chem 120: introduction to inorganic chemistry instructor: upali siriwardane (ph.d., ohio state...

CHEM 120: Introduction to CHEM 120: Introduction to Inorganic ChemistryInorganic Chemistry

Instructor: Upali Siriwardane (Ph.D., Ohio State University)

CTH 311, Tele: 257-4941, e-mail: upali@chem.latech.edu

Office hours: 10:00 to 12:00 Tu & Th ; 8:00-9:00 and 11:00-12:00 M,W,& F

Chapters Covered and Test datesChapters Covered and Test dates• Tests will be given in regular class periods  from  9:30-10:45 a.m. on

the following days:

September 22,     2004 (Test 1): Chapters 1 & 2

• October 6,         2004(Test 2):  Chapters  3, & 4

• October 20,         2004 (Test 3): Chapter  5 & 6

• November 3,        2004 (Test 4): Chapter  7 & 8

• November 15,      2004 (Test 5): Chapter  9 & 10

• November 17,      2004 MAKE-UP: Comprehensive test (Covers all chapters

• Grading:• [( Test 1 + Test 2 + Test3 + Test4 + Test5)] x.70 + [ Homework + quiz average] x 0.30 = Final Average

•                               5

Chapter 4: Structure and properties of ionic and covalent

compoundsWe now put atoms and ions together

to form compounds

Chapter 4. Structure and Properties of Ionic and Covalent Compounds

1. Classify compounds as ionic, covalent, or polar covalent bonds. 2. Write the formulas of compounds when provided with the name of the

compound. 3. Name common inorganic compounds using standard conventions and

recognize the common names of frequently used substances. 4. Predict the differences in physical state, melting and boiling points,

solid-state structure, and solution chemistry that result from differences in bonding.

5. Draw Lewis structures for covalent compounds and polyatomic ions. 6. Describe the relationship between stability and bond energy. 7. Predict the geometry of molecules and ions using the octet rule and

Lewis structure. 8. Understand the role that molecular geometry plays in determining the

solubility and melting and boiling points of compounds. 9. Use the principles of VSEPR theory and molecular geometry to

predict relative melting points, boiling points, and solubilities of compounds.

Start learning the formulas and the names and charges of the

ions found in table

• Why have we been so interested in where the electrons are in an atom? And what is the importance of valence electrons?

• Valence e’s are involved in_______--the no of valence e’s has an important influence on ______ of bonds formed. The filled inner core does not directly affect bond formation.

Compound

• Bonds are formed by a transfer of ________ from one atom to another or by a ______ _________ between 2 atoms.

Lewis (dot) Symbols

Lewis (dot) symbols

• Introduced by G. N. Lewis

• Useful for representative (sp block) elements only

• Group no. = no of valence e-’s (no of dots)

Lewis symbols for A groups

• The elements’ symbol represents the inner core of electrons. Put a dot for each valence electron around the symbol.

• Remember that the no. of valence electrons for the A groups is equal to ?

• Each unpaired electron may be used in bond formation

Remember the octet rule from chapter 3

• So the ions formed by the elements in:

• IA

• IIA

• IIIA

• VA

• VIA

• VIIIA

Ionic bonding

• Extra stability has been noted for the noble gas configuration (8 e-s in valence shell)--(for A elements)

• Ionic bonding

• Each atom in the ionic bond

• Ionic compounds are formed between

• And

• When forming an ionic bond each atom in the bond attains a noble gas configuration by a “complete” transfer of

• An ionic bond is the electrostatic force that holds ions together in an ionic compound

• An ionic bond is a very strong bond; ionic cmpds have high m and b pts.

Typical ionic reactions with Lewis structures

+ -Na + F Na F

What about Li and S?

+ 2-Li + S 2 Li S2

What about Ca and O

• Formula is

What about Ca and N?

• Formula is

Covalent bonding

• Not all bonds are ionic.

• ________ bonds are bonds in which two (or more) electrons are ______ by two atoms.

• One shared electron pair is

• A reminder:• Only valence electrons are involved in

bonding. Group No. = # valence e-s for A elements.

• Covalent bonds are formed

• Each atom in bond attains noble gas configuration by sharing of e- pairs (H2 bond only has 2 e-’s)

Covalent bond formation

• Look at formation of H2 molecule.

• H. + .H ----> H:H (H-H)

1s1 1s1 bond formed by overlap of 1s orbitals

What about F2 or Cl2?

____ _____ - pairs of valence electrons not involved in covalent bond formationLewis structure - representation of covalentbonding in which lone pairs are shown as pairs of dots and bonding pairs are (usually) shown as lines

Cl Cl2 Cl Cl Clor

Lonepairs

Bondingpairs

Usualrepresentation

Polar covalent bonding and electronegativity

• Not all covalent bonds are formed btn the same 2 atoms (as H2, homonuclear diatomic: _______sharing of e-’s in bond)

Polar covalent bonds

• What about the bond in H-F?

