Chemistry I
PD Dr. Christoph Wutz
christoph.wutz@chemie.unichristoph.wutz@chemie.uni-hamburg.de
http://www.chemie.unihttp://www.chemie.uni-hamburg.de/tmc/wutz/
hamburg.de/tmc/wutz/
Page 1
Contents
1. Introduction
Books, Internet, History, Basic Concepts
2. General Chemistry
2.1 Atomic Structure
2.2 Periodic System
2.3 Chemical Bond, Molecules
3. Physical Chemistry
3.1 State of Matter, Phase
Transitions, Separation Proc.
3.2 Chemical Reaction
Thermodynamics, Kinetics
4. Inorganic Chemistry
4.1 Acids/Bases, pHpH-value
4.2 Salts, Solubility
4.3 Redox reactions
4.4 Electrochemistry
Page 2
1
1. Introduction
classic
€ 13.40
ISBN:
basic
advanced
336 Pages; € 25.1397 Pages; € 53.25.53.ISBN: 978978-0-1919-927789927789-6 ISBN: 978978-0-1919-927789927789-6
978978-0486656229
Page 3
OnOn-line
Introduction to Chemistry (Mark Bishop)
http://preparatorychemistry.com/
On-line textook (R. Dickerson; CalTech)
http://caltechbook.library.caltech.edu/178/
On-line lectures and notes (MIT):
http://ocw.mit.edu/courses/chemistry/
Test your knowledge on:
http://www.chemistry-drills.com/
Page 4
2
History I:
Unintentional Chemistry
3000 BC.: Smelting of lead, copper and tin (bronze)
1. Roasting of metal ore at about 300°
300°C :
Malachite: CuCO3
2. Reduction at >1000°
>1000°C: 2 CuO + C
metal oxide + charcoal
→ metal + carbon dioxide
CuO + CO2
2 Cu + CO2
Page 5
History II:
Alchemy
Intentional, traditional, nonnon-scientific chemistry;
Protoscience:
Protoscience: Precursor of modern chemistry:
• Creation of the "philosopher's stone," (catalyst?)⇒
• Conversion of base metals into solver or gold
• Search for the universal solvent "alkahest"
Notion of Greece
elements: Fire (hot)
Water (wet)
Earth (dry)
Air (cold)
Spirit/Ether
Buddhism
Fire
Water
Earth
Wind
Space
China (Tao)
Fire
Water
Earth
Wood
Metal
Some Discoveries/Inventions:
Discoveries/Inventions:
•Sal ammoniac (NH4Cl), cinnabar (HgS
):
(HgS):
by Abu Musa Jabir ibn Hayyan,
Hayyan, Arabia 8. Jh.
Jh.
•Alcohol (ethanol C2H5OH), Italy ~1100 from wine
•Medicals; mercury, arsenic, antimony comp:
by Paracelsus, 16. Jh.
Jh.
•Phosphorous: Hennig Brand; Hamburg 1669
•Porcelain: Johann Friedrich Böttger 1707/08
Discovery of Phosphor
by Hennig Brand (1669)Page 6
3
History III
Chemistry as Natural Science
After 17. Jh.:
Jh.:
Scientific chemistry:
Exemption from dogma and religion.
Rational conclusions
based of observation
and experiments.
Robert Boyle (1627(1627-1691)
Irish researcher
"The Sceptical Chymist"
Chymist" (1661)
Joseph Priestley
(1733(1733-1804)
Engl. researcher
Invent.: NH3,
N2O, CO, SO2
Carl W. Scheele
(1742(1742-1786)
Swed.
Swed. Pharm.
O2, Cl2, Ba, Mn;
HCN, lactic acid
Antoine Lavoisier
(1743(1743-1794)
French Chem.
Oxidation
J. GayGay-Lussac
(1778(1778-1850)
French Chem.
Ideal gas law
Justus v. Liebig
(1803(1803-1873)
German Chem.
Chem.Analysis
Page 7
Difference ChemistryChemistry-Physics
Chemistry:
Chemistry: Change of ComposiComposiPhysics:
Physics: (Change of) State of
tion
of
Matter
Matter
Example:
Iron: metallic, shiny, conductive, malleable.
Heating iron ⇒ melting
Iron in humid air ⇒ gets rusty
cooling ⇒ solidification.
Rust: redred-brown, nonnon-conductive
Properties unchanged.
not moldable
Reversible change of state
⇒ Change of matter
Another example:
Platinum wire heating to redHeating magnesium⇒
red-hot
magnesium⇒
Cooling: No change in prop.
combustion with light emission⇒
emission⇒
Emission of light: physical proc.
magnesium oxide (white powder)
⇒ change of matter
Last Example:
Heavy cooling of air ⇒ liquefaction Respired oxygen converts
Slow warming up (distillation) ⇒
carbohydrates into CO2+H2O.
Separation of oxygen and nitrogen Energy is released.
Properties unchanged.
⇒ Chemical process.
Liquefaction, evaporation, melting,
distillation: physical processes
Combustion, oxidation,
chem. synthesis: Chemical proc.Page 8
4
Beginning and Resources of
Chemical Industry
before 1820: Production of:
• Soda (sodium carbonate) for processing of
glass, bleaching agents, detergents, etc.
• Synthetic fertilizers
• Dyes
Resources of chemical industry before 1870:
• Coal
• Minerals
• Plant products
• Animal products
Resources of chemical
industry after 1870 - ?
• Crude oil
Page 9
Chemical Hazards
21. Sept. 1921; BASF, Oppau,
Oppau, : Explosion of fertilizer plant:
561 Dead, 2000 Injured, 900 apartments destroyed, 100 m crater
2 NH4NO3 → 4 H2O + 2 N2 + O2
Page 10
5
Population
Benefit from Chemical Industry
Demographics
Development of chemical industry
• Synth. fertilizer
• Pesticides
• Medicals
• Hygiene items
Page 11
Why bother with Chemistry?
Chemistry affects our life's:
Carbohydrates,
Proteins,
Fats,
Vitamins, etc.
Toy
Sports
equipequipment
Mobiles
Coal/Gas
Petrol
Batteries
Medicine
Glasses
Medical devices
Detergents
Soap
Tooth paste
Dyes
Carpets
Furniture
Page 12
6
Significance of Chemistry for the Living
from atom
via molecules
(>15 mio.!)
mio.!)
to organism
Ascorbic Acid
Chlorophyll
Page 13
Analytical/Synthetic Chemistry
Analytical chemistry:
Understanding nature:
What is it made of?
What is in there?
and how much?
Synthetic Chemistry:
Can we produce the same products
(or even better) in the lab?
Wikipedia:
Science of matter and
changes it undergoes.
Page 14
7
Basic Terms and Concepts
Measures and Units (SI
(SI))
Measure
Length
Mass
Volume
Density
Temperature
Force
Pressure
Energy
Amount of substance
Molar Mass
Molar concentration
Molar Enthalpy
Molar Entropy
Symbol
l
m
V
ρ
T
F
p
E
n
M
c
H
S
Unit
Meter
Gramm
Liter
Gramm per cm3
Kelvin
Newton
Pascal
Joule
Mole
Mass per Mole; M = m/n
Mole per volume; c=n/V
Joule per Mol
Symbol
m
g
L
g/cm3
K
N = kg·m/s2
Pa = N/m2
J = kg·m2/s2
mol
g/mol
mol/l
J/mol
J/mol·K
Page 15
General Structure of Matter
Chemical Compound
solid
Substance
Atomic level
Na+
Cl-
Salt
liquid
Sodium chloride
gaseous
Water
Hydrogen
atom
Water molecules
Oxygen
atom
Air
Oxygen a. nitrogen
molecules (mixture)
nitrogen atom
Page 16
8
Mixtures, Substances, Compounds,
Elements
Mixture
e. g. vodka
water/alcohol
separation by
physical methods
e.g. distillation
filtration, extraction
Chemical
compounds
Substance e. g. water
or pure
alcohol
are
either
or
Decomposition
by chem. methods
Elements
one type oft atom;
no decomposition
by chem. methods
composed of two or more different
chemical elements (types of atoms)
connected by chemical bonds
e. g. water H2O
two atoms of hydrogen (H)
one atom of oxigen (O)
e. g. hydrogen (H)
oxygen (O)
iron (Fe)
Page 17
Symbols of Chemical Element
Symbols used for abbreviation of chemical
elements by Berzelius; today by IUPAC:
IUPAC:
OneOne- or twotwo-letters long, only first capitalized.
Mostly derived from Greek or Latin words.
H
He
Li
Be
B
C
N
O
F
Ne
Hydrogen
Helium
Lithium
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
Na
Mg
Al
Si
P
S
Cl
Ar
K
Ca
Sodium
Magnesium
Aluminum
Silicon
Phosphorus
Sulfur
Chlorine
Argon
Potassium
Calcium
Jöns Jakob Berzelius
Swedish chemist
(1779-1848)
First 20 elements
Page 18
9
Natural Abundance of Elements
Over 90% of the universe is hydrogen (H);
second most is helium (He).
Earth inner core
Earth crust(0-40 km)
3%
3%
1% 1%
2% 2%
2%
2%
4%
5%
5%
7%
Fe
8%
Al
O
49%
Si
Fe
80%
26%
O
Si
Al
Fe
Ca
Na
K
Mg
H
Res.
Common compounds:
SiO2, Silicates, AlAl-, iron oxides,
carbonates, sulfates, hydroxides
Fe
Si
Ni
O
S
Res.
Earth inner core (3000 km)
mainly composed of iron
(2900°
(2900°C, liquid). Fe/NiFe/Ni-convecconvection at mantle ⇒magnetism Page 19
Change of Matter
by Chemical Reaction
Iron
(Fe)
Sulfur
(S)
Chemical reaction
IronIron-SulfurSulfurMixture: Fe + S
How can it be
separated?
IronIron-sulfursulfurcompound
Iron sulfide (FeS
(FeS))
No physical
separation into
iron and sulfur
possible
Page 20
10
Chemical Reaction
A chemical reaction is a process that leads to the transformation
transformation
of one set of chemical substances (reactants; reagents)
to another (products).
Chemical equations used to graphically illustrate chemical reactions:
reactions:
Reactant 1 (+ reactant 2)
product 1 + (product 2)
A chemical reaction alters the combination of atoms in substances,
substances,
the number of atoms of each elements remains constant.
"The mass of matter remains constant
during chemical reactions".
(Law of Conservation of Mass)
Antoine Lavoisier
fr. Chem. (1743-1794)
Page 21
Chemical Reaction
Example: Sodium Chloride
Example:
Chlorine
Sodium (Metal)
Salt (NaCl)
+
Chemical reactions involve a "rearrangement" of atoms.
New compounds are formed with different properties.
Which substance can react with which other substance
yielding which compound in which molar ratio?
Page 22
11
Elemental Particles
Matter composed of atoms, composed of three elemental particles:
+
0
-
Proton (p+): "big", "heavy" (relatively!), positive
mass ~ 1 u
Neutron (n): "big", "heavy" (relatively!), neutral
mass ~ 1 u
Electron (e-): tiny, lightweight, negative
mass ~ 1/1800 u
u extremely small unit: 1 u = 1,66⋅
1,66⋅10-27 Kg
Page 23
Radiation
Radiation is a process in which waves or particles travel through
a medium or space ⇒ Transport of energy or mass.
Dualism: Each radiation has wave and particle character.
Electromagnetic waves (light, micromicro-, radioradio-, xx-ray)
Energy ~ frequency ~1/wave length
Radioactive decay:
decay:
Atomic nucleus emits particle or wave
Name
Character
α-radiation He nucleus:
2 P+; 2 N
β-radiation Electron
γ-radiation
electromag.
electromag.
wave
Mass
Charge
4u
+2
1/1823 u
-1
-
Page 24
12
2. General Chemistry
Atomic Structure: Early Atomic Models
Stoichiometry
Law of defined proportions = const. composition (J. Proust, 1794):
1794):
"A chemical compound always contains exactly the same proporproportion of elements by mass". E. g. water: oxygen:hydrogen = 8:1
Law of multiple proportions (J. Dalton, 1803):
"The proportions of mass of two elements in different compounds
are rations of small whole numbers". E.g.:
100 g of carbon reacts with 133 g oxygen to carbon monoxide
100 g of carbon reacts with 266 g oxygen to carbon dioxide
266:133=2:1
Daltons Atomic Theory (180
8):
(1808):
• Elements are made of tiny particles called atoms.
•The atoms of a given element are different from those of any other
other element;
they can be distinguished by their respective relative atomic weights.
•All atoms of a given element are identical.
•Atoms of one element can combine with others to form chemical compounds;
compounds;
a given compound always has the same relative numbers of types of atoms.
•Atoms cannot be created, divided into smaller particles, nor destroyed;
destroyed;
a chemical reaction simply changes the way atoms are grouped together.
together.
Page 25
GeigerGeiger-Marsden Gold Foil Experiment
Rutherford Model
1910: α-particles (positive
(positive HeHe-nuclei) directed onto very
thin gold foil: Very few particles deflected ⇒
1. Atoms are almost "empty"
2. Atomic mass concentrated in positive nucleus
3. Atomic volume represented by negative shell
(Radium)
ZnS
Sir E. Rutherford
New Zealand Chem.
Rutherford
Model
Page 26
13
Rutherford Atomic Model
Atoms have a positive core (nucleus),
(nucleus),
surrounded by negative electron cloud (shell).
(shell).
Protons and neutrons located
in nucleus (⇒
(⇒nucleons).
nucleons).
