Chemistry: The Molecular Science Moore, Stanitski and Jurs

Chapter 3: Chemical Compounds

Homework Chapt. 3,12,28,29,75,76

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• Contain 2 or more elements. • Nanoscale: a discrete molecule. • Form when non-metals combine.

Molecular formula • shows the number and kind of elements used

e.g. water H2O ammonia NH3

benzene C6H6

Molecular Compounds

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Molecular Ions

• Collections of atoms, but with overall charge:

e.g.: [CO3]2- the -2 charge applied to the whole anion:

Usually written CO3

2-

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Lecture Outline: I. Molecules and Forces: Inter vs. Intra molecular

II. Empirical and Molecular Formula, problem

solving

III. Hydrocarbons (Example of Covalent Compounds)

IV. More on ionic compounds, nomenclature, Oxidation States

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Molecules and Forces: Types of molecules: A. Small, Discrete– e.g. H2, HCl, NH3, B10C2H12

B. Large, discrete: , e.g., proteins

C. Extended NaCl, Si,

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CH4 B10C2H12 (Ortho-carborane)

Discrete Molecules

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NaCl

Na+ Cl-

Extended Systems: Both ionic and covalent extended systems exist

Ionic- system is a collection of ions

Covalent—atoms close to neutral

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_ +

graphene

INTRA MOLECULAR FORCES—HOLD ATOMS TOGETHER IN THE MOLECULE: 2 TYPES: Covalent, Ionic INTERMOLECULAR FORCES—Forces between molecules (Van der Waal’s –long range, much weaker) talk about these later

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Covalent Forces: Sharing of (Valence) electrons between nuclei:

+

e-

+

e-

H H

+

e-

+

e-

RO

Sharing of (valence) electrons between nuclei.

H2

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www.chemicool.com/definition/morse_potential.html

Atoms in molecules adopt an equilibrium internuclear distance (R0)

R > R0, electron overlap draws nuclei together

R< R0, nucleus-nucleus repulsion pushes nuclei apart

Atoms vibrate about this equilibrium distance 10

Polar molecules are covalent, but the sharing is not equal!

+

e-

+

H, 1 proton, 1 electron Cl, 17 protons, 17 electrons, (7 of them valence electrons)

e-

+ + δ+ δ- This end of the molecule has a partial positive charge

This end of the molecule has a partial negative charge

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Ions - charged units - formed by transfer of e- between elements.

Cation - positive ion. Metals form cations Na Na+ + e-

Anion - negative ion. Nonmetals form anions S + 2 e- S2-

11 e- → 10 e-

Ne has 10 e-

16 e- → 18 e-

Ar has 18 e-

Main group elements: • Transfer e- to achieve the nearest noble gas arrangement.

Number of e- transferred = group A# or (8 – group A#).

Ions and Ionic Compounds

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Properties of Ionic Compounds Ionic compounds • Not individual molecules. Crystal lattices

Each ion is surrounded by many others

Formula unit = smallest ratio of anions to cations

NaCl sodium chloride

Properties of Ionic Compounds

External force displaces layers

Repulsion occurs

Na+

Cl-

Ionic crystals can be cleaved:

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Electrostatic forces hold ionic compounds together:

F = k Q1Q2

d2

Properties of Ionic Compounds

High melting points strong forces. high charge = high m.p.

ions m. p. (°C) NaF +1 -1 993 CaO +2 -2 2572

Similar sized ions:

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Properties of Ionic Compounds Ionic compounds are electrical insulators when SOLID

Many are soluble in water

will conduct if molten

Ionic Compounds

Ions Compound Charges

Charges are always balanced.

Ionic compounds are always neutral!

Mg2+ and F- MgF2 (2+) + 2(1-) = 0 Mg2+ and SO4

2- MgSO4 (2+) + (2-) = 0 Mg2+ and PO4

3- Mg3(PO4)2 3(2+) + 2(3-) = 0

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Ionic Compounds: Electrolytes If an ionic compound dissolves in water:

• It dissociates breaks apart into its ions.

