Molecular Bonds

(Putting Elements Together)

Molar Mass

• Each atom has an atomic mass

• Molar mass is the atomic mass of all the atoms in the molecule summed together

• For Example:

H2O = 2 x Atomic Mass of H +

1 x Atomic Mass of O

Counting Atoms in a Molecule

In the example, NH3, the subscript 3 only applies to the hydrogen.– Therefore: there is 1 N and 3 H in ammonia

In the example, 3Ca3(PO4)2, the number of atoms changes due to the Coefficient in front of the molecule

The 3 is multiplied to the Ca, P and O

The subscript 2, multiplies the P and O

3Ca3(PO4)2

3 Ca3 ( P O4 ) 2

This means that

there are 3 x 3 Ca,

3 x 2 P and 3 x (4 x 2) O

Bonds. . .

No, not that kind – bonds between atoms to form molecules

It all depends upon the atom’s valence (outer shell) electrons

These are the e- in the last Energy Level (n = 1 through 7)

Figure these out using the Periodic Chart and/or Lewis Dot Diagrams

The Roman Numerals Tell You How Many Valence Electrons for the Primary or Representative Elements;

The Valence Electrons for the Transition Elements Vary

I II III IV V VI VII VIII

Group I is monovalent; II is divalent; III is trivalent; IV is tetravalent; V is back to being trivalent (since three e- openings); VI is divalent; VII is monovalent and VIII has a complete octet, so these seldom react or bond

Bond Types (In General):• Pure or Non-Polar Covalent

Χ difference = 0 to 0.5 on the Pauling EN ScaleThe pair of e- shared are done so equallyTwo nonmetals bonded together

• Polar CovalentA shared pair of e-, but not equallyχ difference = 0.5 to 1.6Molecule has Partial + and – Charges

• Ionic Bondsχ difference = 1.7 or higher to the maximum of 4.0Metal bonded with a nonmetal

• Metallic Bonds are similar to Ionic Bonds

Metallic Bonds• Two or more metals mixed are called

alloys

• Two major formats– Interstitial and Substitutional

These bonds permit the

roaming of e- which creates

a sea of dissociated e-

Called the Electron Sea

Model

Ionic Bonds

• These are the bonds between a metal and a nonmetal

• The metal Ion is positively charged and called a cation

• The nonmetal Ion is negatively charged and called an anion

• The bonded molecule should be neutrally charged when finished

Knowing where the metals and nonmetals are on the table will make your life easier

Let’s take a moment to discuss polyatomic ions. . .

• This is a molecule that acts as a cation or anion

• For example:NH4

+ ammonium N3- azide

ClO4- perchlorate CN- cyanide

HCO3- bicarbonate OH- hydroxide

CrzO7-2 chromate NO3

- nitrate

ClO3- chlorate C2H3O2

- acetate

• Don’t PANIC – I gave a list to you!

In an Ionic Bond – one or more electrons are lost or gained by the atoms involvedThis allows the atoms to have a complete valence shell – following the octet rule

In an Ionic Compound – balance the molecule using the criss-cross rule

Mg +2 + Cl-1

Mg Cl2 The one is understood.

This applies even if using a polyatomic ion

NH4+ + O-2

(NH4)2O The parentheses are used to keep

the polyatomic together

Pb+4 + CO3-2

Pb2 (CO3)4 and this can be simplified by

reducing the subscripts to

Pb(CO3)2

Naming Ionic Compounds is really simple:

1. Name the cation (metal) using its proper name; if it is a polyatomic, do the same2. Then, using the stem of the anion (nonmetal), simply add the suffix “ide”

Zinc + Chlorine = Zinc ChlorideIron + Oxygen = Iron OxideLithium + Cyanide = Lithium CyanideAmmonium + Fluorine = Ammonium FluorideCobalt + Phosphorous = Cobalt Phosphide

Transition Metals present an issue for balancing and naming molecules since they can have varying oxidation states

For example:

Manganese can be a +2 or +3

Iron can be a +2 or +3

Lead can be a +2, or even a +4

Copper is a +1 or +2

Gold is usually a +1 or +3

And Hydrogen is a +1 or a -1!