Neutrons: no charge,
Protons: charge = +1 ⇒
Number of protons = number of
(pos.) charges in nucleus
= Atomic number (Z) in
periodic system of elements (PSE)
The number of protons determines the element
Z=3
Lithium
Atomic diameter ~ 10-10 m, nucleus only 1/100000 of it:
Nucleus: tiny, "heavy", positive
Shell: "bigger", light, negative
negative Page 27
Atomic Mass
Number of nucleons (protons + neutrons)= atomic mass (ma),
Mass of electrons can be neglected.
The number of protons determines, which element (type of atom)
(= atomic number Z)
The number of neutrons N plus protons Z determines,
which mass is has. (atomic mass ma = Z + N)
Example: Lithium
3 protons ⇒ atomic number Z = 3
4 neutrons ⇒ atomic mass = 7
Proton
Neutron
Electron
Page 28
14
Molecular Mass
The mass of molecule mM of a chemical compound is the
sum of the atomic masses of the constituting elements
Example: Glucose C6H12O6
The empirical formula indicates the number of atoms
of each element in the molecule
mM = 6·
6·ma(C) + 12·
12·ma(H) + 6·
6·ma(O)
ma(C)
(C) = 12 u
ma(H)
(H) = 1 u
ma(O)
(O) = 16 u
mM (Glucose) = 6·
6·12u + 12·
12·1u + 6·
6·16u = 180 u
Page 29
Amount of Substance, Mole
The mass unit u measures extremely small quantities (single molecules):
molecules):
The Amount of Substance n (or Chemical Amount) with unit Mole is
a quantity that measures the size of an ensemble of elementary entities
entities
(atoms, molecules other particles) on lablab-scale.
The Amount of Substance is no mass, nor a number of entities
- it is, however, closely related to both:
The Amount
of Substance n is the mass m of a portion related to the
m
n = Mass M of the substance. The Molar Mass M is a substanceMolar
substanceM quantity with the unit [g/mol].
specific
Examples:
Atomic/molecular mass
Li = 7 u
H2O = 18 u
H2 = 2 u
C6H12O6 = 180 u
Molar Mass M
M(Li)
M(Li) = 7 g/mol
M(H2O) = 18 g/mol
M(H2) = 2 g/mol
M(C6H12O6) = 180 g/mol
One Mole Li weights 7 g, 1 Mole water ≈ 18 g; 2 Mol H2O ≈ 36 g.
Page 30
15
Avogadro´
Avogadro´s Number
Illustration of the Mole
An amount of substance n = 1 mol contains always
the same number of particles: Avogadro´
Avogadro´s Number NA:
1 mol ≈ NA = 6·
6·1023 particles
1 mol lithium ≈ 6·1023 Li atoms
1 mol hydrogen ≈ 6·1023 H2 molecules = 12·
12·1023 H atoms
23
1 mol helium ≈ 6·10 He atoms
Representation: 1 particles ⇒ 1023 particles
1 mol lithium
6·1023 atoms
(7 g)
1 mol hydrogen
6·1023 molecules
(2 g)
1 mol water
6·1023 molecules
(2+16 = 18 g)
4 mol carbon
4·6·1023 atoms
(48 g)
1 mol carbon
6·1023 atoms
(12 g)
Page 31
Calculating with the Amount of Substance
The mass m [g] and the amount of substance n [mol)]
can be converted by the molar mass M [g/mol]:
M=
m
n
m = M ⋅n
Examples:
What is the mass of 0,2 mol glucose?
n=
m
M
m = 180 g / mol ⋅ 0,2mol = 36 g
What is 5,85 g Sodium chloride (NaCl) in Mol?
mA(Na)
(Na) = 23 u; mA(Cl)
(Cl) = 35,5 u ⇒ MNaCl= 58,5 g/mol
n=
5,85 g
= 0,1mol
58,5 g / mol
Page 32
16
Illustration of
Stoichiometry and Molar Mass
2 H2 + O2
2 diatomic
2 H2O
1 diatomic
hydrogen molecules oxygen molecules
2 water molecules
oxygen atom (O):
8 P+;8 N⇒ma=16 u
hydrogen atom (H):
1P+⇒ ma = 1u
mH2O = 16+2=18u
A water molecule (H2O) consists of 2 hydrogen and 1 oxygen atoms.
The reactants must contain twice as much H atoms as O atoms.
Both elements are diatomic molecules.
Molar Masses: M(H2)=2 g/mol; M(O2)=32 g/mol; M(H2O)=18 g/mol
Hydrogen +
2 molecules (2⋅
(2⋅2u = 4u )
2000 molecules
2 mol (2⋅
(2⋅6 ⋅1023
molecules)
4g
+
Oxygen =
Water
1 molecule (2⋅
(2⋅16u=32u) 2 molecule (2⋅
(2⋅18u=36 u)
1000 molecules
2000 molecules
1 mol (6⋅
2 mol (2⋅
(6⋅1023
(2⋅6⋅1023 molecules)
molecules)
32 g
=
36 g
Conservation of mass!
Page 33
Molar Volume of Gases
1 Mole of any gas contains always 6⋅
6⋅1023 particles,
either atoms (noble gases, e. g. He),
or
diatomic molecules (other gaseous elements: H2, N2, O2, F2, Cl2)
or
compound molecules (e. g. CO2).
1 Mole of ideal gas has a volume of V = 22,4 L under STP.
Molar volume of gases: Vm = 22,4 L/mol :
Example:
Example: A volume
V of 1,12 L of gas contains
which amount
V of=substance n [mol]?
m
n
V
1,12 L
V
n =⋅ n =
n == 0,05mol
V =V
m V
22,4 L mol Vm
m
Page 34
17
Concentration - Molarity
Many chemical reactions in solvent (e. g. water)
How many reactant particles are in the solution?
n
c=
V
The molar concentration c defined as amount of a consticonstituent n divided by the volume of mixture (or solvent*) V
*Dissolving a substance in a solvent does not change the
volume significantly. The density is increased.
Examples:
Examples:
2mol
2 Mole of a substance are dissolved c =
= 4mol / L
in 500 mL solvent. c = ?
0,5 L
Which amount of substance has to be dissolved in 200 ml,
yielding a solution ?
n = c ⋅V = 0,5mol / L ⋅ 0,2L = 0,1mol
n
2mol
2 Mole have to be dissolved in
= 20 L
how much solvent, so that c = 0,1 mol/L? V = =
c 0,1mol / L
1,8 g glucose in 100 ml water; c = ?.
n=
m
1,8 g
=
= 0,01mol = 10− 2 mol
M 180 g / mol
c=
0,01mol
= 0,1mol / L
0,1L
Page 35
Isotopes
Isotopes are variants of atoms of a particular chemical element,
differing in the numbers of neutrons and in the atomic mass.
They contain the same number of protons,
but a different number of neutrons.
Examples:
Examples:
2
1
Deuterium: 1 H
Hydrogen: 1 H
Carbon:
14
12
6
C
13
~99% nat. abundance; C ~1%
6
Traces of 6 C ; instable ⇒ radioactive
(Radiocarbon dating)
235
Uranium: 92 U fissile; 238U non fissile, both radioactive
radioactive
92
Isotopes have identical chemical properties (bonding),
but different physical properties (mass)
Page 36
18
NonNon-whole Number Masses
The atomic masses of many elements in the PSE are,
nonnon-whole numbers, since they consist of mixtures of isotope.
Example:
Example: Chlorine, average atomic mass = 35,45 u
consists of:
Calculation:
- 35Cl: 75.7%
0,757 ⋅ 35 + 0,24 ⋅ 37 = 35,45
- 37Cl: 24.2%
Elements that form only one stable type of atom (isotop
(isotop))
are named pure elements.
elements.
Example:
Example: 19Fluorine
Page 37
Ions
Atoms contain as many electrons in the shell as protons
in the nucleus and are consequently neutral.
Ions are atoms or molecules in which the number of electrons
is not equal to the total number of protons,
giving it a net positive or negative electrical charge.
Ions have completely different properties than the resp. atoms !
Examples: Na → Na+ + e- Positive ions = cations,,
cations
Cl + e → Cl
Negative ions = anions.
anions.
Cations derived from metal name (sodium (cat)ion
(cat)ion,, silver (cat)ion
(cat)ion))
or have the suffix –ium (NH4+ = ammonium ion)
Anions from nonnon-metal atoms have the suffix –ide (chloride)
in compounds with oxygen –ate (SO42-=sulfate) or –ite (SO32-=sulfite)
Ions can carry multiple charges:
Examples:
Examples: Al3+, O2Page 38
19
Review of Rutherford Model
Rutherford´
Rutherford´s model makes no statement about the electronelectron-structure of
the atom or the energy of the electrons in the shell; it does not
not explain
the differences in chemical behavior of the elements .
Experiment:
Experiment: Hydrogen absorbs specific
colors = wave length from the spectrum of
visible light
H2
Prism
Absorption spectrum of hydrogen
Vice versa, thermally activated hydrogen emits
light of exactly the same wave lengths.
Emissions spectrum of hydrogen
H2
Why specific wave length = energies?
Page 39
Discrete Energy Levels
of Electrons
Electrons can occupy diffedifferent wellwell-defined (=discrete)
energy levels in the shell.
Transition between level
by absorption or release
of energy
(net energy conservation)
Excitation: By additional energy
Excitation of atom by light
(heat, photon, electricity) the
electron can be moved to
higher level (outer orbit).
If photon energy too high
⇒ ionization
Deactivation of excited
electrons under emission of
electromagnetic wave and
return to lower level
(inner orbit).
Deactivation by emission of light
Page 40
20
Bohr Model
n=1
n=2
Nucl
n=3
Shell
n
elecelectrons
K
1
2
L
2
8
M
3
18
N
4
32
Electrons travel around nucleus in circular orbits
(shells) associated with definite energy levels
with the principle quantum number n.
The orbit radius r and the electron energy
increases as n2.
The shells are filled with electrons
from inside (lowest energy level)
to outside.
Each shell can carry a certain
number of electrons:
2 ·n 2
In the heaviest atoms
7 shells are occupied.
Only the electrons of the outermost (valence(valence-) shell
can participate in formation of chemical bonds.
Niels Bohr, Dan. phys.
Atomic model 1913
Nobel prize 1922 Page 41
Electron Configuration
Element
Z
H
He
Li
Be
F
Ne
Na
Mg
Cl
Ar
1
2
3
4
9
10
11
12
17
18
n=1
max. 2
1
2
2
2
2
2
2
2
2
2
n=2
n=3
max. 8 max 8+10
1
2
7
8
8
8
8
8
n=4
1
2
7
8
Complete (closed) shells The number of valence electrons
determines chemical properties!
Valence electrons
Page 42
21
Atomic Emission Spectrometry
Flame Photometry
Each element has its unique electron configuration
and hence characteristic electron transitions and corresponding
lines in the emission spectrum ⇒ atomic emission spectrometry.
spectrometry.
Slit
Sample
Eye
Prism
Optic
Light
source
Magnesium
Silicon
Sodium
Copper
Even without spectrometer some elements
can be distinguished by the color of the flame:
Na
Li
Cu
Page 43
Periodic Table (System) of Elements
Prediction of Elements by Mendeleev
Ga
Dimitri Mendeleev
Periodic table
(1869)
Properties of Gallium (Ga)
Prediction
Atomic mass ~ 68
Density
~ 5,9 g/cm3
Melting Temp. ~ 30°C
Oxide
X2O3
Chloride
XCl3
Reality
69,72
5,91 g/cm3
29,8°C
Ga2O3
GaCl3
Page 44
22
Periodic Table of Elements
Elements listed with increasing value of atomic number
= no. Protons in nucleus = no. electrons in shell.
Page 45
Structure of Periodic Table
Main groups
Main group number
= number of
valence electrons
I.
Periods
1.
2.
3.
4.
5.
6.
7.
Main groups
Groups
VIII.
III. IV. V. VI. VII.
II.
Auxiliary groups
La
Ac
Lanthanids
Actinids
Page 46
23
Metals,
Metals, Metaloids (Semimetals)
Semimetals),
NonNon-metals
•isolators
(exept carbon
as graphite)
•gaseous or
low melting
•brittle solids
•conductive
•solid
(except mercury)
some high melting
•malleable, ductile
Metallic character of elements in the periodic table
decreases from bottom left to top right.
> 80% of the elements are metals.
Page 47
Main Group Elements
I.
Auxilliary Groups
Alkali metals:
metals: silver, low melting, soft
Highly reactive,
reactive, natural abundance
only in compounds,
compounds, not elemental.
II. Alkali earth metals:
metals: like alkali, but less
reactive (stable in air).
III. Boron group:
group: B hard metalloid,
Al→
Al→Tl soft metals. AlAl-compounds
highly abundant in Earth´
Earth´s crust.
IV. Carbon group:
group: C nonmetal, Si,Ge metametaloids;
loids; Sn,
Sn, Pb met.; different properties.
Si high abundance (quartz); C→natural
compounds; Pb high density.
VII. Halogens:
Halogens: all nonmetals
V. Nitrogen group:
group: N=nonN=non-metall (gaseous)
F/Cl
F/Cl gas, Br liquid, I solid
P/As met. and nonmet.
nonmet. allotropes,
reactive; form salts with I and II
Sb/Bi=met.;
Sb/Bi=met.; N main ingredient air 78%
VIII. noble (inert) gases:
gases:
VI. Chalcogens:
Chalcogens: O gaseous, nonmetal,vital
extremely low reactivity
in water; 21% of air). S yellow solid
→ elemental state: as traces in air
Se, Te, Po metalloid/metal allotropes
Page 48
24
Orbitals
Flaw of Bohr´
Bohr´s model: A circulating charged particle emits radiation
→ energy loss; and: experiments indicate electron has wave character.
character.