• It is a strong electrolyte the solution is a good electrical

conductor.

Molecular & Ionic Compounds Property Formation

Molecular Most are non-metal combinations

Ionic Metal/non-metal combinations

Physical state

mp & bp

Conductivity

Solubility

Gases, liquids & solids. Brittle & weak or soft & waxy

Low

Poor heat & electrical conductors

Few soluble in water

Crystalline solids Hard & brittle

High

Poor heat & electrical. Good electrical if molten

Many soluble in water

In solution Remain molecular Dissociate 19

• Do not contain C or (C and H). Inorganic compounds

Organic compounds • always contain C, usually H • may contain many other elements.

e.g. benzene C6H6 ethanol C2H6O

water H2O ammonia NH3

carbon dioxide CO2

• most (but not all) are molecular.

Molecular Compounds

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Ethanol has the formula C2H6O … • Doesn’t show atom connections. • A structural formula does.

C2H6O may not be ethanol. • Two C2H6O structural formulas:

ethanol

H – C – C – O – H

H |

| H

H |

| H

dimethyl ether

H – C – O – C – H

H |

| H

H |

| H

Molecular Formulas

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More elaborate models:

Ball-and-stick model

Space-filling model

Molecular Formulas

C

H O

Empirical vs. Molecular Formula: Ethylene (C2H4) Molecular Formula: The number and types of atoms in the molecule: Example: C2H4

C = C H

H H

H

Empirical Formula, lowest ratio of elements to each other: Example: CH2

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hydrogen peroxide H2O2 HO

borane (boron trihydride) BH3 BH3

octene C8H16 CH2

butene C4H8 CH2

Examples Compound mol. formula emp. formula

diborane (diboron hexahydride) B2H6 BH3

Empirical and Molecular Formulas

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Moles of Compounds A mole of XmYn contains:

m moles of atom X and n moles of atom Y

1 mol of H2O contains: 2 mol of H atoms and 1 mol of O atoms

Molar mass = sum of the atomic masses

Mass of 1 water molecule: = 2(1.008 amu) + 1(15.999 amu) = 18.015 amu

Molar mass of water: = 2(1.008 g/mol) + 1(15.999 g/mol) = 18.015 g/mol

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Ionic compounds do not contain molecules.

Don’t use “molecular weight” to describe mass.

atomic wts Formula wt. Molar mass Compound amu amu g/mol NaCl 22.99 + 35.45 58.44 58.44 Ca(NO3)2 40.08+2(14.01)+6(16.00) 164.10 164.10

Formula weight is the correct name • Molar mass can be used

Molar Mass of Ionic Compounds

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Percent Composition Two names used:

• percent composition by mass, or • mass percent of the compound.

Example What is the mass percent of each element in sodium chlorite, NaClO2?

molar mass = (22.990 g) + (35.453 g) + 2(15.999 g) = 90.441 g

%Na = mass of Na in 1 mol NaClO2

mass of NaClO2 in 1 mol NaClO2 x 100 %

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= 22.990 g 90.441 g

x 100 %

% Cl = 100 − 25.42 − 35.38 = 39.20%

%Na = mass of Na …

mass of NaClO2 …

x 100 %

= 25.42%

= 2(15.999 g) 90.441 g

x 100 % = 35.38%

%O = mass of O …

mass of NaClO2 …

x 100 %

Percent Composition

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Last example: molecular formula percent composition

Empirical formula = the simplest ratio of atoms in a compound.

Not molecular formula

The process can be reversed: percent composition empirical formula

Empirical and Molecular Formulas

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Example An orange compound was found to be 26.6% K, 35.4% Cr and 38.0% O. Determine its empirical formula. • Assume a 100.0 g sample. % becomes mass in grams

• Divide each mass by its atomic mass. Gives the number of moles of each (in 100 g).

