2013 General Chemistry I

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Transcript 2013 General Chemistry I

Week

1st 2nd 3rd 4th 5th

Period

3/4-3/7 3/11-3/14 3/18-3/21 3/25-3/28 4/1-4/4

Chapter

2.1-2.16

3.1-3.7

6th 7th 8th 9th 4/8-4/11 4/15-4/18

4/22-4/25

4/29-5/2 3.8-3.12

Mid-term Exam

4.1-4.11

10th 11th 12th 13th 5/6-5/9 5/13-5/16 5/20-5/23 5/27-5/30 4.12-5.6

5.7-5.16

7.1-7.12

7.13-8.8

14th 6/3-6/5 8.9-8.16

15th 6/10-6/13 16th

6/17-6/20 Final Exam * The coverage of the Quiz (Q) 2013 General Chemistry I Chemical Principles / Atkins, Jones

Ionic Bonds, Lewis Structures and Covalent Bonds VSEPR and Valence-Bond Theory Molecular Orbital Theory

(Chapters 1-3)

The Gas Laws, Molecular Motion Real Gases, Intermolecular Forces Liquids and Solids Enthalpy Chemical Change and Entropy Gibbs Free Energy Additional Class / Review

(Chapters 4, 5, 7, 8) Note SS: Chap 2 (3/25) Q: Chap 1 (4/1) SS: Chap 3 (4/8) Q: Chap 2 (4/15) SS: Chap 4 (5/6) SS: Chap 5 (5/13) Q: Chap 4 and 5 (5/20) SS: Chap 7 and 8 (5/27) Q: Chap 7 (6/3)

Chapter 2. CHEMICAL BONDS IONIC BONDS

2.1 The Ions That Elements Form 2.2 Lewis Symbols 2.3 The Energetics of Ionic Bond Formation 2.4 Interactions Between Ions

COVALENT BONDS

2.5 Lewis Structures 2.6 Lewis Structures of Polyatomic Species 2.7 Resonance 2.8 Formal Charge

Chapter 2. CHEMICAL BONDS (화학결합)

Lowering of energy by rearranging valence electrons Understanding the bond formation between atoms ⇒⇒⇒ understanding properties and reactivity of materials

Designing new materials

Ionic bond( 이온결합 ); electron transfer + electrostatic attraction, NaCl Covalent bond( 공유결합 ); sharing electrons, NH 3 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 )

2013 General Chemistry I

Chapter 2. CHEMICAL BONDS



Chemical bond

is the link between atoms.

Key Ideas

Bond formation by lowering of energy

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

Goal

Understanding the bond formation between atoms ⇒⇒⇒

Designing new materials

Lowering of energy by rearranging valence electrons Ionic bond; electron transfer + electrostatic attraction, NaCl Covalent bond; sharing electrons, NH 3 Metallic bond; cations held by a sea of electrons, copper

2013 General Chemistry I

IONIC BONDS (Sections 2.1-2.4)



ionic model

: the description of bonding in terms of ions

ionic solid

: three-dimensional crystalline solid an assembly of cations and anions stacked together in a regular pattern 55 Metal + nonmetal

2013 General Chemistry I

2.1 The Ions That Elements Form



Cations

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

block and on the left of the

block in Period 3 (Al) ⇒ lose electrons down to their noble-gas cores; 1

2 ( duplet ) or

6 ( octet ) Li + , Be 2+ , Na + , Mg 2+ , Al 3+ , ··· Metallic elements of the

block in Periods 4 and later ⇒ lose electrons down to their noble-gas cores surrounded by

10 ; (

– 1)

6 The

electrons, in most cases, cannot be lost.