• It is known that F is more likely to attract e-’s to itself than H, leading to an unequal sharing of the e- pair.

• The covalent bond in which there is unequal sharing:

H F FH

Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms

electron richregion

electron poorregion e- riche- poor

+ -

9.5

H Cl

Cl Cl

+ -Na Cl

Continuum of bond polarity

•(Nearly) complete e- transfer = ionic bond

•Unequal sharing of e- pair = polar covalent bond.

e-s are polarized toward Cl

•Equal sharing of e- pair = nonpolar covalent bond

Electronegativity

• Electronegativity:

• .

• Eneg is a relative concept. Elements with

Lanthanides 1.1-1,3Actinides 1.3-1.5

Electronegativity differences

• 0.2 - 0.5 will be a ________________ bond

• 0.5 - 1.6 will be a ________________ bond

• > 1.6 will be a ________________ bond

Electronegativity differences

• In general the _______ the difference in eneg btn the 2 atoms in the bond, the ____ ______ the bond.

• If the difference is zero, bond (equal sharing of electron pair(s)

(H2, Cl2, O2, F2, N2)

• If the difference is >0 and <1.9, have a :

HCl (3.0 - 2.1); HF (4.0-2.1); OH (3.5-2.1)

• If the difference is > 1.9, have

NaCl (3.0-0.9); CaO (3.5-1.0)

Classify as ionic or covalent

• NaCl

• CO

• ICl

• H2

• Which bond is the most polar (most ionic), which the least polar (most covalent)?

• Li-F Be-F B-F C-F N-F O-F F-F

• Classify the following bonds as ionic, polar covalent, or covalent.

A) the CC bond in H3CCH3

• B) the KI bond in KI

• C) the NB bond in H3NBCl3

• D) the CF bond in CF4

Chemical formulas

• Express composition of molecules (smallest unit of covalent cmpds) and ionic compounds in chemical symbols– H2O, NaCl

Writing formulas for ionic cmpds

• Compounds are neutral overall. Therefore

– NaCl is array of Na+ and Cl- ions – Na2S is array of Na+ and S2- ions

Predict the formulas for the cmpd formed btn

• Potassium and chlorine

• Magnesium and bromine

• Magnesium and nitrogen

Symbol Name Symbol Name

H+ Hydrogen ion H- Hydride ion

Li+ Lithium ion F- Fluoride ion

Na+ Sodium ion Cl- Chloride ion

K+ Potassium ion Br- Bromide ion

Be2+ Beryllium ion I- Iodide ion

Mg2+

Magnesium ion O2- Oxide ion

Ca2+ calcium ion S2- Sulfide ion

Ba2+ barium ion N3- Nitride ion

Zn2+ zinc ion P3- Phosphide ion

Formula Name Formula Name

NO3- nitrate CO3

2- carbonate

NO2- nitrite SO4

2- sulfate

CN- cyanide SO32- sulfite

MnO4- permanganate PO4

3- phosphate

OH- hydroxide PO33- phosphite

O22- peroxide ClO4

- perchlorate

HCO3- hydrogen carbonate ClO3

- chlorate

HSO4- hydrogen sulfate ClO2

- chlorite

HSO3- hydrogen sulfite ClO- hypochlorite

HPO42- hydrogen phosphate CrO4

2- chromate

H2PO4- dihydrogen phosphate C2H3O

- 2 acetate

Symbol (Stock system) Common Symbol (Stock system) Common

Cu+ copper(I) cuprous Hg22+ mercury(I) mercurous

Cu2+ copper(II) cupric Hg2+ mercury(II) mercuric

Fe2+ iron(II) ferrous Pb2+ lead(II) plumbous

Fe3+ iron(III) ferric Pb4+ lead(IV) plumbic

Sn2+ tin(II) stannous Co2+ cobalt(II) cobaltous

Sn4+ tin(IV) stannic Co3+ cobalt(III) cobaltic

Cr2+ chromium(II) chromous Ni2+ nickel(II) nickelous

Cr3+ chromium(III) chromic Ni4+ nickel(IV) nickelic

Mn2+ manganese(II) manganous Au+ gold(I) aurous

Mn3+ manganese(III) manganic Au3+ gold(III) auric

Polyatomic ions Table

• Just have to memorize

• NH4+ ammonium ion

• CO32- carbonate ion

• CN- cyanide ion

• HCO3- hydrogen (or bi) carbonate ion

• OH- hydroxide

• NO3- nitrate ion

• NO2- nitrite ion

• PO43- phosphate ion

• SO42- sulfate ion

• HSO4- hydrogen sulfate ion

• SO32- sulfite ion

• CH3COO- (C2H3O2-) acetate ion

• These polyatomic ions also form ionic cmpds when they are reacted with a metal or a nonmetal in the case of the ammonium ion (or with each other as ammonium sulfate). These polyatomic species act as a

• So the formula for the cmpd formed btn the ammonium ion and sulfur would be:

••• and between calcium and the phosphate ion:

•

• Ionic cmpds do not exist in discrete pairs of ions. Instead, in the solid state, they exist as a three dimensional array--crystal lattice --of cations and anions--are neutral overall,

Given name, write formula

• potassium oxide

• magnesium acetate

Naming ionic cmpds

• Name the cation and anion but drop the word ion from both. This includes the polyatomic ions.

