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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
2
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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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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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 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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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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