Matter is composed of either

Matter is composed of either

CHEMICAL BONDING

(1) Metals(2) Nonmetals(3) Metals and Nonmetals

- Atoms- Molecules- Ions

- Metallic Bonding

- Covalent Bonding

- Ionic Bonding

2A-1 (of 15)

Chemical bonding involves the valence electrons of atoms

1904 ARNOLD SOMMERFELDProposed that metal atoms release their valence electrons, and share them between large numbers of metal atoms

(2) MATTER COMPOSED OF METALS

2A-2 (of 15)

METALLIC BOND – The electrostatic attraction of the shared valence electrons to the nuclei of the many bonding metal atoms

Metallic bonding forms crystalline networks containing billions of metal ATOMS that are strongly attracted together

2A-3 (of 15)

1916 GILBERT NEWTON LEWISProposed that nonmetal atoms share valence electrons to achieve the electron configurations of Noble Gases

(3) MATTER COMPOSED OF NONMETALS

2A-4 (of 15)

Diatomic chlorine

: Cl

:: Cl ::

:

: Cl – Cl :

: :

: :

LEWIS STRUCTURE – A representation of chemical bonding using electron dot notation

2A-5 (of 15)

Cl Cl

: :

: :

BONDING PAIRS: in redLONE PAIRS: in green

::

COVALENT BOND – The electrostatic attraction of the shared electrons to the nuclei of the bonding nonmetal atoms

Covalent bonding forms individual units called MOLECULES that are weakly attracted to each other

2A-6 (of 15)

To draw a proper Lewis Structure for a covalently bonded species:

1 – Add up the valence e-s for all of the atoms in the molecule or ion

2 – Draw a skeletal structure by using pairs of electrons to make bonds

4 – If octets are not produced, make the atoms that have octets share more e- pairs with atoms that do not have octets

3 – Complete octets (or duets for H) for all atoms, outer atoms first, using the remaining valence e-s

LEWIS STRUCTURES

2A-7 (of 15)

Oxygen difluoride, OF2

6 + 7 + 7 = 20 valence e-s

F O F

2A-8 (of 15)

Nitrogen tribromide, NBr3

5 + 7 + 7 + 7 = 26 valence e-s

Br N BrBr

2A-9 (of 15)

1904 RICHARD ABEGGProposed that atoms gain or lose valence electrons to achieve the electron configurations of Noble Gases

(1) MATTER COMPOSED OF METALS AND NONMETALS

2A-10 (of 15)

Metal atoms easily lose valence e-s, forming positive ionsAl 1s22s22p63s23p1

Fe [Ar]4s23d6

Nonmetal atoms gain e-s to their valence shells, forming negative ionsO 1s22s22p4

Once these ions are formed, they are stable (or unreactive)

Al3+ 1s22s22p6

Fe2+ [Ar]3d6

O2- 1s22s22p6

Fe3+ [Ar]3d5

2A-11 (of 15)

2A-12 (of 15)

IONIC BOND – The electrostatic attraction between positive metal ions and negative nonmetal ions

Ionic bonding forms crystalline networks containing billions of positive and negative IONS that are strongly attracted together

Sodium chloride

Na . .. Cl : . .

Na+ . .: Cl :

-

. .

A sodium chloride crystal is a symmetrical array of sodium and chloride ions in a 1:1 ratioEMPIRICAL FORMULA – The simplest whole number ratio of ions of different elements in a compoundEmpirical Formula: NaCl

2A-13 (of 15)

Calcium fluoride

Empirical Formula: CaF2

Ca . .. F : . .

Ca2+ . .: F :

-

. .

. .. F : . . . .: F :

-

. .

2A-14 (of 15)

K3N

K .

. . N : .

K . K

.

2A-15 (of 15)

REPRESENTING IONIC BONDING WITH ELECTRON DOT NOTATION

K+ . . : N : 3-

. . K+ K+

Fluorine, F2

7 + 7 = 14 valence e-s

F F

SINGLE BOND – One shared pair of e-s between two atoms

2B-1 (of 15)

Oxygen, O2


O O

DOUBLE BOND – Two shared pairs of e-s between two atoms

2B-2 (of 15)

Nitrogen, N2


N N

TRIPLE BOND – Three shared pairs of e-s between two atoms

2B-3 (of 15)

BOND ORDER – The number of shared pairs of electronsBOND ENERGY – The energy needed to break a bondBOND LENGTH – The distance between the nuclei of the 2 bonding atoms

Bond Order

Bond Energy (kJ/mol)

Bond Length (nm)

F2 O2 N2

1

154

0.142

2

495

0.121

3941

0.110

2B-4 (of 15)

H H P PS SCl Cl I I

Longest Bond Length?Shortest Bond Length?Highest Bond Energy?Lowest Bond Energy?