• Divide each by the smallest answer found. The smallest whole number ratio is the empirical formula.

Empirical and Molecular Formulas

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Unknown: 26.6% K 35.4% Cr 38.0% O

38.0 g O = 2.375 mol O 1 mol O 16.00 g O

35.4 g Cr = 0.6808 mol Cr 1 mol Cr 52.00 g Cr

In 100.0 g

= 0.6803 mol K 1 mol K 39.10 g K

Empirical and Molecular Formulas

26.6 g K

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Choose a multiplier to make integer

0.6803 mol 0.6808 mol 2.375 mol

Empirical formula = smallest integer ratio.

x2 2

x2 2

x2 7

The empirical formula is K2Cr2O7

Divide every number by the smallest value (ratios stay the same!)

K Cr O

0.6803 mol 0.6803 mol 0.6803 mol

= 1.000 = 1.001 = 3.491

Empirical and Molecular Formulas

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The molecular formula can be determined if the molecular mass is known.

Example Vitamin C has the empirical formula C3H4O3 and molecular mass = 175 g/mol.

Empirical mass: 3(12.01) + 4(1.008) + 3(15.99) = 88.03 g/mol

Empirical mass ≈ ½(molecular mass) Mol. formula = 2(emp. formula) = C6H8O6

Empirical and Molecular Formulas

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Hydrocarbons Binary molecules (C and H).

Alkane Hydrocarbon with CC single bonds only.

• Use an –ane ending.

• Have formula CnH2n+2 (unless cyclic).

Butane, C4H10

Hydrocarbons

H C H H C C H H C C C H H C C C C H

H

H

H H H H H H

H H

H H H

H H H H H H H

methane ethane propane butane

Boiling points -161.6 °C -88.6 °C -42.1 °C -0.5 °C

(Intermolecular forces determine melting, freezing points!) 35

# of C 1 2 3 4 5 6 7 8 prefix meth eth prop but pent hex hept oct

Rings use a “cyclo-” prefix:

Example

4 carbon alkane ring = cyclobutane All C-C bonds are single.

Hydrocarbons

Similar to molecular compound prefixes

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Line-angle structures • Lines represent C-C bonds. • Each junction and end is a C • Each C needs 4 bonds. • C–H bonds are omitted C with 3H

(missing 3 bonds)

C with 2 H (missing 2 bonds)

ethane becomes | H

―C―C― H H

H |

| H

H |

propane becomes | H

―C―C―C― H H

H |

| H

H |

| H

H |

Hydrocarbons

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Branched alkanes occur.

Isomer: same formula, different atom arrangement.

isomers

butane

C4H10

| H

―C―C―C―C― H H

H |

| H

H |

| H

H |

| H

H |

Alkanes and Their Isomers

methylpropane

C4H10

Ι H

―C―C―C― H H

H |

H C H

| H

H |

H |

Ι H

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Ι H

―C―C―C― H H

H |

H C H

| H

H |

H |

Ι H

methylpropane

Alkyl functional groups • An alkane with a H atom removed. • Named by replacing “-ane” with “-yl”

a methyl group

-CH3 methyl -CH2CH3 ethyl -CH2CH2CH3 propyl etc.

Alkanes and Their Isomers

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Alkanes and Their Isomers

• Find and name the longest continuous alkane. • Name the alkyl “branches”. • Add numbers to show the branch locations on the

base chain. Use smallest possible numbers.

• Use prefixes of mono-, di-, tri, etc. as necessary.

Naming branched alkanes

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Alkanes and Their Isomers CH3 CH3

CH3 C CH2 CH CH3 CH3

2,2,4-trimethylpentane

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Alkanes and Their Isomers Formula Isomers Formula Isomers

CH4 1 C9H20 35

C2H6 1 C10H22 75

C3H8 1 C12H26 355

C4H10 2 C15H32 4.4 x 103

C5H12 3 C20H42 3.7 x 105

C6H14 5 C30H62 4.1 x 109

C7H16 9 C40H82 6.3 x 1013

C8H18 18

C6H14

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Transition metals: – lose varying number of e-.