⇒ Orbital model: Atomic orbitals are (energy) states
or wave forms of electrons in the atom (quantum mechanics)
The probability of finding an electron in a certain
volume of space is the electron density.
s-, pp-, dd-, and ff-orbitals exist:
s-orbitals are spherical.
The KK-shell (n=1; 2 e-) has only 1s1s-orb.
-)
In the 2. shell (Lshell,
max.
8
e
(L
at first the 2s2s-orbital (max. 2e-), afterwards
the pp-orbitals are occupied.
p-orbitals dumbbell shaped.
There are three: px, py, pz
having the same energy state.
Each contains max. 2 eEvery next shell ⇒ new type of orbital
(more complex shape). s, p, d hist., no meaning
Page 49
Energy Levels of Orbitals
Main energy levels ("shells"): 1, 2, 3, 4, 5, 6
Within each shell the energy level ss- < pp- < dd-orbital.
Electrons fill low energy orbitals before they fill higher energy
energy ones.
3. and 4. main energy level overlap ⇒
4s4s-orbitals are filled before the 3d3d-orbitals
(see PSE: After K, Ca (sblock),
introduction of transition metals with
(s
valence electrons in dorbitals.
Afterwards
Ga,Ge,As etc. (pd
(p-block)
Page 50
25
Quantum Numbers
Quantum numbers describe the energy of electrons in atoms.
Name
Symbol Orbital meaning
Range
of Values
n≥1
Value
Example
Principal (prim.) q.n.
q.n.
n
shell (main
energy level)
Azimuthal q.n.
q.n.
l
subshell:
subshell: s ⇒ 0 0 ≤ l ≥ n-1 for n = 3:
l = 0,1,2
(orbital) p ⇒ 1
(s,p,d)
s,p,d)
d⇒2
orientation of
-l ≤ ml ≥ +l for l = 1 (p(p-Orb.)
subshell shape
ml = -1,0,+1
px,py,pz
Spin of electron ½ : clockwise
-½: countercounter-clockwise
(angular momentum)
Magnetic q.n.
q.n.
(Projection of
ml
angular momentum)
Spin q.n.
q.n.
ms
n = 1,2,3,...
1,2,3,...
Page 51
Fitting Electrons into Orbitals
Pauli Principle/Hund
Principle/Hund´´s rule
Pauli principle: Two electrons my not occupy same quantum state⇒
state⇒
Each orbital occupied by max. two electrons with different spin ↑↓.
↑↓.
Hund´
Hund´s rule: When several orbitals of equal energy are available,
electrons enter singly with parallel spins (minimizes repulsion).
repulsion).
Examples:
Phosphorus 15P
n=3, l=1, ml=+1
Sulfur 16S
"electrons"electrons-inin-boxes"
Page 52
26
"Full" Outer Valence Shell
Noble Gas Configuration
The number of valence electrons (VE) in the outer shell
determines the chemical behavior.
"Full", closedclosed-shell configuration have a favorable low energy.
Elements with a closedclosed-shell (no valence electrons)
(18. group or VIII. main group: He, Ne, Ar,
Ar, etc.) show very
little tendency to participate in chemical reactions
Since they are all gases they are called noble (or inert) gases.
gases.
Electron configurations with a closedclosed-shell are called
noble gas configuration.
Atoms of other elements can reach lower energy states
by undergoing chemical reactions. By loss or gain of electrons
they can reach a noble gas electron configuration.
configuration.
Page 53
Noble Gas Configuration/Octet Rule
Examples
Octet rule (of thumb): Atoms tend to combine in a way that they
each have eight electrons in their valence shells, giving them
the same electronic configuration as a noble gas
(applicable to the mainmain-group elements, except helium).
Examples:
Noble gas:
1s2
-e-
11Na
1s2 2s2 2p6
9F
1s2 2s2
3s1
+e2p5
10Ne
2s2 2p6
Na+
1s2 2s2 2p6
F1s2 2s2 2p6
Atoms that lose electrons
have less negative charges
in shell than positive protons
in the core and become
positive cations.
cations.
Atoms that gain electrons
become negative anions.
anions.
Page 54
27
More Examples for
Noble Gas Configuration
Potassium (19K, alkali) looses 1 e-, becomes K+ cation,
with the same electron configuration like Argon 18Ar.
Its chemical and physical properties are of cause
completely different (different number of protons).
The same electron configuration achieves chlorine
(17Cl, halogen) by gaining 1e- becoming Cl- (chloride ion).
In a similar manner calcium (20Ca, alkali earth metal)
can lose two e- ⇒Ca2+, or 16S + 2 e- → S2both ions have the same electron configuration as 18Ar
By loss or gain of electrons particles reach the favorable,
low energy closedclosed-shell (noble gas configuration).
Page 55
Chemical Bonds + Molecules
Types of Bonds
When atoms of different elements form chemical compounds,
their chemical and physical properties change significantly.
Example: Sodium (metal)
Chlorine
Salt (NaCl)
+
Three types of chemical bond
nonmetalmetal with nonnonmetalnonnonmetal
metal (often)
Covalent bond Ionic bond
molecules
e. g. H2O
metal with
metal
metallic bond
salts
metals (alloys)
e.g. Na+Cl- e. g. bronze
coordinate
bond
metal complexes
Page 56
28
Metallic Bonding
Metallic bonding: Electrostatic attractive force between positively
positively
charged cations (nucleus + inner e-), which form a 3D lattice and
negative, delocalized, shared valence electrons ("
("electron sea").
The free mobility of the valence
electrons causes electrical conconductivity of metals.
Metal bonding exists between
atoms of a pure metal,
as well as in blends of two
or more metals (alloys).
Common alloys: bronze = copper/tin; brass= copper/zinc
Rose gold=gold/copper; white gold=Augold=Au-Ag/Ni/Pd
Amalgam: mercury/silver
Page 57
Ionic Bond
Ionization
When atoms lose or gain electrons, ions are formed:
Ions with positive charge are called cations.
cations. Na They have lost electrons.
In return, anions have gained electrons
to fill their valence shell.
They are negatively charged.
e- Na+
Cl + e Cl-
The atomic number (number of protons
in nucleus) does not change.
The mass does not change significantly.
Whether an atom tends to lose or gain one or more electrons
depends on the related change of its energy level.
The system as a whole strives to achieve the lowest energy state.
state.
Metals tend to lose electrons –
nonmetals tend to gain electrons.
Page 58
29
Ionization Energy
Within a period the number of protons
increases binding electrons stronger ⇒
Ionization energy increases.
Full and semi-filled shells are
favorable.
Within a group,
group, resp.
from period to period,
period,
the outer electron are shielded
agaist the nuclear charge
by the inner electrons
⇒ I.I.-Energy decreases.
decreases.
Element 2s
Li
Be
B
C
N
O
2p
Ionization
energy
eV]]
Ionization energy [[eV
Ionization I energy required
to remove an electron
from an atom.
⇒ cation
(Unit: eV = 1,6·
1,6·10-19 J)
F
Ne
atomic number →
Page 59
Electron Affinity
The electron affinity EA is the counter part of the ionization energy.
energy.
Amount of energy released when an electron is added to an atom
⇒ negative anion.
anion.
Positive values:
System would go to higher
energy state ⇒ gain of electrons
not possible.
Negative EAEA-values:
Energy is released.
Halogens and chalcogens
have high electron affinities
⇒Large amounts of energy is
released, if they gain electrons.
Electron affinities (eV
(eV))
Page 60
30
Electronegativity (EN)
Electronegativity (EN) describes the tendency of an atom (or a funcfunctional group) to attract electrons (or electron density) towards itself.
I
II
III
IV
V
VI
VII
H
2,20
VIII
He
Li
Be
0,97 1,47
B
C
N
O
F
2,01 2,50 3,07 3,50 4,10
Ne
Na
Mg
1,01 1,23
Al
Si
P
S
Cl
1,47 1,74 2,06 2,44 2,83
Ar
K
Ca
0,91 1,04
Ga
Ge
As
Se
Br
1,82 2,02 2,20 2,48 2,74
Kr
Rb
Sr
0,89 0,99
In
Sn
Sb
Te
I
1,49 1,72 1,82 2,01 2,21
Xe
EN increases from
bottom left to
top right in the PSE
Alkali (I. MG) and alkali
earth metals (II. HG)
electropositive. ⇒
Ionic bonds with
halogens (VII. MG) and
chalcogens (VI. HG)
For organic chem.:
Hal/O >> C > H
Larger difference in EN ⇒ more polar bond; ∆EN>1,5 ⇒ ionic
Compounds of alkalialkali-/alkali earth metals and chalcogens
resp. halogens are ionic as a rule of thumb.
Page 61
Atomic and Ionic Radius
Gain or loss of electrons alters the radius of the particle.
The radii of atoms and their respective ions are significantly different.
Loss of electrons (⇒
(⇒ cation) decreases the radius,
since the residual electrons are attracted more by the nuclear charge.
charge.
Gain of electrons (⇒
(⇒ anion) increases the radius, since the electron
shell becomes wider by the additional negative charge.
Examples:
F-
Li+
Na+
Mg2+
Cl-
K+
Ca2+
BrBr-
I-
Page 62
31
Salts, Ionic Lattice
Salts are solids, that consist of ions (ionic bond).
Cations (+) and anions ((-) are attracted by electrostatic
forces.
forces. The ions are positioned in a threethree-dimensional crystal lattice
with regular longlong-range order.
Such ionic crystals have high melting temperatures, are hard + brittle.
brittle.
Example: NaCl crystal (excerpt)
Each ion is surrounded by six counter
ions in octahedral geometry
With consideration of
the ionic
radius.
Page 63
Chemical Formulas of Salts
The ionic crystal lattice is a continuous structure with no overall
overall charge
(electrical neutral). Positive charges of cations are compensated by
negative charges of anions ⇒ defined molar ratio (stoichiometry)
of cations and anions.
anions.
Since cations and anions can carry different amounts of charge,
their ratio can differ from 1:1.
x resp. y = 1
General composition of a salt: {x⋅An+ y⋅Bm-} = AxBy
omitted
Ionic notation
Empirical formula
(displays charges)
(only molar ratio)
Examples:
Elements
in compound
Na
Combination of
K
ions in comcompound so that Ca
charges are
Ba
compensated
Al
Al
Br
S
O
Cl
F
O
Ions with
closedclosed-shell
Na+
K+
Ca2+
Ba2+
Al3+
Al3+
BrS2O2ClFO2-
Formula
Name
of salt
NaBr {Na+Br-} sodium bromide
K2S {2K+S2-} potassium sulfide
CaO {Ca2+O2-} Calcium oxide
BaCl2 {Ba2+2Cl-} Barium chloride
AlF3 {Al3+3F-} Aluminum fluoride
Al2O3 {2Al3+3O2-}Aluminum oxide
Page 64
32
Covalent Bond
Atoms with low EN differences can form shared pair of electrons
The electron shells (or orbitals) overlap ⇒ covalent bond.
bond.
⇒ molecules
+
+
H-atoms
+
+
H2-molecule
Emp. formula
H2
HF
+
Structural formula
H:H
H-F :
O ::
H H
H-H
H-F *
: :
Examples:
Examples: Compound
Hydrogen
Hydrogen fluoride
: :
+
Water
H2O
H
O
H
* H forms exclusively covalent bonds, because no free H+ (proton=elemental part.)
Valence electrons not involved in covalent bonds
form free (or lone) electron pairs.
pairs.
Octet rule (of thumb): Atoms tend to combine in a why that they each
have 8 electrons (4 electron pairs) in their valence shell
Page 65
Bond Character
The presented types of bonds are only models.
Real bonds can be intermediates; bond character
varies continuously throughout the PSE.
Examples:
Compounds of chlorine with elements of the 3. period:
Covalent bond
Ionic bond
Compounds of sodium with elements of the 3. period:
Metallic bond
Ionic bond
Elements of the 3. period
Metallic bond
Covalent bond
Page 66
33
Chemical Compounds
Molecules, Molecular Mass
A molecule is an electrically neutral group of at least two atom
(of same or different element) held together by covalent chemical
chemical bonds
All gaseous elements*
elements* form diatomic molecules:
H2, N2, O2, F2, Cl2
*besides noble gases
The nonmetals bromine (liquid) and iodine (solid) too: Br2, I2
The molar mass of a molecule is the sum of the molar masses of
the constituting atoms.
Examples:
Examples:Formula Molar mass of atom Molar mass [g/mol]
HCl
H=1,00 Cl=35,45
Cl=35,45
1,00 + 35,45 = 36,45
CH4
H=1,00 C=12,01
12,01 + (4·
(4·1,00)=16,01
H2SO4
S=32,07 O=16,00
(2·
(2·1)+32,07+(4
1)+32,07+(4··16)=98,07
Page 67
Stoichiometric Valency
The valency is a measure of the number of bonds formed
by a given element. It depends on the number of valence
electrons.
