Chem 1151: Ch. 4

Chem 1151: Ch. 4

Forces between particles

Noble Gas ConfigurationNoble Gas Configuration• The first widely-accepted theory for chemical bonding was based on noble

gas configurations.• Noble gases have a filled valence shell (2e- for He and 8 e- for everything

else).• Because noble gases are very unreactive, and chemical reactivity depends on

electronic structure, two scientists (Lewis and Kossel) concluded that this represented stable (i.e., low energy) configuration.

• Formed the basis for the “octet rule” (1916): 8 electrons in valence shell results in greater stability.

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7

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5

1

Noble gases have filled s and p subshells.

Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7th Edition, 2011

Valence Shell e-Valence Shell e-• Valence shell is the outermost shell where electrons reside.• From the electronic configuration (or other means) you can determine

that the valence shell will have the highest number n.• This means that for transition metals, even though the last electrons may

be added to the d subshell for n, the valence shell e- will be found in the s subshell at n+1.– Ex. Yttrium (Y) 1s22s22p63s23p64s23d104p65s24d1

1s 1s 2s 1s 2s 3s

2p

Valence Shell e- in Lewis StructuresValence Shell e- in Lewis Structures• Lewis Structure (electron dot formula): Simplified way to represent

valence shell e-.• Element symbol represents nucleus, with valence shell e- represented by

dots surrounding it.• The number of valence e- can be found by looking at the group number,

except for the transition metals, which all have filled s orbitals at n+1.

O

C

Sr

In

Po

4

Atom # Val Shell e-

2

5

6

C

Sr

In

Po

O 6

Lewis Structure

IonsIons Ion: Atom or molecule that has either lost or gained electrons from

valence shell resulting in a net charge (positive or negative) compared to the number of protons.

Ionization Energy: Energy required to remove an e- from an atom.

Common atomic ions you should know: H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, N3-, P3-, O2-, S2-, F-, Cl-, Br-

For Group A elements, the number of e- gained or lost results in an electronic configuration like that of the noble gases (valence shell octet). If Na lose e-, has electronic configuration of Ne. If Cl gains e-, has electronic configuration of Ar. If Ca loses 2 e-, has electronic configuration of Ar.

Ionic CompoundsIonic Compounds Ionic compounds are formed when valence electrons lost by metal are

gained by non-metal with which it is reacting. Electron(s) cannot be lost from one atom unless there is another atom

available to accept the electron(s). Common atomic ions you should know: H+, Na+, K+, Mg2+, Ca2+, Fe2+, Fe3+, Ag1+, Pb2+, N3-, P3-, O2-, S2-, F-, Cl-, Br-

Common molecular/polyatomic ions you should memorize: OH- (hydroxide), NH4

+ (ammonium), SO42- (sulfate), SO3

2- (sulfite), PO4

3- (phosphate), NO2- (nitrite), NO3

- (nitrate), CO32- (carbonate)

Ionic compounds formed by these ions will have neutral charges.

Pb(OH)2

NaOH

(NH4+)2SO4

NH3

MgCl2

H3PO4

HBr

HCl

AgNO3

Chemical BondingChemical Bonding1. Ionic bond: Attractive force that holds ions of opposite charge together.

– Involves transfer of e- from one component to the other.– Occurs between positively-charged metal (loses 1 or more e-) and non-metal

atom or molecule (gains 1 or more e-).– Usually satisfies octet rule– Common to inorganic chemistry

2. Covalent bond: Formed by sharing of electrons.– Occurs between:

• Two non-metals• Nonmetal and metalloid• Two metalloids

– Usually satisfies octet rule– Common to organic chemistry

Chemical BondingChemical Bonding1. Ionic bond: Attractive force that holds ions of opposite charge together.

2. Covalent bond: Formed by sharing of electrons.

Na Cl+ ClNa

1+ 1-

EN: 0.9 EN: 3.0

0.79 Å 0.91 Å

2.23 Å 0.97 Å

Na Cl

H C

Na Cl

H C

4H C+ CHEN: 2.1 EN: 2.5

H

HH

0.79 Å 0.91 Å

2.23 Å 0.97 Å

Na Cl

H C

Na Cl

H C

Electron-half-equations (for redox rxns)Electron-half-equations (for redox rxns)

Na Cl+ ClNa

1+ 1-

EN: 0.9 EN: 3.0

0.79 Å 0.91 Å

2.23 Å 0.97 Å

Na Cl

H C

Na Cl

H C

Na Na+ + 1e-

Cl + 1e- Cl-

oxidation

reduction

Each of these represents ½ of the rxn

Naming Ionic Binary CompoundsNaming Ionic Binary Compounds

potassium (K+) + chlorine (Cl-)