Ti Ti2+ or Ti+4 (grp 4B) Cr Cr2+ Cr3+ or Cr+4 (grp 6B) Fe Fe2+ or Fe3+ (grp 8B) Cu Cu+ or Cu2+ (grp 1B) Mn Mn2+ Mn5+ or Mn7+ (grp 7B)

Monatomic Ions

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Monatomic Ions Common monatomic ions

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Naming Ionic Compounds

NaCl

Name the ions and add together…

sodium chloride

… cation then anion (drop “ion” from both)

MgCO3 magnesium carbonate

KHSO4 potassium hydrogen sulfate

strontium oxide

magnesium hydroxide

SrO

Mg(OH)2

Single charge metal-ion examples

KMnO4 potassium permanganate 45

Naming Ionic Compounds

FeCl2 FeCl3

iron(II) chloride

Cu2O

CuO

iron(III) chloride

copper(I) oxide

copper(II) oxide

Multiple charge examples

Cu2O CuO

Different charges on metal ions; Different oxide formula:

Ti+4 TiO2 Cr+4 CrO2 Cr+3 Cr2O3 Fe+2 FeO Fe+3 Fe2O3

Note: the “O” ion always has a (-2) charge: The Compound is neutral (charges balance)

A few problems with this simple picture: e.g., magnetite (Fe3O4) Fe3O4 Average Charge on each Fe = 8/3 = 2.67 (1/3 of Fe sites are +2: 2/3 of Fe sites are +3) There are a number of such examples, so the picture can be a little more complicated than in the book, but not much.

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Multiple atom “units” with a net electrical charge.

NH4+ ammonium ion

NO3- nitrate ion

SO42- sulfate ion

OH- hydroxide ion

CN- cyanide ion

HSO4- hydrogen sulfate ion

Polyatomic Ions

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Incr

ease

O SO4

2- sulfate ion SO3

2- sulfite ion

NO3- nitrate ion

NO2

- nitrite ion

Two forms exist: –ate and –ite endings used. More oxygen = “-ate” Less oxygen = “-ite”

If they contain H, add a prefix “hydrogen”

HSO4- hydrogen sulfate ion

(common name=bisulfate ion) HCO3

- hydrogen carbonate ion (common name=bicarbonate ion)

Oxoanions

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Oxidation States: The Rules! 1. Assign oxidation numbers to elements in a compound: 2. Sum of (oxidation number x # atoms of that element) = 0 neutral system = overall charge of molecular anion 3. Pure elements oxidation number is zero: e.g., O2 4. In compound, oxidation # of O = -2 Oxidation number of H = +1 (except in metallic hydrides) 5. Halogens (Cl, F, Br, I…) -1 except when bonded to oxygen

NH3 oxidation # of H = +1; therefore, oxidation # of N = -3 (3x(+1) -3) = 0 SO2 oxidation # of O = -2; therefore, oxidation # of S = +4 (2x(-2) + 4 = 0 CO3

2- oxidation # of O = -2; oxidation # of C = +4 (3x(-2) +4 = -2)

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Oxidation States in Oxides Al2O3 Al+3 aka Al(III) CoO Co+2 aka Co(II) SrTiO3 Sr+2 , Ti+4

Oxidation States in other compounds FeS (S = -2, Fe = +2) FeS2 (S = -2, Fe= +4) CoCl2 (Cl = -1; Co = +2) ClO4

1- (O = -2 Cl =?)

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The ability of transition metals to cycle between oxidation states is vital to catalysis.

CO CO2

C+2 C+4

RuO2 RuO: Ru+4 Ru+2

www.iap.tuwien.ac.at/.../stm_gallery/nonmetals

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