
QUANTUM MECHANICALMODEL OF THE ATOM
Mathematical laws can identify the regions outside of the nucleus where electrons are most likely to be found.
These laws are beyond the scope of this class…
Differs from Bohr’s model:
The kinetic energy of an electron is inversely related to the volume of the region to which it is confined.
It is impossible to specify the precise position of an electron in an atom at a given instant.
Erwin Schrödinger (1887 – 1961), an Austrian physicist, made major contributions. Ψ (psi) is known as the wave function.
QUANTUM MECHANICAL MODEL
25
Schrödinger’s wave function:
Relates probability (Y2) of predicting position of e- to its energy.
dt
dihU
dx
d
m
hE
YY
Y
2
22
2
Where: U = potential energy
x = position t = time
m = mass i =√(-1)
Erwin
Schrödinger
(1887 – 1961)
26
ELECTRON DENSITY DISTRIBUTION IN H ATOM
4 quantum numbers (like an address)
n, l, ml , ms
Quantum number
1. n= principle Energy level (period)
2. l = sublevel orbital
3. ml = orientation of orbital
4. ms= spin = ½, -½
QUANTUM NUMBERS, ENERGY LEVELS AND ORBITALS
ELECTRON ENERGY LEVEL (SHELL)
Generally symbolized by n, it denotes the probable distance of the electron from the nucleus. “n” is also known as the Principle Quantum number
Number of electrons that can fit in a shell: 2n2
Orbital shapes are defined as the surface that contains 90% of the total electron probability.
AN ORBITAL IS A REGION WITHIN AN ENERGY LEVEL WHERE THERE IS A PROBABILITY OF FINDING AN ELECTRON.
Electron Orbitals
The angular momentum quantum number,
generally symbolized by l, denotes the orbital
(subshell) in which the electron is located.
30
s orbital
p orbitals
REPRESENTATIONS OF ORBITALS
d orbitals
32
F ORBITALS
33
EnergyLevel (n)
Sublevels inmain energy
level (n sublevels)
Number oforbitals per
sublevel
Number ofElectrons
per sublevel
Number ofelectrons permain energylevel (2n2)
1 s 1 2 2
2 sp
13
26
8
3 spd
135
2610
18
4 spdf
1357
261014
32
ENERGY LEVELS, SUBLEVELS, ELECTRONS
ELECTRON SPIN
The Spin Quantum Number describes the behavior (direction of spin) of an electron within a magnetic field.
Possibilities for electron spin:
1
2
1
2
A maximum of 2 e- can be in an orbital
Number of shapes in a shell = n
Number of orbitals in a shell = n2
Number of electrons in a shell = 2n2
s has 1 orbital, p has 3, d has 5, f has 7
s can hold 2 e-, p:6, d: 10, f:14
37
FILLING ORDER OF ORBITALS1. Aufbau principle: e- enter orbitals of lowest
energy first (* postulated by Bohr, 1920)
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
6d
4f x 7
5f x 77p
• Relative stability & average distance of e- from nucleus
38
1. Aufbau principle: e- enter orbitals of lowest
energy first
1s
2s
3s
4s
5s
6s
7s
3d
4d
5d
6d
4f x 7
5f x 7
2p
3p
4p
5p
6p
7p
• Relative stability & average distance of e- from nucleus
Filling Order of Orbitals
39
2. Pauli exclusion principle (1925): no two e- can
have the same four quantum numbers; e- in same
orbital have opposite spins (up and down)
3. Hund’s rule: e- are added singly to each equivalent
(degenerate) orbital before pairing
Ex: Phosphorus (15 e-) has unpaired e- in
the valence (outer) shell.