Rules: • H has always valence number 1
• Alkali metals (I. MG.) have always valence number 1
• Alkali earth metals (II. MG) always 2
• Elements of III. MG have valency 3
• Halogens (VII. MG: F, Cl,
Cl, Br, I) are always single valent
• O very often valency 2
• Other elements different valency (MG. elements max. 4)
• Products of lower case indices and valency in binary
compounds constant (last example)
Examples:
Examples: H2O: 22-valent oxygen binds two single valent H
NaCl: Both single valent ⇒ combine 1:1
MgCl2: 22-valent Mg binds 2 single valent Cl
Al2O3: 2·Al (3(3-valent): 2·3 = 6; 3·O (2(2-valent): 3·
3·2 = 6
Page 68
34
Chemical/Structural Formulas of Molecules
Single/Double/Triple Bonds
Chemical (empirical) formula:
formula: Simplest whole number ratio
of atoms of each element present in a compound.
Structural formula:
formula: Graphical representation of molecular structure,
showing how the atoms are arranged.
= Valence structure = LEWIS formula
All valence electrons (bonding and lone electron pairs) represented.
represented.
Hydrogen
S H
Ammonia H N H
sulfide
More examples:
NH3
H2S
H
H
Some compounds (in particular of C, N, and O) contain
double or triple bonds:
Examples:
Oxygen
O O
Nitrogen
N N
Carbon dioxide O C O
Check valency and octet rule!
Page 69
Polyatomic Compound Ions
Ions can also be formed from groups of joined atoms: Compound ions
ions
Most common compound ions:
Ammonium ion NH4+
Hydroxide ion OH-
H
H
+
N H
Nitrate ion NO3O
-
O H
H
O
S
-
O
O
O
Sulfate ion SO42O
N
O
Carbonate ion CO32O
-
O
C
O
-
Hydrogen carbonate
ion HCO32O
-
O H
C
O
Page 70
35
Chemical Models
Mesomeric Resonance
"This is no carbonate ion !"
It is a (crude!) model
of a carbonate ion.
Experiments indicate:
All OO-atoms same charge;
all bonds identical;
all bond angles identical.
Rene Margritte:
Margritte:
"This is no apple"
Christopher Ingold (1933): Mesomerism =
Representation of delocalized electrons
by different contributing structures.
Page 71
Double Bond Rule
(empirical)
empirical)
Elements of the second period (C, N, O) often form double bonds.
The analogous elements of the 3. period (and higher), however,
form in analogous compounds form preferably single bonds.
Main group: → IV.
V.
VI.
(N2)
(O2)
2. Period:
(CO2)
3. Period:
(S8)
(SiO2)
(P4)
Page 72
36
Bond Theory
Orbital Hybridization (sp3)
Carbon essential element in organic compounds; forms covalent bonds.
bonds.
C has 4 electrons in valence shell: configuration 2s2 2p2 (ground state)
1 lone electron pair (s2) and 2 pp-electron could form only 2 bonds.
Prediction: CH2 But: Carbon is 44-valent! Example.: CH4 ?!
Explanation: The valence electrons are rearranged in a way,
that 4 electrons can form bonds.
The ss-orbital and the three pp-orbitals are mixed to form
four sp3-hybrid orbitals of identical energy.
Hybridization is not a physical fact -,
it's only an model to explain experimental observations!
(e. g. molecule geometry) !
Page 73
sp3-Hybrid Orbitals
Geometry of hybrid orbitals
+
p
s
hybrid orbital
The 4 sp3 hybrid orbitals
arranged in space with a
maximum separation:
tetrahedral
Page 74
37
Single Bonds
Two hydrogen atoms form a single bond (σ -bond)
by headhead-on overlapping between the atom orbitals.
Highest electron density between the two nuclei.
σ-bond is symmetrical with respect to rotation about bond axis.
Other orbitals
can also form
σ-bonds:
Page 75
Single Bonds in Hydrocarbons
C-H-bonds and CC-C single bonds are σ-bonds.
Four σ-bonds ⇒ complies with octet rule.
Example:
Methane
4 CC-Hσ-bonds
Example:
Ethane
3 CC-H1 CC-Cσ-bonds
Free rotation about σ-bonds
Page 76
38
sp2-Hybridisation
Example: Boron
hybridization:
Boron (ground state):
1s2 2s2 2p1
E
One ss- and two pporbitals mix to form
tree sp2 hybrids of
equivalent energy
each occupied
by one electron
E
2p
2sp2
2s
1s
1s
TrigonalTrigonal-planar
geometry of
three sp2-hybrid
orbitals in space.
F
F Boron trifluoride
(BF3) does not
B
comply with
F octet rule
⇒ electron gap
⇒ Lewis acid
Page 77
sp2-Hybridisation of Carbon
E
carbon
2p
2sp2
1s
The 2s2s-orbital and the two 2p2p-orbitals mix
to form three sp2-hybrid orbitals.
The third pp-orbital unaffected.
unaffected.
Each occupied with single electron.
In an ethene molecule three sp2-hybrid
electrons form three σ-bonds:
Two with the H-atoms and one with
the other CC-atom (trigonal
(trigonal--planar).
planar).
The fourth electrons occupy the
remaining, unhybridized pz-orbitals,
sidewise overlapping above and
beneath the plane to form a π-bond.
bond.
H
C
sp2-hybrid
orbitals of carbon
and bonds formed in ethene.
H
H
C
C-C-σ-bond and π-bond
together form double bond.
H
Page 78
39
spsp-Hybridization
Which elements can mix one ss- and one pp- orbital to two sp hybrids
to form two σ-bonds and how are they arranged in space?
Beryllium (II. main group) can form two spsp-hybrid orbitals,
which can form two σ-bonds (linear, 180°
180°) e. g.: HH-BeBe-H
The two spsp-hybrids of each CC-atom
form two σ-bonds (C(C-C and CC-H).
The two remaining pp-orbitals
overlap crossbred sidewise
to form two CC-C π-bonds.
bonds.
One C-C σ-bond
H and
two π-bonds together form a
triple bond:
H
sp-hybrid orbitals of carbon
and bonds formed in ethyne.
H
C
C
H
Page 79
Polar Covalent Bonds
Dipolar Molecules
A polar bond is a covalent bond between two atoms
with different electronegativities ∆EN>0.
EN>0.
The electron density of the bonding electron pair is higher
in the vicinity of the atom with the higher electronegativity.
Partial charges and a permanent dipole are formed.
δ+ δ-
H Cl
H Cl
EN: 2,2 2,8
Dipolar momentum
O
δ-
δ+
C
δ+
O
_
δO
EN (H): 2,2 H H
+
δ+
Water molecule angled ⇒
partial charges asymmetrical
⇒ dipole molecule
⇒strong intermolecular forces
⇒ liquid (high boiling temp.)
EN (O): 3,5
δ-
Carbon dioxide: linear
partial charges symmetrical
⇒ no dipolar molecule
⇒weak intermolecular forces
⇒ gas (low boiling temp.)
Page 80
40
Cleavage of Covalent Bonds
Before new bonds can be formed in a chemical reaction,
existing bonds must be cleaved by high temp. or energetic radiation.
radiation.
Distinction between:
Homolysis: Cleavage of nonnon-polar covalent bond into two radicals:
Radicals are atoms or molecules with a nonnon-shared electron.
They are very energetic, have high reactivity and short lifetime.
lifetime.
UV, heat
Initiation of chlorine/
Example: Cl Cl
Cl⋅
Cl⋅ + ⋅Cl hydrogen gas explosion
Heterolysis:
Heterolysis: Cleavage of a polar covalent bond into cation + anion:
H
O
H Cl
H
δδ+
+
Proton (H ) transfer
H O⊕ H
H
+ Cl
Page 81
2.3.4 Intermolecular (physical) Forces
strengh
Different attractive interactions can occur between
molecules, distinguished by strength and origin:
• HydrogenHydrogen-(bridge)(bridge)-bonds: ≈ 10kJ/mol; by polar H
• DipoleDipole-dipole interactions: ≈ 5kJ/mol by dipoles
• VanVan-derder-Waals interactions: ≈ 1kJ/mol by all atoms
In contrast to intramolecular chemical (primary) bonds inside molecules,
these intermolecular forces between molecules are physical
(secondary) bonds.
At the boiling temperature the kinetic energy of the molecules
exceeds the attractive forces.
Page 82
41
Hydrogen Bonds
Hydrogen bond = attractive interaction of a hydrogen atom with an
electronegative atom (N, O, F, ...) that comes from another molecule
or chemical group. The hydrogen must be covalently bonded to
another electronegative atom to create the bond.
These phys. bonds can occur between molecules (intermolecular),
or within different parts of a larger single molecule (intramolecular
).
(intramolecular).
This type of bond occurs in both inorganic molecules such as water
water
and organic molecules such as proteins, cellulose or DNA
molecule molar mass boiling temp.
Hydrogen (bridge)
bonds in water
CH4
16 g/mol
-161 °C
NH3
17 g/mol
-33 °C
H2O
18 g/mol
100 °C
Boiling temperature is increased
by hydrogen bonds
Page 83
DipoleDipole-Dipole Interactions
Dipole–
Dipole–dipole interactions = electrostatic interactions of permanent
dipoles in molecules (asymmetrical polar bonds; e. g. H2O, HCl, HF).
Attractive force between positive part of molecule and negative
part of neighbored molecule
Polar atomic group without hydrogen:
e. g. carbonyl group (org. chem.)
Melting temperature
Boiling temperature
O
δ+ C
δ-
asym.
asym.⇒ dipole
17 °C
sym.⇒
sym.⇒ no dipole
-76 °C
44 °C
-10 °C
Phase transition temperatures increases by dipolar interactions.
Page 84
42
London Dispersion Forces
(VanVan-derder-Waals Interactions)
Interactions)
London dispersion forces (LDF)
also called VanVan-derder-Waals forces
in the strict sense are weak
intermolecular interactions
arising from induced instantaneous
polarization by random fluctuation of
the electron density.
The LDF between molecules increases with surface:
Example:
n-pentane
old name: pentane
boiling T: 36°C
2-methyl butane
2,2-dimethyl propane
old name: iso-pentane old name: neo-pentane
28°C
9,4°C
Page 85
physical
chemical
Chemical Bonds and Physical Bonds
Summary
Bond type
Energy (kJ/mol)
positive charge
negative charge
Ionic bond
Lattice enthalpy
≈1000
Cations
anions
metallic bond
Cations
delocalized electrons
(electron sea)
Covalent bond
≈ 300 per bond
Nucleus
Electron shell
Hydrogen bond
≈ 20 kJ/mol
H atom (polar bond to
Electronegative atom
electronegative atom, N, O,) with lonelone-pair (O, Cl)
Cl)
DipoleDipole-dipole
dipole
≈ 5 kJ/mol
permanently positively
polarized atom in mol.
VanVan-derder-WaalsWaalsbond ≈ 1 kJ/mol
instantaneously positively
polarized atom
permanently positivepositively polarized atom
instantaneously posipositively polarized atom
Page 86
43
Polar/nonPolar/non-polar Solvents
hydrophilic/lipophilic
Several types of chemical reactions occur in solution, if a suitable
suitable
solvent for the reactants can be found. The solubility depends on
on
polarity of solvent and solute. The number, potency and arrangement
arrangement
of polar bonds in the molecule influences the polarity of the solvent:
solvent:
Benzene Ether Ester Alcohol Water
Polarity
Hydrophilicity
Hydrophobicity = Lipophilicity
Polar solutes or atomic groups
are water soluble = hydrophilic
Fats are nonnon-polar and not
water soluble (hydrophobic
(hydrophobic))
NonNon-polar compounds
are lipophilic.
Ionic (Na
(Na+Cl-) or polar compounds (methanol) exhibit good solubility
in water. NonNon-polar compounds (fat, oil, grease) are soluble in nonnon-polar
solvents (white spirit, ether, chloroform, etc.)
Page 87
3. Physical Chemistry
Topics:
• State of matter/phase transitions
• Phase diagrams
• Homogeneous/heterogeneous systems
• Separation processes
• Thermodynamics
• Kinetics
• Chemical equilibrium
Page 88
44
State of Matter
Phase Transitions; Separation
Sublimation
Melting
solid
Evaporation
liquid
gaseous
Temp.
Condensation
Solidification
Resublimation
A phase is a region of space (in a thermodynamic system) throughout
which all physical properties of a material are essentially uniform.
uniform.
States of matter are phases. The change of the state of matter is
is a
phase transition.
The state of matter of a material or compound depends on
temperature T and pressure p ⇒ p-t-phase diagram
Page 89
pressure p
Phase Diagram
Position of
melting curve,
curve,
boiling curve and
sublimation curve
substancesubstance-specific.
liquid
solid
Triple point:
solid, liquid, and
gaseous phase
coexist.
Normal
pressure
Triple point
Ambient
temp.
Tm
Example:
Example: Solid
compound at nornormal conditions
annealed until it
Tb Temperature T melts (Tm) and
evaporates (Tb)
gaseous
Page 90
45
Pressure p
Liquefaction of Gases
Gas can be
liquefied by
cooling or
compression
solid
liquid
Normal
pressure
gaseous
Temperature T
Ambient
temp.
Page 91
Pressure p [hPa]
Phase Diagram of Water
Anomaly:
Anomaly:
Liquid water has a
higher density than
ice.
Water
Ice
1013
Water vapor
0
(Negative slope of
melting curve)
Skating:
Skating:
At const. low temperatemperature (~ -5°C) pressure
is increased
⇒ ice melts
Pressure cooker:
Higher pressure ⇒
100 Temperature T [°C] higher boiling temp.
⇒ faster cooking
Page 92
46
Multi Phase Systems
heterogeneous Systems
heterogeneous Systems:
Systems: Two or more components (= phases).
