2013 General Chemistry I * The coverage of the Quiz (Q) * The coverage of the Study Summary (SS) WeekPeriodChapterChemical Principles / Atkins, JonesNote

2013 General Chemistry I* The coverage of the Quiz (Q)* The coverage of the Study Summary (SS)

Week Period Chapter Chemical Principles / Atkins, Jones Note

1st 3/4-3/7

2nd 3/11-3/14

3rd 3/18-3/21

4th 3/25-3/28 2.1-2.16Ionic Bonds, Lewis Structures and Covalent

Bonds SS: Chap 2 (3/25)

5th 4/1-4/4 3.1-3.7 VSEPR and Valence-Bond Theory Q: Chap 1 (4/1)

6th 4/8-4/11 3.8-3.12 Molecular Orbital Theory SS: Chap 3 (4/8)

7th 4/15-4/18 Q: Chap 2 (4/15)

8th 4/22-4/25Mid-term

Exam(Chapters 1-3)

9th 4/29-5/2 4.1-4.11 The Gas Laws, Molecular Motion

10th 5/6-5/9 4.12-5.6 Real Gases, Intermolecular Forces SS: Chap 4 (5/6)

11th 5/13-5/16 5.7-5.16 Liquids and Solids SS: Chap 5 (5/13)

12th 5/20-5/23 7.1-7.12 EnthalpyQ: Chap 4 and 5

(5/20)

13th 5/27-5/30 7.13-8.8 Chemical Change and EntropySS: Chap 7 and 8

(5/27)

14th 6/3-6/5 8.9-8.16 Gibbs Free Energy Q: Chap 7 (6/3)

15th 6/10-6/13 Additional Class / Review

16th 6/17-6/20 Final Exam (Chapters 4, 5, 7, 8)

2013 General Chemistry I

Chapter 2.CHEMICAL BONDS


IONIC BONDS

COVALENT BONDS

2.1 The Ions That Elements Form2.2 Lewis Symbols2.3 The Energetics of Ionic Bond Formation2.4 Interactions Between Ions

2.5 Lewis Structures2.6 Lewis Structures of Polyatomic Species2.7 Resonance2.8 Formal Charge


Understanding the bond formation between atoms ⇒⇒⇒understanding properties and reactivity of materials

Designing new materials

Chapter 2. CHEMICAL BONDS

(화학결합 )Lowering of energy by rearranging valence electrons

Ionic bond( 이온결합 ); electron transfer + electrostatic attraction, NaClCovalent bond( 공유결합 ); sharing electrons, NH3

Metallic bond( 금속결합 ); cations held by a sea of electrons, copper

Lewis structure ( 루이스구조 )

Octet rule ( 옥텟 규칙 ) ---- ---- coordinate covalent bond ( 배위결합 )

Resonance ( 공명 )

Formal charge ( 형식전하 )

Oxidation number ( 산화수 )


Key Ideas Bond formation by lowering of energy

Goal

Understanding the bond formation between atoms ⇒⇒⇒ Designing new materials

Chapter 2. CHEMICAL BONDS Chemical bond is the link between atoms.

Lowering of energy by rearranging valence electrons

Ionic bond; electron transfer + electrostatic attraction, NaCl

Covalent bond; sharing electrons, NH3

Metallic bond; cations held by a sea of electrons, copper

How are we going to apply Q.M. knowledge ?


IONIC BONDS (Sections 2.1-2.4)55

ionic model: the description of bonding in terms of ions

ionic solid: three-dimensional crystalline solidan assembly of cations and anions stacked together in a regular pat-

tern

Metal + nonmetal


2.1 The Ions That Elements Form56

Cations: Remove outermost electrons in the order np, ns, (n-1)d

Metallic elements in the s block and on the left of the p block in Period 3 (Al)

⇒ lose electrons down to their noble-gas cores; 1s2 (duplet) or ns2np6 (octet)

Li+, Be2+, Na+, Mg2+, Al3+, ···

Metallic elements of the p block in Periods 4 and later

⇒ lose electrons down to their noble-gas cores surrounded by d10; (n – 1)d10ns2np6

The delectrons, in most cases, cannot be lost.