[Name of Metal] + [nonmetal stem + ide] =

potassium chloride (KCl)

strontium (Sr2+) + oxygen (O2-) strontium oxide (SrO)

3 calcium (Ca2+) + 2 nitrogen (N3-) calcium nitride (Ca3N2)

Some metals may form more than 1 type of charged ion. Exs: Cu+ and Cu2+; Fe2+ and Fe3+

Compounds with these ions would be named by adding a roman numeral equivalent to charge in parentheses after metal name:

copper (Cu+) + chlorine (Cl-)

[Name of Metal] + [nonmetal stem + ide] =

copper(I) chloride (CuCl)

iron (Fe2+) + 2 chlorine (Cl-) iron(II) chloride (FeCl2)

iron (Fe3+) + 3 chlorine (Cl-) iron(III) chloride (FeCl2)

Units of Ionic CompoundsUnits of Ionic Compounds

• Stable form of ionic compound is not a molecule, but a crystal lattice where ions occupy lattice sites.

• Molecular compounds have molecular weight

• Ionic compounds have formula weight

• Although ionic compounds form crystal lattices, we still represent them the same as molecular compounds when discussing formula weight


Lewis Structures for Covalent CompoundsLewis Structures for Covalent Compounds1. Use molecular formula to determine how many atoms of each type.2. Draw a structure with the elements relative to each other.3. Determine total number of valence e- for all atoms.4. Put one pair of e- between each bonded pair of atoms, subtract this number of e-

from total. 5. Use remaining e- to form octets (except for H and He) for all atoms.6. If all octets cannot be satisfied with available e-, move nonbonding e- pairs between

bonded atoms to complete octets (will form double or triple bonds).

Cl + Cl Cl2 7 + 7 = 14 e-

Cl ClUse 2 e- to form bond

Cl ClNow fill in remaining

12 e-

Lone pairs or nonbonding e-

More Lewis Structures for Covalent CompoundsMore Lewis Structures for Covalent CompoundsN + N N2

N + N N N

N + 3H NH3

N + H N H3H

HO + 2H H2O

H + O + H OH

H

Still More Lewis Structures for Covalent CompoundsStill More Lewis Structures for Covalent Compounds

S + 3O SO3

SO

O

6 + 6 + 6 + 6 = 24 e-

O24 - 6 = 18 e-

SO

O

O SO

O

O

1. Confirm all octets satisfied.2. Confirm all e- accounted for.

……And One More Lewis Structures for Covalent And One More Lewis Structures for Covalent CompoundsCompounds

C2H2

C

4 + 4 + 1 + 1 = 10 e-

C


HH10 - 6 = 4 e-

C C HH

C C HH C C HH

Lewis Structures for Polyatomic IonsLewis Structures for Polyatomic Ions1. Follow the same instructions indicated for covalent

compounds, but add or subtract the number of e- indicated by the ion charge.

2. Use molecular formula to determine how many atoms of each type.

3. Draw a structure with the elements relative to each other.4. Determine total number of valence e- for all atoms.5. Put one pair of e- between each bonded pair of atoms,

subtract this number of e- from total. 6. Use remaining e- to form octets (except for H and He) for all

atoms.7. If all octets cannot be satisfied with available e-, move

nonbonding e- pairs between bonded atoms to complete octets (will form double or triple bonds).

Lewis Structures for Polyatomic IonsLewis Structures for Polyatomic Ions

SO42-

SO

O

6 + 6 + 6 + 6 +6 + 2 = 32 e-

O

32 - 8 = 24 e-

SO

O

O


O O

2-

Lewis Structures for Polyatomic IonsLewis Structures for Polyatomic Ions

PO43-

PO

O

5 + 6 + 6 + 6 +6 + 3 = 32 e-

O

32 - 8 = 24 e-

PO

O

O


O O

3-

Shapes of molecules and Polyatomic IonsShapes of molecules and Polyatomic Ions

Molecules have 3-D shapes.•VSEPR (valence-shell electron-pair repulsion) Theory: Electron pairs in valence shell of atoms are repelled by other pairs and try to get as far away from each as possible. •Shapes around central atom (any atom bonded to 2 or more other atoms) can be predicted by VSEPR.

2 Rules for determining shape:1.All valence shell electron pairs around central atom are counted equally (both bonding and non-bonding).2.Double or triple bonds are treated like a single bond when predicting shapes.