1s 2s 2p 3s 3p
Wolfgang
Pauli
(1900 – 1958)
Friedrich
Hund
(1896 - 1997)
40
Paramagnetic: a substance that is drawn to a
magnetic field (contains one or more unpaired
electrons)
Diamagnetic: substance that is repelled by a
magnetic field (contains no unpaired electrons)
6.9: PERIODIC TABLE & ELECTRONIC CONFIGURATIONS
s block p blockd blockf block
s1 s2
p1p2p3p4p5 p6
d2d3d5d5d6d7d8d10d10
f1 f2 f3 f4 f5 f6 f7 f8 f9f10f11f12f13f14
s2
1s2s3s4s5s6s7s
2p3p4p5p6p7p
4f5f
3d4d5d6d
3d4d5d6d
d1
42
Element Standard ConfigurationNoble Gas
Shorthand
Nitrogen
Scandium
Gallium
ELECTRONIC CONFIGURATIONS
[He] 2s22p3
[Ar] 4s23d1
[Ar] 4s23d104p1
1s22s22p3
1s22s22p63s23p64s23d1
1s22s22p63s23p64s23d104p1
Valence electrons are the electrons in the outer s and p shell that are on the same period or row.
Examples:
Nitrogen: 1s22s22p3
5 valence electrons
1s22s22p63s23p64s23d104p1
3 valence electrons
Group 1 has 1, group 2 has 2, group13 has 3, etc.
VALENCE ELECTRONS
PERIODIC PROPERTIES OF THE ELEMENTS
Unit 2.2
Chapters 7.1-7.5, 8.3
Crash course: chapter 4
45
1. Which of the following frequencies correspond to light with the longest wavelength?
A) 3.00 x 1013s-1 C) 9.12 x 1012 s-1
B) 4.12 x 105 s-1 D) 3.20 x 109 s-1 E) 8.50 x 1020 s-1
2. The lines in the emission spectrum of hydrogen result from
A) electrons given off by hydrogen as it cools B) decomposing hydrogen atoms C) electrons given off by hydrogen when it burns D) energy given off in the form of visible light when an electron moves from a
higher energy state to a lower energy state 3. The electron configuration of a ground-state Cd atom is
A)[Kr]5s14d10 B) [Kr]5s24d10 C) [Ar]4s24d10
D) [Kr]5s23d10 E) [Ar]4s14d10
5) Calculate the smallest increment of energy that can be emitted or absorbed at a wavelength of 548 nm.
PRACTICE
46
Coulombic attraction is the attraction between oppositely charged particles. (electrons attracted to protons)
Coulombic attractions become stronger when
Charges increase
Distance decreases
COULOMBIC ATTRACTIONS
In an atom each electron is simultaneously attracted to the nucleus and repelled by other electrons.
Any electron between the nucleus and the electron of interest will reduce the nuclear charge acting on the electron.
The net positive charge attracting the electron is called the effective nuclear charge.
The positive charge experienced by a valence electron is always less than the full nuclear charge because the core electrons shield the outer electrons.
EFFECTIVE NUCLEAR CHARGE
PERIODIC TRENDS
Definition: Half of the distance between
nuclei in covalently bonded diatomic
molecule
❖Radius decreases across a period
❖ Increased effective nuclear charge due to
more protons
❖Radius increases down a group
❖ Each row on the periodic table adds a
“shell” or energy level to the atom
ATOMIC RADIUS
TABLE OF ATOMIC RADII
PERIOD TREND:ATOMIC RADIUS
Place the following in order of increasing size:
Na, Be, Mg
Answer: Be < Mg < Na
EXAMPLES
Ionization Energy: the energy required to remove an
electron from an atom
Increases for successive electrons taken from the same atom
Tends to increase across a period
Electrons in the same energy 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 the nucleus
Table of 1st Ionization
Energies
PERIODIC TREND:IONIZATION ENERGY
the first ionization energy (I1)is the amount of energy it takes to remove the first electron from an atom
The second ionization energy (I2) is the amount of energy is takes to remove a second electron from an ion, etc.