I2

H2

P2

I2

biggest atomssmallest atomsmost bonding electronsleast bonding electrons, and they are most shielded from the nuclei

2B-5 (of 15)

Formaldehyde, CH2O

4 + 1 + 1 + 6 = 12 valence e-s

H C OH

2B-6 (of 15)

Sulfate, SO42-

6 + 4(6)

OO S O

O

+ 2 = 32 valence e-s

2-

2B-7 (of 15)

NO3-

5 + 3(6) + 1 = 24 valence e-s

O N OO

-O N O

O

- -O N O

O

RESONANCE – When more than one Lewis structure can be drawn for a molecule or ionRESONANCE STRUCTURES – The Lewis structures that can be drawn for the molecule or ionThe bonding in the real nitrate ion is an average of its resonance structuresThe average N-O bond order is (1+1+2) / 3 = 11/3

↔ ↔

2B-8 (of 15)

1932 LINUS PAULINGDescribed how atomic orbitals are involved in covalent bonding

2B-9 (of 15)

VALENCE BOND THEORY – Two atoms share electrons by overlapping a valence atomic orbital from each atom, creating a region of space between the nuclei where the electrons reside

H atom

1s atomic orbitalwith 1 valence e-

H atom

1s atomic orbitalwith 1 valence e-

H2 molecule

The attraction of the e-s in the molecular orbital to the 2 nuclei bonds the atoms together

2B-10 (of 15)

2 valence e-s in aMOLECULAR ORBITAL

Metals – Low EN’s (the most active metals having the lowest EN’s)Nonmetals – High EN’s (the most active nonmetals have the highest EN’s)

ELECTRONEGATIVITY – A property developed by Pauling, measuring the attraction of an atom for shared electrons

2B-11 (of 15)

Atom with the highest EN?F (4.0)

Atom with the lowest EN?Cs (0.7)

Nonpolar Covalent0

EN Difference Bonding

Polar CovalentSmall (0.1 – 1.6)Large (1.7 – 3.3) Ionic

EN differences between atoms indicates their type of bonding

2B-12 (of 15)

2 atoms with the same EN’s have an EN difference of 0N – N (EN of N = 3.0)

NONPOLAR COVALENT BOND – A bond between 2 atoms in which the electrons are shared evenly

2B-13 (of 15)

2 atoms with close EN’s have an EN difference that is smallH – Br (EN of H = 2.1, EN of Br = 2.8)

POLAR COVALENT BOND – A bond between 2 atoms in which the electrons are shared unevenly

Dipole Moment Arrow

2B-14 (of 15)

2 atoms with extreme EN’s have an EN difference that is largeNa – Cl (EN of Na = 0.8, EN of Cl = 3.0)

IONIC BOND – A bond between 2 atoms in which the electrons are transferred, creating ions

2B-15 (of 15)

FORMAL CHARGE

While atoms that covalently bond are not charged, they can be given charges based upon where the bonding electrons are assigned

2C-1 (of 11)

FORMAL CHARGE – The charge given to an atom assuming one electron in each bond is assigned to that atom

..

0 0

F S F

F S F

0

. .

S naturally has 6 valence e-s , and now 6 0F naturally has 7 valence e-s , and now 7 0

2C-2 (of 11)

Quick way to determine formal charge:(natural number of valence e-s – 1 e- per bond – each lone pair e-)

S:F:

6 – 2 – 4 = 07 – 1 – 6 = 0

C O

C O

-1 +1

2C-3 (of 11)

Formal charges are used to determine the validity of a Lewis structure -the most accurate Lewis structures are those with atoms that have formal charges as close to 0 as possible

C:O:

4 – 3 – 2 = -16 – 3 – 2 = +1

S C N

thiocyanate, SCN-

6 + 4 + 5 + 1 = 16 valence e-s

S C N-

S C N- -

↔ ↔

S:C:N:

S:C:N:

S:C:N:

6 – 2 – 4 = 04 – 4 – 0 = 05 – 2 – 4 = -1

6 – 3 – 2 = +14 – 4 – 0 = 05 – 1 – 6 = -2

6 – 1 – 6 = -14 – 4 – 0 = 05 – 3 – 2 = 0

2C-4 (of 11)

thiocyanate, SCN-

6 + 4 + 5 + 1 = 16 valence e-s

S C N-

S C N S C N- -

↔ ↔

The best Lewis structures have(1) formal charges for the most atoms as close to 0 as

possible(2) negative formal charges go on the atom with the greatest EN

2C-5 (of 11)

S:C:N:

S:C:N:

S:C:N:

6 – 2 – 4 = 04 – 4 – 0 = 05 – 2 – 4 = -1

6 – 3 – 2 = +14 – 4 – 0 = 05 – 1 – 6 = -2

6 – 1 – 6 = -14 – 4 – 0 = 05 – 3 – 2 = 0

COVALENT COMPOUNDS THAT DO NOT OBEY THE OCTET RULE

(1) Molecules with hypovalent central atoms

Covalent compounds with B and Be

BeH2

2 + 1 + 1 = 4 valence e-s

H Be H

(atoms with less than 4 valence electrons)

2C-6 (of 11)

BF3

3 + 7 + 7 + 7 = 24 valence e-s

F B FF

NO!

F B FF

F is too electronegative to share more than 1 pair of e-s

2C-7 (of 11)

(2) Molecules with hypervalent central atoms (atoms that have empty d orbitals in their outer shell)

Nonmetal atoms in the 3rd, 4th, 5th, or 6th Periods

PF5

5 + 5(7) = 40 valence e-s

FPF F

F F P can make 5 bonds using empty d orbitals in its outer shell

3s 3p

↑↓ ↑ ↑ ↑ ___ ___ ___ ___

___ ___ ___ ___ ___

3d

2C-8 (of 11)

ClF3

7 + 3(7) = 28 valence e-s

FF Cl F Only 26 valence e-s

FF Cl F

2C-9 (of 11)

Sulfate, SO42-

6 + 4(6)

OO S O

O


2-

Experimental data shows the S-O bonds are stronger than single bondsReducing the formal charge on atoms that can exceed the octet rule can produce a more accurate Lewis structureS must make 2 double bonds to reduce its formal charge to 0

S:O:

6 – 4 – 0 = +26 – 1 – 6 = -1

2C-10 (of 11)

Sulfate, SO42-

6 + 4(6)

OO S O

O


2-

Experimental data shows the S-O bonds are stronger than single bondsReducing the formal charge on atoms that can exceed the octet rule can produce a more accurate Lewis structureS must make 2 double bonds to reduce its formal charge to 0

S:O:O:

6 – 1 – 6 = -16 – 2 – 4 = 0

+ 5 other resonance structures

6 – 6 – 0 = 0

2C-11 (of 11)

MOLECULAR SHAPE

VSEPR THEORY (Valence Shell Electron Pair Repulsion) – All atoms and lone pairs attached to a central atom will spread out as far as possible to minimize repulsionA Lewis structure must be drawn to use the VSEPR Theory

2D-1 (of 15)

CO2

4 + 6 + 6

O C O

= 16 valence e-s

STERIC NUMBER (SN) – The sum of the bonded atoms and lone pairs on a central atom The steric number of carbon is 2 (SN = 2): 2 bonded atoms and 0 lone pairs

LinearBond angle is 180° O C O

2D-2 (of 15)

H B HH

SN = 3 3 bonded atoms and 0 lone pairs

Trigonal Planar Bond angle is 120°

BH3

3 + 1 + 1 + 1 = 6 valence e-s

H

H HB

2D-3 (of 15)

SO2

6 + 6 + 6

O S O

= 18 valence e-s


Bent Bond angle is 120°

O

OS

2D-4 (of 15)

HH C H

H

SN = 4 4 bonded atoms and no lone pairs

Tetrahedral Bond angle is 109.5°

HC

H

H H

CH4

2D-5 (of 15)

H N HH


Trigonal Pyramidal Bond angle is 108°

N

HH H

NH3

2D-6 (of 15)

H2O


BentBond angle is 105°

O

HH

. .H – O : H

2D-7 (of 15)