• Na2S

• Ca3N2

Name

• Na3PO4

• NH4Cl

• K2S

Cations with more than one charge

• Cu+ copper(I); Cu2+ copper(II)

• So Cu2O is and

• CuO is

Given name, write formula

• Ammonium chloride

• potassium cyanide

• silver oxide

• Magnesium chloride

• Sodium sulfate

• Iron(II) chloride

To name covalent cmpds

• Name the parts as for ionic cmpds (CO: carbon and oxide) but tell how many of each kind of atom by use of Greek prefixies. (Table 4.4)

• The mono- (for 1) may be omitted for the first element

• Prefix meaning• Mono- 1• Di- 2• Tri- 3• Tetra- 4• Penta- 5• Hexa- 6• Hepta- 7• Octa- 8• Nona- 9• Deca- 10

• CO •

• CO2

• P4S10

• • Boron trichloride

• Water H2O Ammonia NH3

Write formula

• Diboron trichloride

• Sulfur trioxide

• Potassium sulfide

Covalent cmpds

• Remember covalent cmpds--• A _________ is the smallest unit of a covalent

cmpd that retains the characteristics of the cmpd. Molecule - two or more atoms in a definite arrangement held together by chemical bonds. (H2O, Cl2) [Cl2 is considered a molecule but not a cmpd]

• Molecular cmpds exist as

Comparison of properties of ionic and covalent cmpds

• Physical state:

• Ionic cmpds are

• Molecular cmpds can be

Comparison continued

• Melting (___________) and boiling (_________) pts

• In general the melting and boiling temps are much _______for ionic cmpds than for molecular (covalent) cmpds. The ionic bond is very strong and requires a lot of (heat) energy to break the bond. The bond btn molecular species is not as strong.

Comparison continued

• Structure in solid state:

• Ionic solids--

• Covalent solids--

Comparison continued

• In aqueous (H2O) solution:

• Ionic cmpds dissociate into the

• Many covalent cmpds when dissolved in water retain their structure and molecular identity

• Learn the names, formulas, charges, etc for those ions highlighted in table 4.3.

• HCO3-: you should learn as bicarbonate

Writing Lewis structures for covalent species

• These rules are for covalently bonded cmpds only (btn 2 or more nonmetals)

• Do not use them for ionic cmpds.

• 1. Count the total no. of valence electrons (the group no. is equal to the no. of valence electrons).

• if the species is an anion, increase the no. of valence electrons by the charge on the ion

• if the species is a cation, subtract the charge of the cation from the total no. of valence electrons.

• 2.Count the total no. of atoms, excluding H, in the molecule or ion. Multiply that no. by 8.

• Exception: multiply the no. of H’s by 2.

• This tells you how many electrons you would need if you were putting 8 electrons around all atoms without any sharing of electrons (and 2 around all H’s).

• 3. Subtract the no. of e-’s calculated in step 1 from the no. in step 2. This gives you the no. of e-’s that must be shared to get an octet around all atoms in the molecule.

• 4. no. of e-’s that must be shared /2 gives you the no. of bonds.

• 5. subtract the no. of e-’s that are shared (from step 3) from the total no. of valence e-’s. This gives you the no. of unshared e-’s.

• If you divide the no. of unshared e-’s by 2 you get the no. of lone pairs.

• Write the skeletal structure and fill in with the info you came up with. After you’ve put in the # bonds calculated, fill in the octets.

• H (and F) form only one bond. Therefore they can only be terminal atoms in a structure.

• So you can not have

• C---H---C

• It has to be H---C--C

• Examples

• CH4

• PCl3

• SO32-

• NO3-

• CN-

• COBr2 (C is bonded to O and Br atoms)

• SO2

• H3O+ (hydronium ion

• N3-

Draw Lewis structure of CO2 i) Valence electrons: 4 + 2 x 6 = 16 ( 8 pairs)ii) Central atom C; O -- C -- Oiii) Give octet to carbon -- O -- C -- O -- Try to fill octet to O iv) Count electrons:4 bond pairs = 4 pairs 4 lone pairs = 4 pairs 8 electron pairs

Multiple bonds

• In general a triple bond (N2) is ________ than a double bond (O2) which is ________than a single bond (F2).