Ga 3+ ([Ar]3

10 ), ··· Elements in the

block ⇒

-electrons first, then variable number of (

– 1)

-electrons Fe 2+ ([Ar]3

6 ), Fe 3+ ([Ar]3

5 ), ···

2013 General Chemistry I

Many elements in the

and

blocks; variable valence ⇒ Inert-pair effect;

-electrons alone or all their valence

- and

-electrons In + , In 3+ , Sn 2+ , Sn 4+ , ··· The order of losing electrons: 

Anions

: Add electrons until the next noble-gas configuration is reached 56

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2.2 Lewis Symbols

valence electrons

– depicted as dots; a pair of dots for paired electrons 58 (1916) - cations and anions

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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: CaCl 2 ; empirical formula

2013 General Chemistry I

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.

2013 General Chemistry I

2.4 Interactions Between Ions

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

2013 General Chemistry I

2.4 Interactions Between Ions

- Coulomb potential energy of the interactions of two individual ions 60 e is the fundamental charge; z two ions; r 12 is the distance between the centers of the ions; e 0 vacuum permittivity.

1 and z 2 are the charge numbers of the is the

2013 General Chemistry I

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

toward right hand side toward left hand side

; 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

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- real potential energy of an ionic solid →

attractive and repulsive interionic interactions in close range.

Short-range repulsions between ions Total potential energy =

+

*

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Born-Meyer equation ; correction for repulsions to the Madelung constant repulsive effect The coulombic interaction between ions in a solid is large when the ions are small and highly charged .

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COVALENT BONDS

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

Quantum mechanical view of covalent bond Chemical bonding in H 2 + by sharing an electron between two protons An electron

between

the two nuclei exerts an attractive force on the nuclei.

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2.5 Lewis Structures

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

octet rule

: atoms go as far as possible toward completing their octets Nonmetal atoms share electrons to complete their octet; lines (bonding pairs), dots (lone pairs) - A line (-) represents a shared pair of electrons :

a bond

2013 General Chemistry I

2.6 Lewis Structures of Polyatomic Species

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

Methane; CH 4 (why not CH 3 nor CH 5 ?) 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.

2013 General Chemistry I

 Writing a Lewis structure -

terminal atom

: bonded to only one other atom

central atom

: bonded to at least two other atoms - The element with the lowest ionization energy (less greedy) as the central atom

electronegativity

is a better indicator Ex. HCN - Atoms symmetrically around the central atom; SOS for S 2 O (Exception; NNO) - OH is attached to the central atom in oxoacids; HO–Cl for HClO

i.e.

H 2 SO 4 ----- (HO) 2 SO 2 - Polyatomic ions; total number of electrons should be adjusted to represent the overall charge 65

2013 General Chemistry I

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.

Form multiple bonds if in short of electron pairs.

Step 5: A line for each bonding pair

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Take care of charges.

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Ex 2.4

Lewis structures for CH 3 COOH (multiple central atoms) # of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12 CH 3 COOH Lewis structures for C 2 O 2 H 4 (multiple central atoms)

How many structures can you draw?

2013 General Chemistry I

Lewis structures for NO 3 # of electron pairs = (5 + 3 6 + 1) / 2 = 12

3 12

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

- 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.

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

Benzene, C 6 H 6

- All the carbon-carbon bonds with the same length - Only one 1,2-dichlorobenzen exists.

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2.8 Formal Charge



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 70 V = the number of valence electrons in the free atom L = the number of electrons present on the bonded atom as lone pairs B = the number of bonding electrons on the atom Formal charge indicates the degree of redistribution of electrons relative to free atoms

not 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.

2013 General Chemistry I

OCO with lower formal charges is more likely for CO 2 than COO.

HCN has lower formal charges than HNC.

NNO with lower formal charges is more likely for ON 2 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

formal charge (2.12) oxidation state ).

2013 General Chemistry I

Formal charges differ from oxidation numbers!

Neither of them is the true charge.

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

Ca 2+ 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 2013 General Chemistry I

Odd number of electrons

2.9 Radicals and Biradicals

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



Radicals

: species with an unpaired electron, highly reactive

Biradicals

: molecules with two unpaired electrons

2.10 Expanded Valence Shells

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

PCl 5

.