Ga3+([Ar]3d10), ···

Elements in the d block

⇒ ns-electrons first, then variable number of (n – 1)d-electrons

Fe2+([Ar]3d6), Fe3+([Ar]3d5), ···


56

Anions: Add electrons until the next noble-gas configuration is reached

Many elements in the p and d blocks; variable valence

⇒ Inert-pair effect; p-electrons alone or all their valence p- and s-electrons

The order of losing electrons:

In+, In3+, Sn2+, Sn4+, ···


2.2 Lewis Symbols58

- valence electrons – depicted as dots; a pair of dots for paired electrons

- cations and anions

(1916)


Formulation of ionic compounds

1) Cations by removing all dots from metal atoms:

2) Anions by adding dots to complete the valence shell (octet or duplet):

3) Adjust the number of each element to conserve the total number of dots.

4) Write the charge of each ion:

CaCl2; empirical formula


2.3 The Energetics of Ionic Bond Formation

NaCl ionic crystal formation; electron transfer from Na to Cl and then Coulomb stabilization

The energy required for the formation of ionic bonds is supplied largely by the attraction between oppositely charged ions.



2.4 Interactions Between Ions59

- In an ionic solid, each cation is attracted to all the anions to a greater or lesser extent. → a “global” characteristic of the entire crystal

i.e. ionic bond is not a bond between two ions !

Lattice energy: the difference in energy between the ions packed together in a solid and the ions widely separated as a gas

- strong electrostatic interactions in ionic solids → high melting points and brittleness


- Coulomb potential energy of the interactions of two individual ions

60

e is the fundamental charge; z1 and z2 are the charge numbers of thetwo ions; r12 is the distance between the centers of the ions; e0 is the vacuum permittivity.

2.4 Interactions Between Ions


61

- In a one-dimensional crystal in which cations and anions alternate along a line,

; A = 2 ln2 or 1.386

- molar potential energy of a three-dimensional crystal

The factor A is the Madelung constant, dependent on how the ions are arranged about one another

toward right hand side

toward left hand side


Short-range repulsions between ions

Total potential energy = EP + EP*

- real potential energy of an ionic solid→ attractive and repulsive interionic interactions in close range.


Born-Meyer equation; correction for repulsions to the Madelung constant

repulsive effect

The coulombic interaction between ions in a solid is largewhen the ions are small and highly charged.


COVALENT BONDSCOVALENT BONDS

"With brilliant insight, and before anyone knew about quantum mechanics or orbitals, Lewis proposed"

Chemical bonding in H2+ by sharing an electron between two protons

An electron between the two nuclei exerts an attractive force on the nuclei.

Quantum mechanical view of covalent bond


2.5 Lewis Structures

Covalent bonding; octet (or duplet) by sharing (Lewis, 1916)

Nonmetal atoms share electrons to complete their octet; lines (bonding pairs), dots (lone pairs)

- octet rule: atoms go as far as possible toward completing their octets

- A line (-) represents a shared pair of electrons : a bond.


2.6 Lewis Structures of Polyatomic Species 64

- Lewis structure does not portray the 3D shape of a molecule or ion, but simply displays which atoms are bonded together.

- bond order: the number of bonds that link a specific pair of atoms.

- Each atom completes its octet by sharing pairs of electrons.

Methane; CH4 (why not CH3 nor CH5?)


Writing a Lewis structure

- terminal atom: bonded to only one other atom central atom: bonded to at least two other atoms

65

- The element with the lowest ionization energy (less greedy) as the central atomelectronegativity is a better indicator

Ex. HCN

- Atoms symmetrically around the central atom; SOS for S2O (Exception; NNO)

- OH is attached to the central atom in oxoacids; HO–Cl for HClOi.e. H2SO4 ----- (HO)2SO2

- Polyatomic ions; total number of electrons should be adjusted to represent the overall charge


Toolbox 2.1: How to write the Lewis structure of a polyatomic species

Step 1: # of electron pairs = total # of valence electrons / 2

Step 2: Write down the most likely arrangements of atoms.

Step 3: One electron pair between each pair of bonded atoms

Step 4: Complete the octet (duplet for H) of each atom using the remaining electron pairs.

Step 5: A line for each bonding pair

Take care of charges.

Form multiple bonds if in short of electron pairs.



Ex 2.4 Lewis structures for CH3COOH (multiple central atoms)

# of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12

CH3COOH

Lewis structures for C2O2H4 (multiple central atoms)How many structures can you draw?


