Laboratory 06

MOLECULAR GEOMETRY AND POLARITY

Background- Lewis structureDiagrams that show the bonding between atoms of a molecule and the lone pairs of electrons that may exist in the molecule

A Lewis structure can be drawn for any covalently bonded molecule, as well as coordination compounds.

Construction of Lewis Structures

Two Rules

1. Total # of valence electrons – the total number of valence electrons must be accounted for, no extras, none missing.

2. Octet Rule – every atom should have an octet (8) electrons associated with it.

Hydrogen should only have 2 (a duet).

Determining the number of valence electrons

Full d-orbitals do not count as valence electrons.They belong to the inner shell.

For example:

Pb [Xe]4f145d106s26p2

This is four (4) valence electrons. The 5d is part of the inner shell (n=5) which is full.

The total number of available valence electrons is just thesum of the number of valence electrons that each atom possesses (ignoring d-orbital electrons)

For H2O, The total number of valence electrons = 2 x 1 (each H is 1s1) + 6 (O is 2s22p4) = 8

For CO2 The total number of valence electrons = 4 (C is 2s22p2) + 2 * 6 (O is 2s22p4) = 16

Central AtomIn a molecule, there are only 2 types of atoms:

1. “central” – bonded to more than one other atom.2. “terminal” – bonded to only one other atom.

Almost always the least electronegative atom is the centralatom. For example, in ClO2, the Cl is the central atom;

in SF5 the S is the central atom.

You can have more than one central atom in a molecule.

Hydrogen never is the central atom. It forms only one bond,so it must generally be in the outer layer of atoms.

Bonds

Bonds are pairs of shared electrons.

Each bond has 2 electrons in it.

You can have multiple bonds between the same 2 atoms.

For example: C-OC=OC O

Each of the lines represents 1 bond with 2 electrons in it.

Lewis Dot Structure

Each electron is represented by a dot in thestructure

.

:Cl: ¨That symbol with the dots indicate a chlorine atom with 7 valence electrons.

Drawing Lewis Dot Structures

1. Determine the total number of valence electrons.2. Determine which atom is the “central” atom.3. Stick everything to the central atom using a single

bond.4. Fill the octet of every atom by adding dots.5. Verify the total number of valence electrons in the

structure.6. Add or subtract electrons to the structure by

making/breaking bonds to get the correct # of valence electrons.

7. Check the “formal charge” of each atom.

Formal Charge of an Atom

Formal charge = number of valence electrons – number of bonds – number of non-bonding electrons.

Dot structure for H2O

1. Total number of valence electrons: 6 + (2 x 1) =8

2. Central Atom – O

3. Stick all terminal atoms to the central atom using a singlebond.

Dot structure for H2O

H – O - H

Dot structure for H2O

..

H – O – H ¨

That is a total of 8 valence electrons used: eachbond is 2, and there are 2 non-bonding pairs.

FC (H) = 1-1-0 = 0FC (O) = 6 – 2 – 4 = 0

Expanded Octets

Example PCl5: .. .. :Cl: :Cl: Total valence e- = 40 .. .. :Cl – P - Cl : FC(P) = 5 – 5 – 0 =0 ¨ | ¨ : Cl: FC (Cl) = 7 – 1 – 6 = 0 ¨

Background - Covalent Bonds• The simplest covalent bond is that in H2

– the single electrons from each atom combine to form an electron pair

– the shared pair functions in two ways simultaneously; it is shared by the two atoms and fills the valence shell of each atom

• The number of shared pairs– one shared pair forms a single bond– two shared pairs form a double bond– three shared pairs form a triple bond

Polar and Nonpolar Covalent Bonds• Although all covalent bonds involve sharing of electrons, they

differ widely in the degree of sharing• We divide covalent bonds into

– nonpolar covalent bonds– polar covalent bonds

Difference in ElectronegativityBetween Bonded Atoms Type of BondLess than 0.50.5 to 1.9Greater than 1.9

Nonpolar covalentPolar covalentIons form

Background – Polarity

An example of a polar covalent bond is that of H-Cl

– the difference in electronegativity between Cl and H is 3.0 - 2.1 = 0.9

– Polarity can be shown by using the symbols ++ and --, or by using an arrow with the arrowhead pointing toward the negative end and a plus sign on the tail of the arrow at the positive end

H Cl+ -

H Cl

Polar bonds and polar moleculesN

HH

HO

H H

Water = 1.85D

Ammonia = 1.47D

directionof dipolemoment

directionof dipolemoment

Laboratory 07

QUALITATIVE ANALYSIS : TESTING THE SOLUBILITY RULES

Background : Ionic Compounds

1. Most ionic compounds are also called salts.

2.Most ionic compounds exist as solids and many dissolve to form aqueous solutions.