Physical (and chemical) properties change instantaneously
at phase boundaries.
boundaries.
Differentiation by combination of states of matter (Example)
(1) solid/solid
Mixture
(Granite = feldspar + quartz+ mica)
(2) solid/liquid
Suspension (Blood = blood cells in blood plasma)
(3) liquid/liquid
Emulsion
(Milk = fat droplets in water)
(4) solid/gaseous
Smoke; foam (cigarette smoke, Styrofoam)
(5) liquid/gaseous
Fog; foam
(Fog, whipped cream)
= Aerosols
Heterogeneous mixtures can of be separated into homogeneous
(uniform) components by physical operations like filtration,
centrifugation, decantation, etc.
Page 93
Homogeneous Systems
In homogeneous systems the properties are uniform throughout the
whole sample; no phase boundaries.
Homogeneous systems can be pure substances or multi component
systems (=solutions).
Different types of solutions:
(1) solid/solid
(2) solid/liquid
(3) liquid/liquid
(4) solid/gaseous
(5) liquid/gaseous
(6) gaseous/gaseous
(metal alloys, z. B. bronze)
(sugar dissolved in water)
(alcohol + water)
(solution of hydrogen in platinum)
(air dissolved in water)
(air)
Homogeneous mixtures can also be separated into their components
by physical operations, e. g.:
• Distillation: Separation by different boiling points; (distilling)
(distilling)
• Extraction: Separation by different solubility's; (coffee cooking)
cooking)
Page 94
47
Distillation
Thermal separation process: Mixture is separated into components by
the differences in their boiling temperatures.
Mixture (2) is annealed (1); More volatile component (lower Tb) evapoevaporates. Thermometer (4) indicates temperature of vapor (3)
= boiling temperature.
In condenser (here Liebig cooler) (5)
with cooling water plugs (6;7)
the more volatile component
condenses and drops into receiving
flask (8).
Volume extension during
evaporation ⇒ pressure equalization (9).
High boiling compounds bp.>200
bp.>200°°C:
low pressure (vacuum) distillation.
⇒ Lower boiling temperatures.
Examples:
Examples: Essential Oils, pure alcohol,
sea water ⇒ freshwater
Page 95
The Chemical Reaction
A chemical reaction is a process that leads to the transformation
of one set of chemical substances to another. Existing bonds between
the atoms of the reactants (=reagents) are partly broken, new bonds
bonds
are formed
Reaction equation: Reactant 1 (+ react. 2)
product 1 + (product 2)
The connections among the atoms are altered, the number of atoms
of the respective element remains constant.
During a chemical reaction the total
mass of matter remains constant.
Law of conservation of mass
Antoine Lavoisier
fr. Chem. (1743-1794)
Page 96
48
Chemical Thermodynamics
Chemical Kinetics
The electronic structure of the reactants (valency, polarity, etc.)
etc.)
indicate which chemical reactions are possible in general,
but do they really take place and under which conditions?
Sometimes side reactions can occur. Which product is formed?
Whether a chemical reaction takes place, to which extent and how
fast is determined by the chem. thermodynamics and kinetics.
kinetics.
Chem. thermodynamics:
thermodynamics:
Under which conditions does a reaction occur spontaneously ?
Is heat released or consumed?
Chem. reactions always involve conversion of energy.
Whether a reaction occurs or not is related to the available
chemical energy stored in the reactants and products.
Reaction kinetics:
kinetics: Study of rate of chemical reactions:
How much product is formed in time under given experimental
conditions (e. g. concentration).
Page 97
Reaction Enthalpy
Exothermal/Endothermal Reaction
Reaction enthalpy (∆HR) (= reaction heat) = how much heat
is released or consumed during reaction.
Enthalpy negative ∆HR<0 ⇒ heat released = exothermal
EnthaIpy pos. (∆
(∆HR>0) ⇒ heating necessary = endothermal.
endothermal.
∆HR<0
exothermal
∆HR>0
endothermal
enthalpy
products II
reactants
∆HR
endothermal
reaction
exothermal
products I
Chemical energy is stored in chemical compounds.
If the products (I) have lower energy state than reactants ⇒
exothermal, energy released. (heat, light, bang)
If products (II) have higher energy content, reaction is endothermal,
endothermal,
feed of energy required (heat, electricity).
Page 98
49
Example for
Exothermal/Endothermal Reaction
Example:
HydrogenHydrogen-oxygen reaction
2 H2 + O2
2 H2O
Electrolysis
Oxyhydrogen reaction
H2-O2-reaction ∆HR = - 286 kJ/mol (released; heat, light, bang)
During electrolysis of water ∆HR = 286 kJ/mol
(consumed; electrical power)
Page 99
Enthalpy of Formation
Calculation of Standard Enthalpy of Reaction
The chemical internal energy = enthalpy H of a system cannot
be measured directly. Change in enthalpy ∆H more useful quantity.
By definition all elements have an enthalpy of zero.
Enthalpies are funct.
funct. of T und p ⇒ Standard cond. (25°
(25°C, 1 bar): H0
Molar Enthalpy of Formation of a compound ∆Hf = amount of heat
released or consumed during its formation from the elements.
Example:
Example: Molar standard enthalpy of formation of water:
H2 + ½ O2 → H2O (l)
∆HR = -286 kJ/mol
∆Hf = 0
0 -286 kJ/mol
The reaction enthalpy can be calculated from the enthalpies of forma
forma-tion of involved compounds:
∆HR0 = ∑∆H0f ( products
) − ∑∆H0f (reactants)
Example:
Example: Combustion of ethene:
∆Hf0
C2H4 + 3 O2 → 2 CO2 + 2 H2O (g)
∆HR = -1322,9 kJ/mol
= 52,3
0
2·
2·(-393,5) 2·
2·(-241,8) kJ/mol
Page 100
50
Hess Law
Hess: the enthalpy change for a reaction carried out in a series of steps
is equal to the sum of the enthalpy changes for the individual steps
steps
Example:
Example: ∆Hf CO not directly measurable, because CO2 is formed too.
∆Hf (CO)= -110,5 kJ/mol (calculated)
(
rea
ctio
n
re
ac ∆HR = -283
tio
n kJ/mol
∆Hf (CO2)= -393,5 kJ/mol
(measured)
∆Hf (CO) = ∆Hf (CO2) - ∆HR (CO→
(CO→ CO2)
Page 101
Hess Law
and Biological Processes
Photosynthesis
Combustion
Respiration
(Glycolysis,
Glycolysis, Citrate cycle,
cycle, ...)
Page 102
51
Entropy
Another thermodynamic property of matter is the entropy:
entropy: Interpretation
as molecular disorder or statistical number of possible arrangements.
arrangements.
Ssolid <
solid
Sliquid <
liquid
Sgas
gas
Spure substance< Smixture
pure subst.
mixture
Page 103
Change of Entropy
In closed thermodynamic systems, the entropy remains
constant for completely reversible processes.
For irreversible processes the entropy increases.
Examples for irreversible processes
with increase of entropy:
Heat exchange:
60 °C
50
40
20 °C
30
40
Diffusion:
Reaction entropy:
∆S R0 = ∑ ∆S 0 ( produkte) − ∑ ∆S 0 ( Edukte)
Page 104
52
Gibbs (Free) Energy
The reactions enthalpy indicates, if process is exoexo- or endothermal,
but not, if reaction occurs spontaneously. Endothermal spontaneous
spontaneous
reactions can occur, which consume heat from the environment,
which is cooled.
Driving force for a reaction is the Gibbs (Free) Energy ∆G.
GibbsGibbs-Helmholtz equation: ∆G = ∆H -T·∆S
Reactions occur spontaneously, if ∆G < 0 (exergonic
(exergonic reaction).
If ∆G > 0 ⇒ endergonic reaction, absorbs energy from surrounding.
„somebody has
to lift the ball“
ball“
Page 105
Activation Energy
Catalysis
energy
The activation energy is the minimum energy required to start a chemichemical reaction („
(„igniting spark“
spark“) (cleavage of existing bonds), even for
exothermal reactions in which the released heat exceeds the activation
activation
energy by far.
A catalyst (chem. compound
added in small amounts
to reaction)
Activation energy •lowers activation energy,
with catalyst •increases reaction rate
reactants
•doesn't influence equilibrium
reaction
•is not consumed by reaction
H-H
reaction
enthalpy
∆HR < O
products ⇒ exothermal
H⋅⋅⋅H
⋅⋅⋅H
reactant
Platinum
Page 106
53
Kinetics
Reaction kinetics: Speed of chemical reactions.
Reaction rate defined as change of concentration with time
For a reaction: A + B → P:
r=−
d [ A]
d [ B] d [ P]
=−
=
dt
dt
dt
i.e.: increase of concentration of product P
resp. decrease of reactant A or B with time.
Monitoring progress of reaction by
• change of color
• spectroscopic
• conductometric
• calorimetric (heat measurement)
reduction of methylene
blue by glucose
Page 107
Rate Law
Reaction Order
The reaction rate depends on the concentration of the involved
compounds ⇒ rate law:
law:
d [ A]
a
b
For A + B → P:
r=−
dt
= k (T ) ⋅ [ A] ⋅ [ B]
The order of reaction is defined as the sum of all the exponents
of the reactants involved in the rate equation.
k is the rate constant, which is temperature dependent as a rule.
reaction
order
rate law
A→P
1. order
-d[A]/dt = k·[A]
[A]
A+B→P
2. order
-d[A]/dt = k·[A]·
[A]·[B]
[B]
Page 108
54
Chemical Equilibrium
Reversible/Irreversible Reaction
In general, a chemical reaction can be reversed:
The products can react with each other to form the original reactants.
reactants.
Such a reaction is called a reversible reaction.
Reversible reaction:
A+B
C+D
Forward and reversed reaction occur simultaneously so that
a dynamic equilibrium is formed, in which A and B react continuously
to C und D, while C and D form A and B again. Once the equilibrium
equilibrium
is reached, the net concentrations remain constant.
A reaction without reversed reaction is called irreversible.
irreversible.
It occurs until A and/or B are consumed completely or, if C and/or
and/or D
are removed from the reaction site.
Irreversible reaction: A + B
C+D
Page 109
Law of Mass Action
The equilibrium state of a chemical reaction can be described
quantitatively by the Law of Mass Action (Guldberg/Waage 1865).
A+B
C+D
Equilibrium constant kc =
k← c(C ) ⋅ c( D)
=
k→ c( A) ⋅ c( B )
Concentrations are multiplied! By convention product ⇒ numerator
in general:
aA + bB ⇔ cC + dD
Almost complete
reaction:
Equivalent
equilibrium:
Barely occurring
reaction:
Application of
Law of Mass Action:
A+ B → C + D
A+ B C + D
A+ B ←C + D
kc =
[C ]c ⋅ [ D]d
[ A]a ⋅ [ B]b
kc>>1 ; z.B. 1012
kc=1
kc<<1 ; z.B. 10-12
• Solubility of salts
• Autoprotolysis of water
Page 110
55
Le Chatelier´
Chatelier´s Principle
Change of concentration:
concentration:
Changing the concentration of an
ingredient (C) will shift the equilibrium
to the side that would reduce the change.
By the reversed reaction C and D are
consumed partly, while the concentration
of A and B is increased until kC reached.
[C ] ⋅ [ D ]
kc =
[ A] ⋅ [ B ]
A+ B
C+D
A+ B ←
C+ D
C+
A+ B
A+ B
D
C + D + Energy
Change of temperature.:
temperature.:
Exothermal reaction has energy (heat) as product. Cooling removes
removes
heat and the equilibrium is shifted to the product side.
Endothermal reactions are enhanced by feeding energy.
Change of pressure/volume:
pressure/volume: In reactions with volume decrease (com(compression)
pression) increasing pressure shifts the equilibrium to product side.
Example:
Example:
N2 + 3 H2
2 NH3
A + B + P /V
C+D
Page 111
4. Inorganic Chemistry
4.1 Acids – Bases
Historically, acids are compounds, which taste sour (vinegar,
lemon). Those compounds
change the color of
certain vegetal ingreingredients to red (red
cabbage; litmus).
(R. Boyle, 1663)
acidic red cabbage extr.
extr. alkaline
Substances, which taste and feel soapy, were called alkalis
(arab.:Plant ash←Potassium hydroxide). Lye = corrosive alkaline
substance. Later such chemicals were called bases,
bases, since
they are (together with an acid) the base for formation of salts.
salts.
Bases can change the color of certain plant ingredients, too.
Lavoisier (~1770) assumed, that acids contain oxygen, because
nonnon-metal oxide (e. g. CO2) together with water form acids.
Observation:
Observation: (Hydrochloric) acid dissolves lime (CaCO3) and zinc
involving gas development. It does not react with copper.
Page 112
56
Acid Rain
The source of acid rain are the The produced acids are harmful for
oxides of carbon, nitrogen and plants, waters, and buildings:
sulfur, which are formed by the
combustion of fossil fuels
(coal, mineral oil, natural gas)
They react with water molecules
to produce acids.
Carbon dioxide:
• CO2 + H2O → H2CO3
Nitroxides:
Nitroxides:
• 2 NO2 + H2O → HNO2 + HNO3
Sulfur oxides:
• SO2 + H2O → H2SO3
• SO3 + H2O → H2SO4
1908
1968
limestone sculptures in industr. area
CaCO3+2H2SO4→CO2+H2O+Ca2++SO42solid
gaseous
soluble
Page 113
AcidAcid-Base Theory by Brø
Brønsted
Comprehensive acidacid-base theory by Brø
Brønsted (1923):
Acids: Proton donators; compounds or specimens,
that release protons.