I1 < I2 <I3 and so forth (once an electron is removed it gets harder to take another one away-think of effective nuclear charge)
IONIZATION ENERGY
I1 I2 I3 I4 I5 I6 I7
Na 496 4560
Mg 738 1450 7730
Al 578 1820 2750 11600
Si 786 1580 3230 4360 16100
P 1012 1900 2910 4960 6270 22200
S 1000 2250 3360 4560 7010 8500 27100
Cl 1251 2300 3820 5160 6540 9460 11000
Ar 1521 2670 3930 5770 7240 8780 12000
IONIZATION ENERGY (KJ/MOL)
A sharp increase in ionization energy occurs when a core
electron is removed. The large jump occurs because the core
electron is much closer to the nucleus and experiences a
much greater effective nuclear charge than valence electrons.
ELECTRONEGATIVITY
Definition: A measure of the ability of an atom in a chemical compound to attract electrons
o Electronegativity tends to increase across a period
o As radius decreases, electrons get closer to the bonding atom’s nucleus
o Electronegativity tends to decrease down a group or remain the same
o As radius increases, electrons are farther from the bonding atom’s nucleus
PERIODIC TABLE OF ELECTRONEGATIVITIES
PERIODIC TREND:ELECTRONEGATIVITY
Metals- on the left of the periodic table
Shiny luster, malleable, ductile, good conductors of heat and electricity, tend to form cations (low ionization energy)
Increasing metallic character as we go left and down on the periodic table
Nonmetals
Brittle solids, poor conductors, tend to form anions (high electron affinity)
Metalloids
semiconductors
METALS, NONMETALS, AND METALLOIDS
When an atom gains or loses an electron, it is called an ion.
Cation if positively charged or giving electrons away. Electrons leave the highest shell first.
Anion if negatively charged or accepting electrons.
IONS
Metals tend to lose their valence electrons when bonding and become cations
Nonmetals tend to gain electrons in their highest energy level when bonding and become anions
The most stable atoms have 8 electrons in their highest energy level
IONS
GENERAL CHARGES
+1 -3Many
different
charges
+2
N
O
B
E
L
G
A
S
E
S
-1-2+4
-4+3
Many Different Charges
IONIC RADII
Cations
Positively charged ions formed whenan atom of a metal loses one or more electrons
Smaller than the corresponding atom
Anions
Negatively charged ions formed when nonmetallic atoms gain one or more electrons
Larger than the corresponding atom
TABLE OF ION SIZES
An isoelectric series are ions that possess the same number of electrons
Example: Na +, Mg 2+ , Al 3+
In an isoelectric series, the ion with the most protons (highest effective nuclear charge) will have the smallest radius.
ISOELECTRIC SERIES
UNIT 2.3
Chapter 8.1, 8.2
69
1. Answer the questions below related to properties of the alkali metal potassium.
a) The atomic radius of potassium is greater than the atomic radius of zinc. EXPLAIN.
b) The second ionization energy of potassium is greater than the second ionization energy of calcium. EXPLAIN.
c) the electronegativity for Oxygen is greater than potassium. EXPLAIN.
PRACTICE
71
Atoms or ions that are strongly attached to one another
Chemical bonds will form if potential energy decreases (becomes more stable)
CHEMICAL BONDS
72
1. Ionic: electrostatic attraction between oppositely charged ions (typically between a metal and a nonmetal)
2. Covalent: sharing of e- between two atoms (typically between nonmetals)
molecules created
3. Metallic: “sea of e-”; bonding e- are relatively free to move throughout the 3D structure
4. Covalent Network: atoms bond with strong directional covalent bonding that lead to giant molecules and networks. Examples: carbon and silicon: diamond,
graphite, silicon oxide (sand)
8.1: TYPES OF BONDS
73
Valence e-:
e- in highest energy level and involved in bonding; all elements within a group on periodic table have same # of valence e-
Valence electrons are the outer s and p
Lewis symbol (or electron-dot symbol):
Shows a dot only for valence e- of an atom or ion.
Place dots at top, bottom, right, and left sides and in pairs only when necessary (Hund’s rule).
Primarily used for main group elements only
Ex: Draw the Lewis symbols of C and N.
LEWIS SYMBOLS
•
• C ••
•
: N ••
Gilbert N.
Lewis
(1875 – 1946)
74
Atoms tend to gain, lose, or share e- until they are surrounded by 8 valence e- (have filled s and p subshells) and are thus energetically stable.