Electron Pair ArrangementsElectron Pair Arrangements

• According to the VSEPR theory, the arrangement of electron pairs around the central atom (represented by E) depends on the number of electron pairs.– Two pairs locate opposite each other. – Three pairs arrange themselves in a flat triangle around the

central atom. – Four pairs become located at the four corners of a pyramid-

like shape called a tetrahedron.


Shapes of molecules and Polyatomic IonsShapes of molecules and Polyatomic Ions

LinearLinear

TriangularTriangular

TetrahedralTetrahedral

OO N+

OO C CH CH ClCl C C

ClCl

CC

CH3

CH3F

F

F

B

Cl

ClCl

Cl

C H

HOH

Cl

CH3

CC

CH3HH

O

H

H

HH N+

One Molecule, Different GeometriesOne Molecule, Different Geometries

CH2OH

O CH3

C

C

CH3

C

C

CH3

Polarity of BondPolarity of Bond1. Nonpolar Covalent Bond: Electrons forming bond between 2

atoms spend nearly equal time between both atoms. 2. Polar Covalent Bond: Covalent bond where electrons

polarized and remain closer to atom with higher EN (bond polarization).– Electronegativity (EN): Ability of an atom to attract shared e- of

covalent bond.

3. Ionic Bond: Electrons are transferred (highly polar).

Cl Cl H Cl

nonpolar polar

Na Cl

ionic

1.2EN1.21.0 EN0EN

Polarity indicated by arrow pointed towards destination of e-

Polarity of MoleculesPolarity of Molecules1. Nonpolar Molecules: Charge distribution from bond

polarizations is symmetric. 2. Polar Molecules (dipoles): Charge distribution from bond

polarizations is not symmetric. 3. Need to consider all charges to determine direction of

polarity.


Binary Covalent CompoundsBinary Covalent Compounds


• Hazards of dihydrogen monoxide• http://www.dhmo.org/facts.html

Naming Binary Covalent CompoundsNaming Binary Covalent Compounds

1. Give the name of the less EN element first.

2. Give the stem of the name of the more EN element next, and add suffix –ide.

3. Indicate the number of each type of atom in molecule using numeric prefix.


CO

CO2

H2O

N2O5carbon monoxide

carbon dioxide

dihydrogen monoxide

dinitrogen pentoxide

CCl4 carbon tetrachloride

S2O7 disulfur heptoxide

Naming ionic compounds containing polyatomic ionsNaming ionic compounds containing polyatomic ions1. Give the name of the metal first.2. Make sure that charges add up to zero.3. Put parentheses around polyatomic ions if more than 1 used.

potassium phosphateK and PO43-

Mg and PO43-

K3PO4

sodium nitrateNa and NO3- NaNO3

Mg3(PO43-)2 magnesium phosphate

NO3- and NH4

+ NH4NO3 Ammonium nitrate

Write formulas for the following and name them:

Crystal Lattices of non-ionic compoundsCrystal Lattices of non-ionic compounds1. Most pure substances (elements or compounds) form crystal lattices in solid state.2. These lattice sites may be occupied by neutral atoms or molecules instead of ions.3. Components of lattice sites held together by covalent bonds.4. Network Solids: Lattice formed by atoms bound in covalent bonds (ex. Si and O).5. Metallic Bond: Lattice of metal ions

– In this structure, valence e- can move about more freely, which is how metals can conduct heat and electricity.

Graphite: Each Carbon is covalently bonded to 3 other carbons in ring

Diamond: Each carbon is bonded to 4 other carbons http://www.eduys.com/Copper-

Molecular-Structure-Model-303.html

Copper

http://en.wikipedia.org/wiki/Image:Graphit_gitter.png

http://en.wikipedia.org/wiki/Image:Diamond_unit_cell.PNG

Interparticle ForcesInterparticle Forces

1. Dipolar forces: Attraction between positive end of one polar molecule and negative end of another (usually weak, ~0.5 -2.0 kcal)

2. Hydrogen Bonding: Attractive interaction of a hydrogen atom with an electronegative atom (e.g. N, O, F) from another molecule or chemical group.– The H must be covalently bonded to

another electronegative atom. – Stronger than dipolar and dispersion (12-

16 kcal)3. Dispersion Forces: Momentary

nonsymmetric electron distributions in molecules (very weak)

O

R R

O

R R

δ-

δ-

δ+

δ+

Intermoleculardipolar

H HO

Hydrogen bond δ+

δ-

H HO

δ+

δ-

Relative Strengths of Interparticle ForcesRelative Strengths of Interparticle Forces


Chem 1151: Ch. 4

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