Bases:Proton acceptors; compounds or specimens,
that absorbs (gains) protons.
Joh.
Joh. N. Brø
Brønsted
Advantage: Definition is
dan.
dan. Chem.
Chem.
• independent of water as solvent
• not restricted to liquid phase ⇒
HCl (g) + NaOH (s) →
• more bases, e. g. ammonia (NH3)
Na+Cl –(s) + H2O (l)
In water Brø
Brønsted
Bases must have a lone
acids
produce
pair to receive a proton
hydronium ions:
ions:
H-A +
O
H
H
A- + H O⊕H
H
Page 114
57
Dissociation Equilibrium
Conjugate Acid/Base
If an acid HA reacts with a base B, a proton is transferred from
the acid to the base (protolysis
).
(protolysis).
conjugate
acid
base
acid
base
In the reversed reaction, a proton is transferred from HB+ to Ai. e., HB+ acts as acid, A- as base.
In such a dissociation equilibrium two sets of
acidacid-base pairs occur: S1/B1 and S2/B2.
A- is called conjugate base to HA
and HB+ conjugate acid to B.
Acid and bases can be neutral molecules or ions.
Example:
Example: Acetic acid + ammonia → ammoniumammonium- + acetate ion
Page 115
Ampholyte
An ampholyte molecule (or ion) can either donate or accept
a proton, thus acting either as an acid or a base.
Example:
Water reacts toward nitric acid
as a base: It receives a proton and
becomes a hydronium ion.
Towards the base ammonia
water reacts as an acid, however,
releasing a proton and becoming
a hydroxide ion.
H2O + HNO3 → H3O+ + NO3H2O + NH3 → OH- + NH4+
More ampholytes:
HCO3- (Hydrogen carbonate ion), HSO4- (Hydrogen sulfate ion)
Acid/base designates a chemical behavior towards a reaction partner,
rather than a class of substance.
Names (e .g. nitric acid) are historical regarding reaction with water.
Page 116
58
AcidAcid-Base Theory by Lewis
1927; G.N. Lewis: Even more general def. for acids and bases:
(independent of protons or hydroxide ions)
LewisLewis-acid: Specimen, with empty outer orbital, ready to receive
electron pair; has electron pair gap (does not comply with octet
octet rule!)
LewisLewis-acid: Electron pair acceptor
LewisLewis-base has a free electron pair, which can add to a Lewis acid
to form a new electron pair bond.
LewisLewis-base: Electron pair donator
Example:
Brö
Brönstedtnstedtacid
New bond by overlap
of empty boron orbital
with occupied nitrogen
orbital
Lewis acid Lewis base Lewis acidacid-base aduct
Page 117
General Structure of Acids
What is the general structure of a Brø
Brønsted acid?
Brø
ø
nsted
acids
contain
polar
bonded
hydrogen.
Br
Which elements form polar bonds with H ?
NonNon-metals (electronegative!) form polar bonds with H.
Example:
Example: Hydrogen chloride H Cl other halogens analogous
Elements with lower electronegativity (S,N,P,C) can increase
polarity by additional bonds to oxygen atoms:
Examples: Nitric acid
Citric acid
Sulfuric acid
H
S
O
O
H
O
O
N
O
H
Page 118
59
O
O
Strong/Weak Acids/Bases
Acids and bases are graded strong or weak regarding their
tendency to release resp. receive protons.
The more polar the hydrogen is bonded, the greater the
tendency for dissociation = protolysis,
protolysis, the stronger the acid.
Hydrogen
Examples: Sulfuric acid
Sulfurous acid
sulfide
is
H
H
is
stronger
stronger
S
S
acid
acid
than:
than:
S
O
O
O
H
H
O
O
O
H
O
H
Strong acids (HCl, H2SO4, HNO3, H3PO4, HClO4)
dissociate in water almost completely:
completely: HA + H2O→H3O+ + AStrong acids have weak conjugate bases;
strong bases have weak conjugate acids.
Page 119
Polyacids
Some acids can release more than one proton via different
protolysis steps. The second step is hampered compared to the first.
Example:
Example: twotwo-proton acid carbonic acid (H2CO3).
H2CO3 + H2O
HCO3– + H3O+ 1. Stufe
step
HCO3– + H2O
CO32–+ H3O+
step
2. Stufe
conjugate acid-base-pairs
H2CO3 / HCO3– und H3O+ / H2O
HCO3– / CO32– und H3O+ / H2O
Intermediate (e. g. hydrogen carbonate ion) = ampholyte !
More
polypolyacids:
Oxalic acid H2C2O4
Sulfuric acid H2SO4
Phosphoric acid
H3PO4
Page 120
60
Concentrate/Dilute Strong Acids
Molarity
Mixing of pure, concentrate acid with water ⇒ dilute acid.
acid.
Warning: During dilution of concentrate acids (in particular sulfuric acid)
acid)
with water large amounts of heat are released;, rapid evaporation
evaporation of
water; uncontrolled spray of acidic liquid
Rule:
Rule: Add acid slowly to water – never add water to conc. acid!
Concentration c (= Molarity M (outdated but still common)) of acid/base
= Number of Mole dissolved acid/base in 1 L.
M > 1 mol/L ⇒ conc. acid; M < 1 mol/L ⇒ dilute acid
Example:
Example: conc. (69%mass) nitric acid has density ρ=1,41 g/cm3
1 L⇔
L⇔1410g; 69%⇔
69%⇔973g pure HNO3; molar mass=63 g/mol;
n=m/M = 973g/63g/mol=15,4 mol ⇒conc. nitric acid = 15,4 M HNO3
How much conc. acid for 100 ml of 0,1 M (dilute) HNO3 ?
n = V·c = 0,01 mol HNO3
1000ml
x
=
⇒ x = 0,65ml ⇔ 0,92g
15,4mol 0,01mol
Page 121
Strong Acids
Sulfuric acid:
acid: Made from sulfur trioxide (SO3) + H2O→H2SO4
100%: ρ=1,8 g/cm3; 1800g/98g/mol = 18 mol/L = 18 M H2SO4
additional phys. solution of SO3 ⇒ Oleum (hygroscopic)
Hydrochloric acid = aqueous solution of hydrochloric gas (HCl)
(HCl)(g)
(g);
max. 40 % ⇒13 mol/L hydrochloric acid
Complete dissociation: HCl + H2O → H3O+ + Cl> 36% = fuming acid. ⇐ HCl gas evaporates
Nitric acid = HNO3 ; concentrate = fuming 90% HNO3
⇒ free NO2 (brown toxic gas);
dissolves silver, but not gold (aqua fortis)
fortis)
Nitrating acid:
acid: Mixture of nitric + sulfuric acid for nitration
of compounds; e g. toluene ⇒ trinitrotoluene (TNT)
Aqua Regia = HCl/HNO3 (3:1) dissolves gold
Page 122
61
Common Acids
and their Conjugate Bases
Hydrochloric acid HCl(aq)
(aq)
Chloride ion
Cl-
Hydrofluoric acid HF
Fluoride ion
F-
Sulfuric acid
Hydrogen sulfate HSO4-
H2SO4
Sulfurous acid H2SO3
Hydrogen sulfite
HSO3-
Nitric acid
HNO3
Nitrate ion
NO3-
Nitrous acid
HNO2
Nitrite ion
NO2-
Phosphoric
acid
Carbonic acid
H3PO4
Dihydrogen phosphate
Hydrogen phosphate
H2PO4HPO42-
H2CO3
Hydrogen carbonate HCO3-
Perchloric acid HClO4
Perchlorate ion
ClO4-
Hydrogen cyanide
HCN
Cyanide ion
CN-
Acetic acid
CH3COOH Acetate ion
Sulfate
SO42-
Sulfite
SO32-
Phosphate PO43Carbonate CO32-
CH3COO-
Ampholytes
Page 123
Common Bases/Alkalis
Caustic soda = alkaline solution of sodium hydroxide
(NaOH) in water; max. 1200 g = 30 mol/L
exothermal process!
Production: Electrolysis of alkali chlorides
(see: Industrial Chemistry)
Chemistry)
Potash lye = alkaline solution of potassium hydroxide
(KOH) in water; exothermal process!
Ammonia water = Ammonium hydroxide = Sal ammoniac
← NH + + OHAmmonia (gas) dissolved in water: NH3 + H2O →
4
Page 124
62
Carbonic Acid
Gaseous carbon dioxide
soluble in water:
partly reacts to carbonic acid:
Carbonic acid is a weak
acid, dissociates only partly
forming hydrogen carbonate
ions (Ampholyt).
(CO2)(g)
CO2 (aq
(aq))+ H2O
(CO2)(aq).
H2CO3
H2CO3 + H2O
HCO3- + H3O+
In a second protolysis step
carbonate ions can be produced: HCO3- + H2O
CO32- + H3O+
In carbonated water all species coexist in a chemical equilibrium
equilibrium
in different concentrations. Addition of lemon (acid, H3O+) shifts all
equilibriums to the left ⇒ more CO2 gas escapes.
Page 125
Autoprotolysis of Waters
Salt solution are electrical conductive, since they contain ions.
ions.
Application of voltage ⇒ ions are moved: Cat+→ cathode (neg. pole)
: An- → anode (pos. pole)
Pure, distilled water is conductive too,
because water always contains ions due to autoprotolysis:
autoprotolysis:
H2O + H2O
[H3O]+ + [OH]-
Water acts as acid and as base,
hydronium ions and hydroxide ions are produced
in small concentrations.
Page 126
63
Ion(ic)
Ion(ic) Product of Water
Law of mass action
for autoprotolysis
Derivation of ion product of water
from law of mass action
The ion product of aqueous solutions is always 10-14 mol²
mol²/L²
/L²
By addition of hydronium
H2O + H2O
ion (from acid)
acid) shifts the
equilibrium,
equilibrium, H3O+ and OHH O + H2O
react until KW =10-14 again.
again. 2
Now,
Now, [H3O+]>10-7 and
H2O + H2O
[OH[OH-]< 10-7
[H3O]+ + [OH]-
[H3O]+ +[OH][H3O]+ +
[OH]Page 127
The pH Value
The concentration of hydronium ions is a measure for
acidic or alkaline character of a solution.
Representation as decimal number confusing
in powers of 10 inconvenient, therefore definition:
The pH value is the negative decimal logarithm of the
hydronium ion concentration. (p
(potentia hydrogenii)
ydrogenii)
pH = − lg[c( H 3O + )]
c(H3O+) mol/L decimal
c(H3O+) mol/L power
pH value
solution
1
100
0
very acidic
0.0001
10-4
4
moderate acidic
0.0000001
10-7
7
neutral
0.0000000001
10-10
10
moderate alkaline
0.00000000000001
10-14
14
very alkaline
Page 128
64
Relationship between Hydronium and
Hydroxide Ion Concentration
In neutral water is: c(H3O+) = c(OH-) = 10-7 mol/L
For every solution is: c(H3O+) · c(OH-) = 10-14 mol²
mol²/L²
/L²
Note: concentrations are multiplied!
Analogous to pH a pOH value can be defined:
The pOH value is the negative decimal logarithm of the
hydroxide ion concentration.
pOH = − log[c (OH − )]
c(H3O+)·c(OH-)=10-14 mol/L logarithm ⇒
Examples:
pH
2
7
11
pOH
12
7
3
pH + pOH = 14
solution
acidic
neutral
alkaline
Page 129
pH calculation
(strong acid or strong base)
base)
Example:
Example:
n =0,1mol of strong acid (HA) added to V=10 L water.
water. pH = ?
HA + H2O→H3O+ + A- (1 mol acid produces 1 mol hydronium)
hydronium)
c( H 3O + ) = nHA / V = 0,1mol / 10L = 0,01mol / L = 10−2 mol / L
pH = − log[c( H 3O + )] = − log10 −2 = 2
Example:
Example:
0,01 mol of strong base added to 100 L of water.
water. pH = ?
c(OH − ) = nB / V = 0,01mol / 100L = 0,0001mol / L = 10−4 mol / L
pOH = − log[c(OH − )] = − log 10 −4 = 4
pH = 14 − pOH = 14 − 4 = 10
Page 130
65
pH Values of Common Solutions
Substance
Battery acid
Gastric acid
Lemon juice
Coke
Vinegar
Orange/apple juice
Wine
Sour milk
Beer
Acid rain
Coffee
Tee
Rain
Mineral water
Milk
Distilled water
pH-Wert
Art
-0,5 Acidic
2,0
2,4
2-3
2,7
2,9
3,5
4,0
4,5
4,5 – 5,0
<5,0
5,0
5,5
5,6
6,0
6,5
7,0
Neutral
Substance
Saliva
Blood
pH-Wert
6,5 – 7,4
7,34 – 7,45 Alkaline
Sea water (today)
8,05
Sea water
(pre-industrial)
8,16
Intestine
Soap
Art
Acidic to
8,3
9,0 – 10,0
Sal ammoniac
11,5
Bleaching agent
12,5
Concrete
12,6
Caustic Soda
13,5
Page 131
pH Indicator
A pH indicator is a halochromic compound added in small amounts
to a solution to visualized acidity resp. basicity by color.
Indicators are colored compounds, which undergo acidacid-basebasereactions ⇒ chemical structure and color is changed.
Indicator enable a rough determination of the pH value:
Examples:
Litmus
Chemical structure of the
indicator methyl orange
Page 132
66
pH Test Strip/pH Meter
pH Test strips contain a mixture
of different indicators.