Example : AgCl insoluble in water but AgNO3 is soluble

3. An ionic compound is made up of a metal and a nonmetal; metals are located on the left side of the periodic table and nonmetals are on the right side.

4. The cation (positive ion) is written first followed by the anion (negative ion).

Nomenclature of binary ionic compounds

Symbol Anion Symbol

Anion Name

Br Br- Bromide

Cl Cl- ChlorideF F- FluorideH H- HydrideI I- IodideN N-3 NitrideO O-2 OxideP P-3 PhosphideS S-2 Sulfide

NaClNaCl

NaFNaFHH22SS

Sodium chloride

Sodium fluoride

Hydrogen sulfide

BaClBaCl22

KK22OOMgMg33NN22

Barium chloride

Potassium oxide

Magnesium nitride

Polyatomic anions

NO3- = nitrate NO2

- = nitrite

SO4 2

- = sulfate SO3

2- = sulfite

PO43- = phosphate CO3

2- = carbonate

HCO3- = hydrogen carbonate or bicarbonate

OH- = hydroxide CN- = cyanide

C2H3O2- = acetate C2O4

2- = oxalate

NaHCO3 = sodium hydrogen carbonate or sodium

bicarbonate

K2SO3 = potassium sulfite

MgSO4 = magnesium sulfate

KCN = potassium cyanide

H2PO4 = hydrogen phosphate

Ca(OH)2 = calcium hydroxide

NH4NO3 = ammonium nitrate

Zn(NO3)2 = zinc nitrate

Li3PO4 = lithium phosphate

HNO3 = hydrogen nitrate

Solubility Rules for Ionic Compounds in Water

Soluble Ionic Compounds Insoluble Ionic Compounds1. All common compounds of Group 1A ions (Li+, Na+, K+, etc.) and ammonium ion (NH4

+) are soluble.

1. All common metal hydroxides are insoluble, except those of Group 1A and the larger members of Group 2A (beginning with Ca2+).

2. All common nitrates (NO3-), acetates

(CH3COO-), and most perchlorates (ClO4

-) are soluble.

2. All common carbonates (CO32-) and

phosphates (PO43-) are insoluble, except

those of Group 1A and NH4+.

3. All common chlorides (Cl-), bromides (Br-), and iodides (I-) are soluble, except those of Ag+, Pb2+, Cu+, and Hg2+. All common fluorides (F-)are soluble, except those of Pb2+ and Group 2A.

3. All common sulfides are insoluble except those of Group 1A, Group 2A, and NH4

+.

4. All common sulfates (SO42-) are

soluble, except those of Ca2+, Ba2+, Ag+, and Pb2+.

Precipitation reactionsGeneral form: Solution A + Solution B → Insoluble Solid C + Solution D.In a precipitation reaction two solutions are mixed together to produce an insoluble solid which is called the precipitate.This type of reaction is also called a double displacement reaction

Lead nitrate(aq) + Potassium iodide(aq) → Lead iodide(s) + potassium nitrate(aq)Pb(NO3)2(aq) + 2KI(aq) → PbI2(s) + KNO3(aq)

Barium chloride(aq) + Sodium sulfate(aq) → Barium sulfate(s) + Sodium chloride(aq)BaCl2(aq) + Na2SO4(aq) → BaSO4(s) + NaCl(aq)

Copper sulfate(aq) + Sodium hydroxide(aq) → Copper hydroxide(s) + Sodium sulfate(aq)CuSO4(aq) + 2NaOH(aq) → Cu(OH)2(s) + Na2SO4(aq)

Silver nitrate(aq) + sodium chloride(aq) →Silver chloride(s) + sodium nirate(aq)AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

Mercury(II) nitrate(aq) + Potassium iodide(aq) → Mercury iodide(s) + Potassium nitrateHg(NO3)2(aq) + 2KI(aq) → HgI2(s) + KNO3(aq)