Lewis structures for NO3- # of electron pairs = (5 + 3 6 + 1) / 2 = 12

3 12

- double-headed arrows (↔), indicating a blend of the contributing structures

- delocalization: a shared electron pair is distributed over several pairs of atoms and cannot be identified with just one pair of atoms.

Identical N–O bond lengths of 124 pm (> 120 pm for N=O, < 140 pm for N–O)


69 Benzene, C6H6

- All the carbon-carbon bonds with the same length

- Only one 1,2-dichlorobenzen exists.


2.8 Formal Charge 70

Formal charge – the charge it would have if the bonding were perfectly covalent in the sense that the atom had exactly a half-share in the bonding electrons

V = the number of valence electrons in the free atomL = the number of electrons present on the bonded atom as lone pairsB = the number of bonding electrons on the atom

Formal charge indicates the degree of redistribution of electrons relative to free atomsnot the real charge of an atom

- A Lewis structure in which the formal charges of the individual atoms are closest to zero typically represents the lowest energy arrangement of the atoms and electrons.


OCO with lower formal charges is more likely for CO2 than COO.

NNO with lower formal charges is more likely for ON2 than NON.

- Formal charge exaggerates the covalent character of bonds by assuming that the electrons are shared equally.

- Oxidation number exaggerates the ionic character of bonds. It represents the atoms as ions, and all the electrons in a bond are assigned to the atom with the lower ionization energy (higher electronegativity (2.12)).

O C O0 0 0

O C O-2 +4 -2

formal charge oxidation state

HCN has lower formal charges than HNC.


Formal charges differ from oxidation numbers!

Neither of them is the true charge.

Quantum mechanically, there is no true localized charge on an atom!

Ca2+ is an oxidation state of calcium with the oxidation number of “+2”.

Oxidation number is important in following the oxidation-reduction reaction.

- octet rule: In covalent bond formation, atoms go as far as possible toward completing their octets by sharing electron pairs.

There are many exceptions to the octet rule


2.9 Radicals and Biradicals 72

Radicals: species with an unpaired electron, highly reactive

CH3 OH NO NO2

Biradicals: molecules with two unpaired electrons

2.10 Expanded Valence Shells

Odd number of electrons # of electron pairs = (5 + 2 x 6) / 2 = 8.5

O N O

Expanded valence shell (expanded octet); large atoms with empty d-or-bitals (Period 3 or later) may accommodate more than 8 electrons.

PCl5 vs. NCl5


72

- hypervalent compound: a compound that contains an atom with more atoms attached to it than is permitted by the octet rule

ST 2.10B Linear I3– ion

3x7 + 1 = 22 electrons, 11 electron pairs 2 bonds, 2 3 + 2 = 8 lone pairs

Remaining one pair into the central I


Lewis structures for C2O2H4 (multiple central atoms)

# of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12

3 3 3

3 3

2


72

5 6 + 2 = 32 valence electrons, 16 electron pairs

Ex 2.8 Dominant resonance Lewis structure of SO42–

- variable covalence: the ability to form different numbers of covalent bonds

most preferred structure

number of resonance structures


THE PERIODICITY OF ATOMIC PROPER-TIES

39


2.11 The Unusual Structures of Some Group 13/III Compounds 76

- incomplete octet: fewer than eight valence electrons

- boron and aluminum

- completing octets by a coordinate covalent bond, in which both electrons come from one of the atoms


2.11 The Unusual Structures of Some Group 13/III Compounds 76

Possible due to the atomic radius of Al in AlCl3 larger than that of B in BCl3.

borane


Chapter 2. CHEMICAL BONDS

Ionic bond; electron transfer + electrostatic attraction, NaCl

Covalent bond; sharing electrons, NH3

Formal charge – the charge it would have if the bonding were perfectly covalent in the sense that the atom had exactly a half-share in the bonding electrons

Oxidation number– exaggerates the ionic character of bonds. It represents the atoms as ions, and all the electrons in a bond are assigned to the atom

with the lower ionization energy (higher electronegativity ).

Nonpolar covalent bond; the average charge on each atom is zero.