Comparison of color with
reference scale ⇒ pH value.
In a pHpH-Meter electrochemical
determination of H3O+ ion
concentration and calculated
pH value displayed.
Glas electrode
Page 133
Neutralization
Combination of an acidic solution with a certain with a certain
c(H3O+) with an identical volume of an alkaline solution with
the same c(OH-) produces a neutral solution (pH=7).
The underlying reaction is a neutralization.
neutralization.
H3O+ + OH-
H2O + H2O
If identical amounts of substance (equimolar
(equimolar amounts)
of H3O+ and OH- are combined, they react in a neutralization
reaction to neutral water.
Attributes of a neutralization:
• represents an acidacid-basebase-reaction, proton transfer H3O+ to OH-.
• exothermal reaction; reaction heat is released!
• used in quantitative analysis for determination
of unknown c(H3O+) or c(OH-): AcidAcid-base titration (later)
Page 134
67
Calculation of Neutralization Reaction
In a submarine 180 L acid are leaked out of the battery cells
with pH = 1. How much sodium hydroxide solution of c = 2 mol/L
is necessary for a complete neutralization (pH = 7) ?
Neutralization means:
amount of substance of acid (H3O+)
is equal to amount of base (OH-)
nA = nB
cA = 10-pH = 0,1 mol/L
cA⋅VA = cB⋅VB
VB = cA⋅VA/cB
VB =
0 ,1mol/L ⋅180 L
= 9L
2mol/L
Page 135
Formation of Salt by Neutralization
In addition to H3O+ and OH- -ions acidic and alkaline solutions
both contain counter ions (electric neutrality).
They can be product of a previous acidacid-basebase-reaction
from production of the acidic solution:
HCl (g) + H2O (l) → Cl- (aq)
aq) + H3O+ (aq)
aq)
Example:
or they result from the dissolution of a compound in water:
H O (l)
Example:
NaOH (s) 2 →
Na+ (aq)
aq) + OH- (aq)
aq)
Total reaction of the neutralization:
Cl-(aq)
(aq) + H3O+(aq) + Na+(aq)
(aq) + OH-(aq)
aq) → Na+(aq)
(aq) + Cl-(aq)
(aq) + 2H2O
If the water is evaporated NaCl (table salt) is remaining.
Acid + Base → Salt (+ Water)
Page 136
68
Quantitative Analysis
Titration
At equivalence point:
nknown = nunknown
cknown·Vknown.= cunknown.·Vunknown
Burette
Standard solution
Determination of the concentration of an unknown solution by defined
defined
addition of another solution (standard solution with known conc.)
conc.)
until equivalence point is reached, visualized by indicator.
cknown·Vadded
cunknown= V
receiving
Reaction has to be:
• complete
• known stoichiometry
• fast
• End point detectable
Suitable reactions:
• Neutralization
• Redox reaction
• (Complexation)
unknown
solution
Page 137
AcidAcid-BaseBase-Titration
(Principle)
Acidic solution, c(H3O+) = ?
Measuring a certain volume
of acid VA; Drop wise addition
of sodium hydroxide solution
with known concentration cB
until pH=7 (neutralization);
Measuring the consumption
of basic solution VB.
At equivalence point:
n A ⋅ ( z A ) = nB
with number of acid protons zA
with n
= c ⋅V
c A ⋅ VA ⋅ ( z A ) = cB ⋅VB
c A = cB ⋅
VB
VA ⋅ ( z A )
Page 138
69
Quantitative AcidAcid-BaseBase-Reaction
(Strong Acid/Base)
If a certain amount of base is added to an existing acidic solution,
solution,
the acid is neutralized partly.
What is the concentration of the remaining acid ?;
Which pHpH-value is reached ?
In the existing acid with a volume VA and a concentration
cA the amount of H3O+ is:
n
= c A ⋅ VA ⋅ z A
+
0 ( H 3O )
The addition of base produces a certain amount of OH- ions:
n(OH − ) = cB ⋅VB
Part of the H3O+ is neutralized; the remaining amout is:
n( H O + ) = n0 ( H O + ) − n(OH − ) with the conc.: c( H 3O + ) =
3
3
c A ⋅VA ⋅ z A − cB ⋅VB
VS + VB
Volumes have to be added
For identical conc. cA=cB und zA=1, simple equation:
c( H O + ) =
3
c ⋅ (VA − VB )
V A + VB
Page 139
Titration Curve
(Strong Acid/Strong Base)
Titration of 20 ml 0.1 molar HCl
with 0.1 M NaOH (c
(cA=cB , zA=1)
c( H O + ) =
3
0,1 ⋅ (20 − VB )
20 + VB
Addition VB c(H3O+) mol/L
NaOH [ml] c(OH-) mol/L
0
0,1
5
0,06
10
0,033
15
0,014
19
0,0025
19,9
2,5·
2,5·10-4
19,99
2,5·
2,5·10-5
20
21
30
40
10-7
0,0025
0,02
Equivalence point
nA = nB
pH value
Neutral point
pH = 7
1
1,22
1,47
1,84
2,6
3,6
4,6
7*
11,4
12,3
12,5
* Ion product of water relevant
c(OH − ) =
0,1 ⋅ (VB − 20)
20 + VB
Page 140
70
Acidity Constant;
pKA Value
In addition to very strong acids, which dissociate completely,
weaker acids exist, which dissociate only partly.
Their equilibrium reaction in water is described by
the law of mass action.
HA + H2O → H3O+ + A-
c( H 3O + ) ⋅ c( A− ) since c(H2O)≈
c( H 3O + ) ⋅ c( A− )
K=
KA =
c( HA) ⋅ c( H 2O ) const. ⇒
c( HA)
The acidity constant KS is a measure for the strength of the acid
Often, the negative decimal logarithm pKA is represented:
pK A = − log K A Analogous KB and pKB for bases
Important note: pH describes individual solution property
In contrast to: pKA describes compound property!
Page 141
Perchloric acid
HCl
Hydrochloric acid
H2SO4
Sulfuric acid
H3O+
pKA
ClO4-
-9
Cl-
-6
HSO4-
-3
Hydronium ion
H2O
Nitric acid
NO3-
-1,74
-1,32
acid
H2S
H2PO4-
7,21
7,25
Hypochloric acid
Ammonium ion
NH3
9,21
HCO3-
Hydrogencarbonate
CO32-
10,4
Hydrogen peroxide
HO2-
11,62
Hydrogenphosphate
PO43-
12,32
Hydrogen sulfate
SO42-
1,93
H2O2
H3PO4
Phosphoric acid
H2PO4-
1,96
HPO42-
HF
Hydrofluoric acid
F-
3,14
CH3COO-
4,75
HCO3-
6,46
(H2CO3) Carbonic acid
Dihydrogenphosph.
7,06
HPO42-
NH4+
Chloric acid
CH3COO
Acetic acid
H
Hydrogen sulfide
pKA
HS-
HClO
HSO4-
0
conjugated
base
Name
ClO-
ClO3-
HClO3
strong
conjugated
base
S2-
12,9
H2O
Hydrogensulfidion
water
OH-
15,74
OH-
hydroxide-ion
O2-
24
HS-
For conjugated acidacid-base pairs is:
weak
HClO4
HNO3
weak
Name
moderate
acid
strong
very strong
pKA Table
pK A + pK B = 14
Page 142
71
Degree of Protolysis
(Degree of Dissociation)
Degree of protolysis α = Fraction of dissociated acid in equilibrium.
α=
concentration of dissociated acid
total concentration of acid
=
c( HA)0 − c( HA)
c( HA)0
Degree of protolysis depends on KA value (resp. pKA)
and on initial concentration of acid c0:
α = K A c0
Ostwald Law of Dilution:
Stronger acid or
lower concentration
⇒ more dissociation
(=protolysis
(=protolysis))
Page 143
pH Calculation II
(Weak Acid)
c( H 3O + ) ⋅ c( A− )
KA =
c( HA)
KA =
c ( H 3O + ) 2
c( HA) 0
logarithm ⇒
⇒
since c(H3O)+=c(A-)
and weak acid ⇒
little dissociation: c(HA)
c(HA) ≈ c(HA)0
c( H 3O + ) = K A ⋅ c( HA) 0
pH =
1
2
[ pK A − log c( HA) 0 ]
Example:
0,1 mol acetic acid (pK
(pKA = 4,75) in 1 L water:
pH =
1
2
[4,75 − log 0,1] =
1
2
( 4,75 + 1) = 2,875
Page 144
72
Buffer Solutions
Buffer solutions contain a weak acid and its conjugate weak
base together.
Buffers stabilize the pH value against external acid/base attack.
attack.
Only minor changes of pH by addition of acid or base.
c( H 3O + ) ⋅ c( A− )
KA =
c( HA)
+
⇒ c ( H 3O ) = K A ⋅
c( HA)
c ( A− )
c( A− )
pH = pK A + log
c( HA)
take logarithm ⇒
HendersonHenderson-Hasselbalch equation:
Calculation of pH of buffer solution
For pH < pKA, more acid than conjugate base;
for pH > pKA, more conjugate base than acid.
Page 145
Buffer Effect
Buffer solution neutralizes added
strong acid or base partly, by
shifting the equilibrium.
Most effective buffer,
if c(HA):c(A-) = 1:1
(so that pH = pKA)
Strong acid (H
(H3O+) is converted
into weak acid
H3O+ + A- → HA + H2O
Buffering of base analogous
Excess of acid or base by
a factor of 10 alters
the pH value only by 1!
Molar fraction of base
Page 146
73
Titration Curve
Weak Acid/Strong Base
Example:
20 ml 0.1 molar acetic acid
(pKA=4,75); 0.1 M NaOH added
Initially:
Neutral point:pH = 7<Equivalence pH=8,72
pH = 1 2 [4,75 − log 0,1] = 2,875
By addition of NaOH solution:
HAc + OH- → Ac- + H2O
−
HAc/Ac
HAc/Ac-- pH = 4,75 + log c( Ac )
Buffer:
c( HAc)
Addition NaOH [ml] pHpH-Wert
0
2,875
5
4,27
10
4,75
19
6
19,86
7
20
8,72
>25 as for HCl/NaOH
X
X
Buffer
region
Equivalence p.: all HAc neutralized;
only 0,05 mol/L Ac- (pKB = 9,25)
pH = 14 − 1 2 [9,25 − log 0,05] = 8,72
Page 147
Titration Curve of Phosphoric Acid
Phosphoric acid dissociates in three steps:
H3PO4 + H2O
→
H2PO4– + H3O+ pKS1 = 2,0
H2PO4– + H2O → HPO42– + H3O+ pKS2 = 7,2
HPO42– + H2O → PO43– + H3O+
value
pKS3 = 12,3
Titration curve 10 mL 0.1 M H3PO4
PO43–
E
(not complete)
E
HPO42–
E
H2PO4–
H3PO4
(0,1 M)
Page 148
74
4.2 Salts, Solubility
The Crystalline State
Microscopic characteristics of the crystalline state is:
• 3D regular spatial arrangement (crystal lattice)
• of ions (salts), atoms (metals), or molecules
• with high symmetry and long range order.
order.
Macroscopic crystals: flat surfaces, high symmetry.
Examples:
CaF2
Quartz
(SiO2)
Such single crystals are transparent, but refractive.
Polycrystalline material (µ
(µm size crystals) are opaque.
Page 149
Crystal Lattice
Different types of atomic crystal lattices exist, which are represented
represented
in the external shape of the macroscopic crystal: Cubic, tetragonal,
tetragonal,
hexagonal, monoclinic, triclinic. Dense sphere packing is sought ⇒
highest lattice energy; arrangement of counter ions electric neutral.
neutral.
Crystal type depends on ratio of ion radii r(cation):r((anion).
r(cation):r((anion).
CsClCsCl-type: Cs+:Cl- > 0,7 ⇒
NaClNaCl-type: Na+:Cl- <0,7 ⇒
cubic bodybody-centered lattice;
cubic faceface-centered.
Each Cs+ surrounded by 8 ClEach Na+ surrounded by 6 Cl(and vicevice-versa).
Na+
Cl-
In ZnSZnS-(Zincblende)(Zincblende)-type:
type:
Zn2+:S2-<0,4 ⇒ coordination number 4.
Page 150
75
Amorphous State
The amorphous state is formed by the same molecular entities
as the crystal lattice (ions, molecules), but without long
range order.
order. Also denoted as glassy state.
state.
Amorphous solids have bent fracture surfaces.
Examples:
Examples: Glasses and sediments like flint stone, opal and
Malachite = (Cu(OH)2·CuCO3)
Anisotropy is the properproperty of being directionally
dependent.
E. g. refractive index
Crystals are anisotropic;
Amorphous bodies
are isotropic.
Page 151
Salts
Salts are ionic solids, in which the cations (+) and anions ((-) are
arranged regularly in a 3D crystal lattice.
lattice.
Due to the strong electrostatic attractions between oppositely char
char-ged ions (lattice energy) salts have high melting points (NaCl 800°
800°C).
Salt crystals are electrically neutral, i. e. the charges of cations
cations and
anions are balanced = no net charge.
mAn + + nB m − → Am Bn
In salts of main group elements (I, II, VI, VII) the monoatomic ions
comply with the octet rule ⇒ defined stoichiometry: (NaCl, MgCl2, Na2S)
Salts of auxiliary group elements (transition metals) often have
variable valancy:
FeCl2 = {Fe2+2Cl-} = iron(II)
valancy:
iron(II)-chloride
FeCl3 = {Fe3+3Cl-} = iron(III)
iron(III)-chloride
Cu2O = {2Cu+O2-} = copper(I)
copper(I)-oxide
CuO = {Cu2+O2-} = copper(II)
copper(II)-oxide
Salts can also be composed of polyatomic ions:
AmmoniumAmmonium- NH4+, CarbonateCarbonate- CO32-, SulfateSulfate- SO42-, PhosphatePhosphate- PO432
DichromateDichromate- Cr2O7 , PermanganatePermanganate- MnO4 .