Transition Metals

• To determine the correct Roman Numeral to place after the metal:

Roman Numeral = - (Charge # anion)(#anions)

(# cations)

This is needed because, for example,

iron chloride can be either FeCl2 or FeCl3;

or iron (II) chloride or iron (III) chloride

Therefore – Ionic Bonds are:

Metal + Nonmetal• + ion - ion• cation anion• monatomic monatomic or

(except NH4+) polyatomic

• left of steps right of steps

Reactions are Exothermic

Form Crystal Lattice Structures

Covalent Compounds

• These can be monatomic or polyatomic compounds

• It is a bond between two nonmetals• They share a pair of electrons• They can be subgrouped into polar or

nonpolar• If a binary compound (2 atoms) – use

the same naming rules as in Ionic Compounds

• If it has more than two atoms – need to use the prefixes

Number Prefix Number Prefix

1 Mono 7 Hepta

2 Di 8 Octa

3 Tri 9 Nona

4 Tetra 10 Deca

5 Penta 11 Undeca

6 Hexa 12 Dodeca

Naming Covalent Compounds

Process:

1. Prefix Indicating # + full name of first

nonmetal

2. Prefix Indicating # + root name of second nonmetal + suffix “ide”

3. Watch for polyatomics and use their proper names

For Example:• P4S10 becomes Tetraphosphorous Decasulfide

• P2O5 Becomes Diphosphorous Pentaoxide

• SF6 becomes Sulfur Hexafluoride

• SiBr4 becomes Silicon Tetrabromide

Covalent Bonds can be Polar or Nonpolar

A nonpolar has no discernable

negative or positively charged sides

(EN difference is 0)

A polar covalent bond means one

side is negative and the other positive

Electronegativity Percent Ionic BondDifference Character Type• 0.2 1 % Non-polar• 0.4 4 Covalent• 0.5• --------------------------------------------------------------------------• 0.6 9• 0.8 15• 1.0 22 Polar• 1.2 30 Covalent• 1.4 39• --------------------------------------------------------------------------• 1.6 47 Ionic if metal/nonmetal• 1.8 55 Polar Cov. if non/nonmetal• 2.0 63• --------------------------------------------------------------------------• 2.2 70• 2.4 76 Pure Ionic• 2.6 82• 2.8 86• 3.0 89• 3.2 92

• Some elements are able to form more than one oxyanion (polyatomic ions that contain oxygen), each containing a different number of oxygen atoms.

• For example, chlorine can combine with oxygen in four ways to form four different oxyanions: ClO4

-, ClO3-, ClO2

-, and ClO- (Note that in a family of oxyanions, the charge remains the same; only the number of oxygen atoms varies.)

• The most common of the chlorine oxyanions is chlorate, ClO3

-. In fact, you will generally find that the most common of an element’s oxyanions has a name with the form (root)ate.

• The anion with one more oxygen atom than the (root)ate anion is named by putting per- at the beginning of the root and -ate at the end. For example, ClO4

- is perchlorate.

• The anion with one fewer oxygen atom than the (root)ate anion is named with -ite on the end of the root. ClO2

- is chlorite. • The anion with two less oxygen atoms

than the (root)ate anion is named by putting hypo- at the beginning of the root and -ite at the end. ClO- is hypochlorite.

Oxyanion Example

• ClO- Hypochlorite

• ClO2-Chlorite

• ClO3-Chlorate

• ClO4-Perchlorate

 

• Some compounds have common names as well as their scientific names – you should learn these and others!

– NO nitrogen monoxide nitric oxide

– H2O dihydrogen monoxide water

– NH3 nitrogen trihydride ammonia

– CH4 carbon tetrahydride methane

– C4H10 tetracarbon decahydride butane

Some atoms are Diatomic – KNOW THESE!

H2 N2 O2 F2 Cl2 Br2 and I2

and P is usually found as P4

while Sulfur is found as S8

Other elements will bond beyond the octet rule – like PCl5, and the noble gas Xe bonds with F in XeF6, XeF2, XeF4, and XeO4 – and this is due to a thing called “hypervalence” or “expanded octet”

Molecular Geometry

• The 3-Dimensional Shapes of Molecules depend upon the valence e-’s of the atoms involved

• Valence Bond Theory and VSEPR Model both use the same shapes– Basically – they focus on covalent bonds with

the shared bonding pairs of electrons (BP)– The assumption is made that the molecule will

adopt a geometry to minimize the repulsion between e-’s

• The General Shapes:

Basic Geometry Bond Angles

• Linear 180o

• Trigonal Planar 120o

• Tetrahedral 109.5o

• Trigonal Bipyramidal 90o and

120o

• Octahedral 90o

Molecular Orbital TheoryMOT uses atomic orbitals (AO), e- λ’s and e-

density regions to examine bonds

This is the end of Part I

• Next:– Van der Waals and London Dispersion Forces– Polarity– Intermolecular Forces– Lewis Dot Diagrams with Covalent Bonds– Determining Molecular Structure– Resonance Structures