• Bond order: BO of 1--single bond, BO of 2-- -double bond, BO of 3 --triple bond.

• The stronger the bond,

Terminology used in describing Lewis structures of moleculesBond pairs: An electron pair shared by two atoms in a bond. Lone pair: An electron pair found solely on a single atom. Single covalent bond - Bond between two atoms when they shared 1 pair Double covalent bond – Bond between two atoms when they shared 2 pairs.Triple covalent bond – Bond between two atoms when they shared 3 pairs.Lewis Structure, Stability, Multiple Bonds, and Bond Energies Bond order

The stability of a covalent compound is related to the bond energy. The magnitude of the bond energy increases and the bond length decreases in the order: single bond > double bond > triple bond.Bond Energy order: single < double < triple

Bond length order: single (1) < double (2) < triple (3)

Resonance

• Resonance structure –1 of 2 or more Lewis structures for a molecule (ion) that can’t be represented with a single structure

• Resonance – use of

• Each resonance structure contributes to the actual structure– no single structure is a complete description– positions of atoms must be the same in each,

only electrons are moved around– actual structure is an “average”

• Draw resonance structures for SO3 and N3-.

Exceptions to Octet Rule

There are three classes of exceptions to the octet rule.

1) Molecules with an odd number of electrons;

2) Molecules in which one atom has less than an octet;

3) Molecules in which one atom has more than an octet.

Let’s do Lewis structures for

• CO2 (CS2)

• O3 (SO2)

• I3-

3D structure of species

• Electrostatic forces in ionic bonds is _____________. But species with covalent bonds have electron pairs concentrated btn 2 atoms and is ..

• We use VESPR theory to predict the shape of the covalently bound species.

VSEPR theory

VSEPR

• Most stable geometry is one in which electron pairs (electron clouds) are as

Shapes of molecules (3D)

• The geometry is determined by the atoms present in the species. See atoms that are bonded to other atoms. Don’t “see” lone pairs but they influence geometry

• I. Diatomics (2 atoms only): always ________

• H2, HCl, CO X----X

• II. Polyatomic (3 or more atoms) species:

Use VSEPR model to predict shapes

Steps in applying VSEPR

• 1. Do Lewis structure

• 2. Count total e- pairs (clouds) around central atom (A). Multiple bonds count as one electron pair (cloud). In reality multiple bonds are bigger than single bonds (electron clouds larger).

• 3. Separate e- pairs into bonded pairs (B) and lone pairs (E)

• 4. Apply table that I give you.

• 5. Remember that lone pairs of e-’s are invisible, but their presence affects the final molecular geometry!!!!!

• Lone e- pair-lone e-pairs are more repulsive than bonded pair-lone pair repulsions or bonded pair-bonded pair repulsions.

VSEPR: valence shell electron pair repulsion

• 2 electron clouds around a central atom (A)

2 electron clouds

Three electron clouds

Three electron clouds

Four electron clouds

Table 4.5 (changed)

• # e # bonded #lone pairs geom angle clouds pairs pairs

• 2 • 3 • 3• 4• 4• 4

Predict geometry

• H2S

• SO2

• CO2

• CF4

• H2CO

• ClO3-

• ClO2-

Polar vs nonpolar cmpds

• A molecule is polar if its centers of positive and negative charges do not coincide. If a molecule is polar we say that it acts as a dipole. In an electric field nonpolar molecules (positive and negative centers coincide) do not align with the field but polar molecules do.

• Next we will see why this happens and the implications.

Molecules are subjected to electric fieldPolar molecules align with fieldNonpolar molecules are not affected

Polar molecules

• I. Diatomics, A-B

• a.If A = B have homonuclear diatomic; has

• b. A ≠ B have heteronuclear diatomic

II. Polyatomic species are more complicated.

• Let’s look at VSEPR cases considered.

• General rule (my rule):

Which of these are polar?

• H2S

• SO2

• CO2

• CF4

• AlCl3

• CHCl3

• SCl2

Properties based on electronic structure and molecular geometry• Intramolecular forces: within a molecule--

bonds

• Intermolecular forces: between molecules--these determine important properties as melting and boiling points and solubility

Solubility

• Like dissolves like:

• Polar cmpds dissolve in polar solvents as ionic and polar cmpds (HCl)

in water

• Nonpolar cmpds dissolve in nonpolar solvents: oils in CCl4

Melting and boiling points

• Stronger the intermolecular forces the higher the melting and boiling points

• In general for cmpds of similar weight: polar moleculaes have stonger forces than nonpolar cmpds

• In general for similar structure the greater the mass the stronger the forces

Which have higher melting (boiling pts)

• CO and NO

• F2 and Br2

• CH3CH2OH and CH3CH3