Exceptions do occur (and will be discussed later.)
THE OCTET RULE
75
Metallic elements have low I.E.; this means valence e- are held “loosely”.
A metallic bond forms between metal atoms because of the movement of valence e- from atom to atom to atom in a “sea of electrons”. The metal thus consists of cations held together by negatively-charged e- "glue.“
METALLIC BONDING
This results in excellent thermal
& electrical conductivity,
ductility, and malleability.
A combination of 2 metals is
called an alloy.
Free e- move rapidly in response to electric fields, thus metals are excellent conductors of electricity.
Free e- transmit kinetic energy rapidly, thus metals are excellent conductors of heat.
Layers of metal atoms are difficult to pull apart because of the movement of valence e-, so metals are durable.
However, individual atoms are held loosely to other atoms, so atoms slip easily past one another, so metals are ductile.
77
An alloy is best defined as a substance that contains a mixture of elements and has metallic properties. There are two types of alloys:
Substitutional alloy- some of the host metal atoms are replaced by other metal atoms of similar size.
Interstitial alloy- is formed when some of the holes in the closest packed lattice are occupied by smaller atoms (changes the properties of the host metal)
METAL ALLOYS
78
Ionic bonds do not form moleculesAn ionic formula is an empirical formula (smallest whole number ratio of atoms) and doesn’t show what the structure looks like
IONIC BONDING
79
•Results as atoms lose or gain e- to achieve a noble gas e-
configuration; is typically exothermic.
–The bonded state is lower in energy (and therefore more stable).
–Electrostatic attraction results from the opposite charges.
•Occurs when diff. of EN of atoms is > 1.7 (maximum is 3.3: CsF)
•Can lead to interesting crystal structures (Ch. 11)
–Ionic compounds are brittle solids with high melting points. Solids do not conduct electricity, but molten form will conduct (ions freely moving)
IONIC BONDING
80
IONIC BONDINGFormulas for ionic compounds (metal with a nonmetal)
are ALWAYS empirical (lowest whole number ratio)
Ionic compounds to not form molecules (they form crystals) so the formula doesn’t show the exact number of atoms in the compound but instead a ratio of how they bond.
Compounds are always neutral so when writing an ionic formula
Make sure the charges add up to zero.
Example: Mg2+ and Cl- so the formula is MgCl2
81
Tells you the amount of energy it takes to break an ionic bond (If the lattice energy is negative its showing the amount of energy released when the ionic bond formed)
Larger lattice energy means stronger ionic bond
LATTICE ENERGY
82
Hlattice = energy required to completely separate 1 mole of solid ionic compound into its gaseous ions
LATTICE ENERGY
rr
QQH lattice
Electrostatic attraction (and thus lattice energy)
increases as ionic charges increase and as ionic radii
decrease.
Ex: Which has a greater lattice energy?
NaCl or KCl NaCl or MgS
83
Arrange the following in order of increasing lattice energy: NaF, CsI, and CaO
Answer: CsI < NaF < CaO
84
Transition metals typically form +1, +2, and +3 ions.
It is observed that transition metal atoms first lose both “s” e-, even though it is a higher energy subshell.
Most lose e- to end up with a filled or a half-filled subshell.
UNIT 2.4 Chapter 8.4-8.8
Crash course: chapter 10
85
86
1. Which of the following would have the highest melting point, KBr, CaO, MgO. Explain.
3. In which of the following processes are covalent bonds broken?
a. C10H8 (s) C10H8(l)
b. C(diamond) C(graphite)c. NaCl (s) NaCl (molten)d. KCl (s) KCl (aq)e. NH4NO3 (s) NH4
+ + NO3-
4. The energy from radiation can be used to cause the rupture of chemical bonds. A minimum energy of 822 kJ mol- is required to break a covalent bond. What is the longest wavelength of radiation that possesses the necessary energy to break one covalent bond?