76

- partial charges: the charges on the atoms

- polar covalent bond: a bond in which ionic contributions to the resonance result in par-tial charges

lower energystructure

- electric dipole: a partial positive charge next to an equal but opposite partial negative charge

size of an electric dipole ---- electric dipole moment (m)

Unit: Debye (D)

the average charge on each atom is not zero.

definition: a dipole between electron and protonseparated by 100 pm is 4.80D


Electronegativity (c) – Electron-pulling power of an atom when it is part of a molecule (by Linus Pauling in 1932)

77

2.12 Correcting the Covalent Model: Electronegativity

Cl-H bond: m = ~1.1 D : Cl has ~23% of an electron’s charge

Rough guide to the charge separation in a bond between two atoms

Average based on measured bond energies from a large range of compounds; can be revised

Measure of extra stability due to ionic contributions


77

- Mulliken scale: c = ½(I + Eea)

average of the ionization energy and electron affinity

2.12 Correcting the Covalent Model: Electronegativity

Mulliken’s electronegativity scale (1934); properties of an isolated atom

Exactly defined


77

- rough rules of thumb

ionicpolar covalent

covalent

i.e. NaCl or KF : ionic C-O : polar covalent Ca-O : ionic



2.13 Correcting the Ionic Model: Polarizability

78

- All ionic bonds have some covalent character.

- highly polarizable atoms and ions: readily undergo a large distortion of their electron cloud

i.e. large anions and atoms such as I-, Br-, and Cl-

- polarizing power: property of ions (and atoms) that cause large distortions of electron clouds

- increases with decreasing size and increasing charge of a cation

i.e. the small and/or highly charged cations Li+, Be2+, Mg2+, and Al3+


THE STRENGTHS AND LENGTHS OF COVALENT BONDS79

2.14 Bond Strength

Dissociation energy (D): energy required to separate the bonded atoms

- The bond breaking is homolytic, which means that each atom retains one of the electrons from the bond.

- average dissociation energy for one type of bond found in different molecules

i.e. C-H single bond:

average strength of bonds in a selectionof organic molecules, such as methane (CH4),ethane (C2H6), and ethene (C2H4)

measured by


2.15 Variation in Bond Strength 80

Strongest bond between two nonmetal atoms; CO (1062 kJ/mol)

Lone pair-lone pair repulsion due to the short F–F distance


80


ATP(aq) + H2O ADP(aq) + H2PO4–(aq) + 174 kJ/mol

bond stiffness ( bond strength); resistance to stretching and compressing

will be discussed in the Major Technique 1: Infrared (IR) Spectroscopy


2.16 Bond Length 82

Bond length: the distance between the centers of two atoms joined by a covalent bond

- corresponding to the internuclear distance at the potential energy minimum for the two atoms

- affecting the overall size and shape of a molecule evaluated by using spectroscopy or x-ray diffraction methods

- Factors influencing bond length


Box 2.2 Bond length measurements

quantum mechanical rotational energy with a rotational quantum number J

microwave spectroscopy

Reproducibility of bond lengths; constant within a few percent in similar arrangements


82 Covalent radius: contribution an atom to the length of a cova-

lent bond

Bond length: Approximately the sum of the covalent radii of the two atoms

- Decreases from left to right (increasing Zeff)

- Increases in going down a group (size of valence shells and better shielding by inner core electrons)


bond stiffness ( bond strength); resistance to stretching and compressing

INFRARED SPECTROSCOPY

Infrared radiation: electromagnetic radiation with longer wavelengths (lower frequencies) than red light ~ 1000 nm or ~ 3×1014 Hz

- Molecules by infrared radiation become vibrationally excited.


90

- “stretching” mode: the atoms moving closer and away again. “bending” mode: bond angles periodically increase and decrease.

Vibrational frequencies

- The stiffness of a bond measured by its force constant, k

Force = -k × displacement by Hooke’s law

- Vibrational frequency, n, of a bond between two atoms A and B of mass mA and mB

n = m = m = effective mass (or reduced mass)


91 Normal modes of vibration of polyatomic molecules

A nonlinear molecule consisting of N atoms→ 3N-6 normal modes

A linear molecule → 3N-5 normal modes

i.e. H2O, n = 3 → 3 normal modes

CO2, n = 3 → 4 normal modes


91 Actual spectrum

Characteristic frequencies of functional groups detectable in a spectrum

- fingerprint region: a complex series of absorptions

Documents

2013 General Chemistry I * The coverage of the Quiz (Q) * The coverage of the Study Summary (SS) WeekPeriodChapterChemical Principles / Atkins, JonesNote