FF P F

F F

SN = 55 bonded atoms and no lone pairs

Trigonal Bipyramidal3 Equatorial F’s in a

plane, 120° apart2 Axial F’s 180° apart,

90° from the plane

F P F

F

FF

PF5

2D-8 (of 15)

SN = 54 bonded atoms and 1 lone pair

SF4

← 2 close 90º interactions

3 close 90º interactions →

← most stable configuration

e- pair in equatorial position

e- pair in axial position

2D-9 (of 15)

: F :F S F

: F :


SF4

: F :F S F

: F :

FP

F FF

See-Sawe- pairs always go in

equatorial positions to minimize repulsion

2D-10 (of 15)

SN = 53 bonded atoms and 2 lone pairs

ClF3

F Cl FF

F Cl F

F T-Shape

2D-11 (of 15)


XeF2

F Xe F

F Xe

F Linear

2D-12 (of 15)

SN = 66 bonded atoms and no lone pairs

SF6

FSF

FS

F FF

Octahedral 90º and 180º F

F

F

F

F

F

2D-13 (of 15)


IF5

FI

FI

FF

Square Pyramidal FF

F

F

F

F

2D-14 (of 15)


XeF4

Xe

FF

Square Planar FF

FF Xe F

F

2D-15 (of 15)

MOLECULAR POLARITYA BOND is polar if it has a positive end and a negative endA MOLECULE is polar if it has a positive end and a negative end

To determine if a molecule is polar or nonpolar:1) Draw the correct Lewis structure2) Draw its correct shape3) Use EN’s to determine if the BONDS in the molecule are polar or

nonpolar4) For the polar bonds, label the positive and negative ends with δ+ and δ-5) If a line can be drawn separating all δ+’s from all δ-’s, the molecule is

polar, if not its nonpolar

2E-1 (of 13)

. .H – O : H

O

HH

δ+

δ-

δ+δ-

EN’s: O = 3.5, H = 2.13.5 – 2.1 = 1.4 the O-H BONDS are polarAll of the δ+’s can be separated from all of the δ-’s, the H2O MOLECULE is polar

2E-2 (of 13)

δ+

δ-

δ+

δ-

EN’s: N = 3.0, H = 2.13.0 – 2.1 = 0.9 the N-H BONDS are polarAll of the δ+’s can be separated from all of the δ-’s, the NH3 MOLECULE is polar

H N HH

NH

H H δ+δ-

2E-3 (of 13)

FF C F

F

FC

F

F F

EN’s: C = 2.5, F = 4.04.0 – 2.5 = 1.5 the C-F BONDS are polarAll of the δ+’s cannot be separated from all of the δ-’s, the CF4 MOLECULE is nonpolarδ+

δ-

δ-δ+

δ-

δ+

δ-δ+

2E-4 (of 13)

A more exact way to determine if a molecule is polar or nonpolar:1) Draw the correct Lewis structure2) Draw its correct shape3) Use EN’s to determine if the BONDS in the molecule are polar or

nonpolar4) For the polar bonds, draw a DIPOLE MOMENT ARROW pointing toward

the negative end of the bond5) If the dipole moments are symmetrical the molecule is NONPOLAR

2E-5 (of 13)

Dipole moments of equal magnitude are symmetrical if:1) there are 2 dipole moments that are linear

Y X Y

2E-6 (of 13)

Dipole moments of equal magnitude are symmetrical if:2) there are 3 dipole moments that are trigonal planar

Y

X

YY

2E-7 (of 13)

Dipole moments of equal magnitude are symmetrical if:3) there are 4 dipole moments that are tetrahedral

Y

X

YYY

2E-8 (of 13)

O C O

O C O Symmetrical dipole moments the CO2 MOLECULE is nonpolar

2E-9 (of 13)

. .H – O : H

O

HH Assymmetrical dipole moments

the H2O molecule is POLAR

2E-10 (of 13)

H N HH

NH

H HAssymmetrical dipole moments the NH3 molecule is POLAR

2E-11 (of 13)

FF C F

F

FC

F

F F

Symmetrical dipole moments the CF4 molecule is NONPOLAR

2E-12 (of 13)

ClF C F

F

ClC

F

F F

Assymmetrical dipole moments because the C-Cl dipole moment is smaller than the C-F dipole moments the CClF3 molecule is POLAR

2E-13 (of 13)

Matter is composed of either

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