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)