CLASS STARTER
87
Covalent bonds form molecules
The formula is not always empirical but shows what the molecule looks like
A molecular formula shows what the molecule actually looks like
Molecular formula: C6H6 empirical: CH
COVALENT BONDS
These molecules are shown in ball and stick form.
These are also represented in structural formulas like this:
H – O H – N – H H
| | |
H H H – C – H
|
H
MOLECULES- 2 OR MORE ATOMS BOUND TOGETHER THAT ACT AS A
SINGLE, DISTINCT OBJECT
Ball and Stick Structural Condensed
CONDENSED STRUCTURAL FORMULA
90
Atoms share e- to achieve noble gas configuration that is lower in energy (and therefore more stable).
Polar covalent: (different elements)
e- pulled closer to more EN atom and are shared unequally
-Nonpolar covalent: (same elements)
e- shared equally
COVALENT BONDING
91
•H2 nonpolar; the hydrogens share the electrons equally
•HF polar: fluorine pulls the electrons closer so they share the electrons unequally
•In a polar molecule, one end is partially positive and one is partially negative (Dipole)
s+
s-
H—F or H—F (vector points to neg. end)
92
• a line between atoms shows that 2 electrons are being shared H—F (single bond)
•Multiple bonds
–A double line shows that 4 electrons are being shared
•O=O (double bond)
–A triple line shows that 6 electrons are being shared
•N=N (triple bond)
Bond length : triple < double < single
Bond energy : triple>double>single
COVALENT BONDS
93
An indication of bond strength and bond length
Single bond: 1 pair of e- shared
Ex: F2
BOND ORDER
•• ••
:F-F:•• ••
O=O
:N ≡ N:
Longest,
weakest
Shortest,
strongest
Double bond: 2 pairs of e- shared
Ex: O2
Triple bond: 3 pairs of e- shared
Ex: N2
94
1. Add up valence e- from all atoms in formula.– If there is a charge, add e- (if an anion) or subtract e- (if a cation).
2. Draw the “molecular skeleton”:– Place the least EN atom(s) in the center. (never H)– Array the remaining elements around the center and connect them with a
single bond. (When in doubt, put the element written first in the formula in the center of the molecule.)
3. Complete the octets of the outer (more EN) atoms first.
4. Place leftover e- on the central atom, even if it violates the octet rule.
5. If the central atom does not have an octet, create multiple bonds by sharing e- with the outer atoms.
DRAWING LEWIS STRUCTURES
95
1.CO2 has 16 valence electrons
2. CO2 wants 24 total electrons
3. 24-16 = 8
4. 8/2=4 bonds
5. O=C=O
6. :O=C=O:
1. add up valence e- from all the elements in the formula
2. Add up the amount of e- each atom in the formula wants (all atoms want 8 except H=2, Be=4, B=6)
3. Subtract #1 from #2 this tells you the number of e- shared
4. Divide by two to find the number of bonds in the Lewis structure
5. Draw the “molecular skeleton” with correct number of bonds.
remember H can only single bond.
6. complete the octet on each atom
LEWIS STRUCTURES
96
FORMAL CHARGES
Formal charge is the charge calculated for an atom in a
Lewis structure on the basis of an equal sharing of bonded
Electron pairs.
Formal charges help us decide which Lewis structure is best
NITRIC ACID
We will calculate the formal charge for each atom in this Lewis structure.
.. :
..H O
O
O
N
:
:..
..
Formal charge of H
NITRIC ACID
Hydrogen shares 2 electrons with oxygen.
Assign 1 electron to H and 1 to O.
A neutral hydrogen atom has 1 electron.
Therefore, the formal charge of H in nitric acid is 0.
.. :
..H O
O
O
N
:
:..
..
Formal charge of H
:
..H O
O
O
N
:
:..
..
NITRIC ACID
Oxygen has 4 electrons in covalent bonds.
Assign 2 of these 4 electrons to O.
Oxygen has 2 unshared pairs. Assign all 4 of these electrons to O.
Therefore, the total number of electrons assigned to O is 2 + 4 = 6.
.. :
..H O
O
O
N
:
:..
..
Formal charge of O
NITRIC ACID
Electron count of O is 6.