NCl 5 72

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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 I 3 – ion 3 x 7 + 1 = 22 electrons, 11 electron pairs 2 bonds, 2 3 + 2 = 8 lone pairs Remaining one pair into the central I

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Lewis structures for C 2 O 2 H 4 (multiple central atoms) # of electron pairs = (4 + 3 + 4 + 2 6 + 1) / 2 = 12 2 3

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3 3 3 3

variable covalence

: the ability to form different numbers of covalent bonds 72 Ex 2.8

Dominant resonance Lewis structure of SO 4 2– 5 6 + 2 = 32 valence electrons, 16 electron pairs

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number of resonance structures most preferred structure

THE PERIODICITY OF ATOMIC PROPERTIES

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2.11 The Unusual Structures of Some Group 13/III Compounds

76 - boron and aluminum - incomplete octet: fewer than eight valence electrons - completing octets by a

coordinate covalent bond

, in which both electrons come from one of the atoms

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2.11 The Unusual Structures of Some Group 13/III Compounds

76 Possible due to the atomic radius of Al in AlCl 3 larger than that of B in BCl 3 .

borane 2013 General Chemistry I

Chapter 2. CHEMICAL BONDS

Ionic bond; electron transfer + electrostatic attraction, NaCl 

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

Covalent bond; sharing electrons, NH 3 

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

2013 General Chemistry I

76 lower energy structure the average charge on each atom is not zero.

partial charges

: the charges on the atoms -

polar covalent bond

: a bond in which ionic contributions to the resonance result in partial charges -

electric dipole

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

electric dipole moment (

)

Unit: Debye ( D )

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

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

2.12 Correcting the Covalent Model: Electronegativity



Electronegativity (

)

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

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2.12 Correcting the Covalent Model: Electronegativity

Mulliken’s electronegativity

scale (1934); properties of an isolated atom Exactly defined - Mulliken scale: c = ½(I + E ea ) average of the ionization energy and electron affinity 77

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- rough rules of thumb ionic polar covalent covalent i.e. NaCl or KF : ionic C-O : polar covalent Ca-O : ionic

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2.13 Correcting the Ionic Model: Polarizability

- 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 + , Be 2+ , Mg 2+ , and Al 3+ 78

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THE STRENGTHS AND LENGTHS OF COVALENT BONDS

2.14 Bond Strength

 measured by

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 selection of organic molecules, such as methane (CH 4 ), ethane (C 2 H 6 ), and ethene (C 2 H 4 )

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2.15 Variation in Bond Strength

Strongest bond between two nonmetal atoms; CO (1062 kJ/mol) 80 Lone pair-lone pair repulsion due to the short F–F distance

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ATP(

) + H 2 O ADP(

) + H 2 PO 4 – (

aq)

+ 174 kJ/mol

bond stiffness

( bond strength); resistance to stretching and compressing will be discussed in the Major Technique 1: Infrared (IR) Spectroscopy

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2.16 Bond Length



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 82

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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

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

Covalent radius

: contribution an atom to the length of a covalent bond - Decreases from left to right (increasing Z eff ) - Increases in going down a group (size of valence shells and better shielding by inner core electrons) 82 Bond length: Approximately the sum of the covalent radii of the two atoms

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INFRARED SPECTROSCOPY



Infrared radiation

: electromagnetic radiation with longer wavelengths (lower frequencies) than red light ~ 1000 nm or ~ 3×10 14 Hz - Molecules by infrared radiation become

vibrationally excited

bond stiffness

( bond strength); resistance to

stretching

and compressing

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“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 m A and m B n = 1 2π k µ m = 𝑚 𝐴 𝑚 𝐵 𝑚 𝐴 +𝑚𝐵 m = effective mass (or reduced mass)

2013 General Chemistry I



Normal modes of vibration of polyatomic molecules

A nonlinear molecule consisting of N atoms → 3N-6 normal modes i.e. H 2 O, n = 3 → 3 normal modes A linear molecule → 3N-5 normal modes CO 2 , n = 3 → 4 normal modes

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

Actual spectrum

 Characteristic frequencies of

functional groups

detectable in a spectrum 91 - fingerprint region: a complex series of absorptions

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