A neutral oxygen has 6 electrons.
Therefore, the formal charge of O is 0.
.. :
..H O
O
O
N
:
:..
..
Formal charge of O
NITRIC ACID
Electron count of O is 6 (4 electrons from unshared pairs + half of 4 bonded electrons).
A neutral oxygen has 6 electrons.
Therefore, the formal charge of O is 0.
.. :
..H O
O
O
N
:
:..
..
Formal charge of O
NITRIC ACID
Electron count of O is 7 (6 electrons from unshared pairs + half of 2 bonded electrons).
A neutral oxygen has 6 electrons.
Therefore, the formal charge of O is -1.
.. :
..H O
O
O
N
:
:..
..
Formal charge of O
NITRIC ACID
Electron count of N is 4 (half of 8 electrons in covalent bonds).
A neutral nitrogen has 5 electrons.
Therefore, the formal charge of N is +1.
.. :
..H O
O
O
N
:
:..
..
Formal charge of N
Formal Charge
Formal charge =
Valence
electrons
number of bonds
divided by 2
number of
unshared
electrons
– –
An arithmetic formula for calculating formal charge.
The most stable Lewis structure will be that in which
-the atoms bear the smallest formal charges
-any negative formal charge reside on the more
electronegative atom
105
SO42- HCN
C2H4 NO31-
EX: DRAW THE LEWIS STRUCTURE AND FIND THEFORMAL CHARGE OF EACH ATOM IN THE STRUCTURE
H
H
H
H
H
H
H
H
H
H
H
H
Resonance is invoked when more than one valid Lewis structure can be written for a particular molecule.
RESONANCE
The actual structure is an average of the resonance
structures.
Benzene, C6H6
The bond lengths in the ring are identical, and
between those of single and double bonds.
H
H
H
H
H
H
H
H
H
H
H
H
Resonance bonds are shorter and stronger than single bonds.
RESONANCE BOND LENGTH AND BOND ENERGY
Resonance bonds are longer and weaker than double
bonds.
O O O
O O O
RESONANCE IN OZONE, O3
Neither structure is correct.
Oxygen bond lengths are identical, and intermediate to
single and double bonds
Resonance in a carbonate ion:
Resonance in an acetate ion:
RESONANCE IN POLYATOMIC IONS
THE LOCALIZED ELECTRON MODEL
Lewis structures are an application of the “Localized
Electron Model”
L.E.M. says: Electron pairs can be thought of as
“belonging” to pairs of atoms when bonding
Resonance points out a weakness in the Localized
Electron Model.
Models are attempts to explain how nature operates on the microscopic level
based on experiences in the macroscopic world.
MODELS
Models can be physical as
with this DNA model
Models can be mathematical
Models can be theoretical or
philosophical
❖A model does not equal reality.
❖Models are oversimplifications, and are therefore often wrong.
❖We must understand the underlying assumptions in a model so that we don’t misuse it.
FUNDAMENTAL PROPERTIES OF MODELS
113
Bond order: single bond = 1, double bond=2, triple bond = 3
To determine bond order with resonance structures:
Pick one bond and add up the integer bond order in one resonance structure to the same bond position in all other resonance structures.
Divide the sum by the number of resonance structures to find bond order.
BOND ORDER & RESONANCE STRUCTURES
114
Which has shorter bonds? What is the bond order in each?
SO3 or SO32-
Answer: SO3
Bond order for SO3 is 1 1/3
bond order of SO32- is 1
115
•Odd-electron molecules:Ex: NO or NO2 (involved in breaking down ozone in the upper atmosphere)
•Incomplete octet:
H2 He BeF2 BF3
EXCEPTIONS TO THE OCTET RULE
116
Expanded octet: occurs in molecules when the central atom is in or beyond the third period, because the empty 3d subshell is used in the bonding
PCl5 SF6
If you find the number of bonds mathematically, the math won’t make sense and you’ll know it has an expanded octet.
Only use single bonds and add extra electrons to the central atom. (outside atoms are usually halogens)