Composition of the salt with according to charge balance; e. g.:
-}
KMnO4={K+MnO4-}, K2Cr2O7={2K+Cr2O72-}; (NH4)2CO3={2NH4+CO32Page
152
76
Salts and Minerals
(Trivial names)
A mineral is a naturally occurring solid chemical NaCl=table salt=
substance formed through biogeochemical prorocksalt=halite
pro- rocksalt=halite
cesses,
cesses, having characteristic composition,
composition,
highly ordered atomic structure,
structure, and specific
physical properties
Common minerals:
Sandstone = sedimented quartz grains (SiO2)
Chalk = calcium carbonate (CaCO3)⇒
Lime stone = Calcium carbonate, too (CaCO3), ⇒
Marble = Calcium carbonate, too (CaCO3)
Gypsum = Calcium sulfate (CaSO4)
Silicates:
Silicates: SiSi-Oxides with Al, Fe, Mg, Na, K, ions.
Sulfidic minerals:
minerals:
Most rockrock-forming minerals
Zincblende (ZnS),
ZnS),
Pyrite (FeS2),
Other important minerals
Galena (PbS
(PbS))
Bauxite: mainly Al(OH)3
Oxides:
Oxides: Corundum (Al2O3); Chile niter: NaNO3
Iron oxides: Hematite (Fe2O3),
Magnetite (Fe3O4), W üstite (FeO)
FeO)
Page 153
Lime Cycle
Page 154
77
Solvation of Salts
Hydrate Shell
Ion crystals can be dissolved preferably in the polar solvent water.
water.
At surface water molecules are attracted by the ions; ions are disdissolved. They are surrounded by a hydrate shell.
shell. (hydration)
Solvation of NaCl in water:
hydrated ions
crystal
Page 155
Solubility Product
The solubility of a certain ionic compound in the respective solvent
solvent is
limited. The solubility of salts quantified by the solubility product.
product.
Heterogeneous equilibrium between saturated solution of the salt
and solid, sedimented salt:
A+ B − → A + + B −
(s)
( aq )
+
( aq )
−
Solubility product: K Sp = c ( A ) ⋅ c( B )
Form of solubility product depends on stoichiometry:
2+
−
2+
−
e. g. CaCl2
Example: A 2 B ( s ) → A ( aq ) + 2 B ( aq )
KSp = c( A2+ ) ⋅ c 2 ( B − )
and for Al2O3 ?
If product of ion concentration < KSp ⇒ unsaturated solution;
If product of ion concentration > as KSp ions crystallize as salt.
The solubility product is temperature dependent.
(like all equilibrium constants)
Page 156
78
Solubility Products
Poor Solubility Salts
Low solubility: KSp < 10-4 mol2/L2
Compound
KSp
Halogenides
Compound
Sulfates
10-10 mol2/l2 CaSO4
AgCl
KSp
mol2/l2
2·10-5 mol2/l2
AgBr
5·10-13
BaSO4
1·10-9 mol2/l2
AgI
8·10-17 mol2/l2 PbSO4
1·10-8 mol2/l2
Sulfides
Carbonates
CaCO3
5·10-9
HgS
10-54 mol2/l2
BaCO3
2·10-9 mol2/l2 CuS
10-36 mol2/l2
CdS
10-28 mol2/l2
mol2/l2
Hydroxides
Mg(OH)2
10-11
mol3/l3
PbS
10-28 mol2/l2
Al(OH)3
10-33 mol4/l4 ZnS
10-22 mol2/l2
Fe(OH)2
10-15
mol3/l3
NiS
10-21 mol2/l2
Fe(OH)3
10-38
mol4/l4
MnS
10-15 mol2/l2
Cr(OH)3
10-30 mol4/l4 Ag2S
10-50 mol3/l3
Hints:
Hints:
• alkali metal salts
have good solubility
• alkali earth salts
have lower solubility
• Heavy metal salts
(Ag, Pb,
Pb, Ba,
Ba, Cd)
Cd)
poorly soluble
• Sulfides (S2-) are
poorly soluble
Page 157
Precipitation Reactions
Silver nitrate (AgNO3) is the best soluble silver salt.
What happens, if a sodium chloride solution is added
to a saturated solution of silver nitrate?
Ag (+aq ) + NO3−( aq ) + Na(+aq ) + Cl(−aq ) → ?
The solubility product of silver chloride (10-10) is exceeded
and the salt AgCl precipitates.
{
}
Ag (+aq ) + NO3−( aq ) + Na(+aq ) + Cl(−aq ) → Na(+aq ) + NO3−( aq ) + Ag + Cl − ↓
Page 158
79
The Solubility
Due to the different unit dimension KSp values are hard to compare.
Therefore, definition of the solubility S of a salt:
For
A+ B − ( s ) → A+ ( aq ) + B − ( aq )
Example:
Example: KSp (AgCL)=10
AgCL)=10-10 mol2/L2
S = c( A+ ) = c( B − ) = K Sp
LAgCl = 10 −10 = 10 −5 mol / L
c( Ag + ) = c(Cl − ) = c( AgCl ) = 10 −5 mol / L
m = M ⋅ n = 143,5 g / mol ⋅10 −5 mol = 1,435mg
For the general case of a salt AxBy:
c( A y + ) c( B x − ) x + y K Sp
S=
=
=
x
y
xx ⋅ y y
Page 159
Temperature Dependence
of the Solubility
The dissociation of ions from the crystal is a highly endothermic
endothermic
(energy consuming) process .
In contrast, energy is released when the ion is solvated.
Lattice energy
- Solvation energy
= Solution energy
Energy
Solution energy > 0 ⇒
endothermic process
Solution energy < 0 ⇒
exothermic process.
Solvation
endothermic
solution
exothermic
Most solution processes are endothermic
⇒ improved solubility at higher temperatures.
←
A+ B − ( s ) + Energie →
A+ ( aq ) + B − ( aq )
Page 160
80
CommonCommon-Ion Effect
The commoncommon-ion effect is the decrease of the solubility of one salt,
when another salt, which has an ion in common with it, is also present.
present.
For example, the solubility of silver chloride,
chloride, AgCl,
AgCl, is lowered when
sodium chloride, a source of the common ion chloride, is added.
c( Ag + ) = c(Cl − ) = 10 −5 mol / L
Example:
Example: An amount of NaClNaCl-solution is added to a saturated AgClAgCl3
solution, so that c(Cl )=10 mol/L.
mol/L. The solubility product of AgCl is
exceeded; AgCl has to precipitate until: K Sp = c( Ag + ) ⋅ c(Cl − ) = 10−10 mol / L
The solubility product remains constant!
The concentration of silver ions
c(Ag+) ≠ c(Cl-) is then:
c( Ag + ) =
K Sp
c(Cl − )
=
10 −10
= 10 −7 mol / L
10 −3
Page 161
4.3 Redox Reactions
Basics, Definitions
Redox reactions are essential for chemical energy
Fuelstorage and conversion:
cell
Both in
Battery
technology:
and
in nature:
Combustion
alcoholic fermentation
Photosynthesis
Page 162
81
Origin of the Terms
Oxidation/Reduction
Originally the term oxidation
oxidation described the reaction
of organic materials or metals with atmospheric oxygen;
i. e. the combustion of
or the rust of iron
wood, oil, wax or coal
4 Fe + 3 O2
→ 2 Fe2O3
C + O2 → CO2
Reduction was the recovery of (precious) metals
2 HgO → 2 Hg + O2
from their respective salts e. g.:
Since there are many similar processes, which don't involve oxygen,
oxygen,
the term oxidation nowadays is used in a more general concept.
Page 163
Today's Definition of
Oxidation/Reduction
Oxidation is the loss of electrons by an atom, ion or molecule
⇒ increase of oxidation number (state)
Oxidation = Electron donation
A transfer of electrons changes material properties substantially:
substantially:
E.g.: Copper (red shiny metal) is transformed into Cu2+ ions.
Cu2+-Ions form together with anions salts or are solvated.
Solvated Cu2+ ions have a hydrate sheath and blue color.
This process
is reversible
Reduction is the gain of electrons by an atom, ion or molecule
⇒ decrease of oxidation number (state)
Reduction = Electron acceptance:
Page 164
82
Redox Reaction
Thermite Reaction
Iron oxide powder mixed with aluminum powder.
Reaction initiated with sparkler or blowpipe (activation energy!).
energy!).
Exothermic
reaction!
Energy released
as heat and light.
How can you prove that the
produced metal is iron,
not aluminum?
Thermite
process used for
welding rails
Page 165
Redox Reactions
Examples
No free electrons exist! Oxidation and reduction coexistent!
Redox reaction = Electron transfer:
Depending on the reactants several electrons can be transferred
In redox reactions the gain of electrons in the reduction must be
be
identical to the loss of electrons in the oxidation.
Electrons never show up in the total equation!
Example: Zinc + sulfur: Oxidation:
Zn → Zn2+ + 2eReduction: S + 2 e-→ S2Redox reaction: Zn + S → Zn2+ + S2- = ZnS
A redox reaction can formally be subdivided into oxidation and
reduction. The two parts reaction never occur separately!
Conservation of mass and charge: On both sides of reaction arror
same number of the atoms of each element and same total charge.
Note: Do not subtract electrons
Zn - 2e- → Zn2+
No half molecules
:
½ Cl 2
Page 166
83
Balancing of Redox Equations
Aluminum reacted with bromine
yielding aluminum bromide:
Interim eq.: Al + Br2 → AlBr3
Bromine diatomic,
Aluminum ion 3 valent,
Bromide 1 valent
+ III
3+
0
Al → Al
Oxidation ⇒ oxidation number increases!
Reduction ⇒ oxidation number decreases!
+ 3e −
0
−I
Br 2 + 2 e − → 2 Br
−
Lost = gained electrons: Ox (·
(·2); Red (·3)
2 Al → 2 Al 3 + + 6 e −
3 Br 2 + 6 e − → 6 Br −
2 Al + 3 Br2 → 2 Al 3 + + 6 Br − = 2 AlBr3
Total equation by
addition of partial
reaction equations;
e- eliminated!
Page 167
More Redox Reactions
Reduction of ironiron-(II)(II)-oxide to iron:
Redox : 2 FeO + C → CO2 + 2 Fe
Oxidation : C + 2 O2→ CO2 + 4 e2+
Reduction: 2 Fe + 4 e → 2 Fe
Redox reaction without oxygen:
Redox:
2 Na + Cl2 → 2 NaCl
Oxidation:
2·(Na → Na+ + e-)
Reduction: Cl2 + 2 e- → 2 Cl-
Na (I. mg) looses 1 e-,
Cl (VII. mg) gains 1 e-,
to obey octet rule.
Chlorine diatomic gas.
Redox reaction with complex stoichiometry:
Redox: 16 Al + 3 S8→8 Al2S3={2Al3+3S2-}
Oxidation:
16·(Al → Al3+ + 3 e-)
Reduction: 3·(S8+16 e-→8S2-)
Al (III. mg) looses 3 e-,
S (VI. mg) gains 2 e-,
to obey octet rule.
sulfur as S8 ring.
The stoichiometric factors result from combination
of the partial reaction steps.
Page 168
84
Oxyhydrogen Reaction
Many redox reaction are exothermic, i. e. release of energy.
Another reaction this kind the oxyhydrogen reaction (hydrogen test)
test)
Biggest hydrogenhydrogen-oxygen reaction
of all times?
2 H2 + O2 → 2 H2O + energy
Zeppelin "Hindenburg"
Lakehurst 1937
Oxidation:
H2 → 2 H+ + 2 eReduction:O2 + 4 e- → 2 O2Redox: 2 H2 + O2 → (4H+ + 2O2-)→ 2H2O
H (I. mg) losses 1 e-,
O (VI. mg) gains 2 e-,
⇒ noble gas configuration.
configuration.
Both diatomic gases!
Page 169
Oxidizer/Reducer
Compounds which are able to oxidize other species (eliminate electrons)
electrons)
are called oxidizers (oxidants, oxidizing agents) = electron acceptors
The oxidizer itself is being reduced!
Common oxidizers:
O2 + 4 e-→ 2 O2• Oxygen
Cl2 + 2 e-→ 2 Cl• Chlorine
• Oxoanions z.B.:
z.B.: MnO4- (Permanganate), Cr2O72- (Dichromate)
• Anions of halogen oxoacids, z.B.
z.B. ClO4- (Perchlorate)
+
• Noble metal ions, e. g. Ag + e-→ Ag
Compounds which are able to reduce other species (donate electrons)
electrons)
are called reducers (reductants, reducing agents) = electron donators
The reducer itself is being oxidized!
Common reducers:
H2 → 2 H+ + 2 e• Hydrogen
• Zn, Fe, Mg
Zn → Zn2+ + 2e• Carbon
Page 170
85
Oxidation Numbers/State
How do you know, its a redox reaction
and which species gains or looses how many electrons?
Redox reaction ⇒ change of oxidation numbers
Oxidation number = hypothetical charge an atom would have
if all bonds to atoms of different elements were 100% ionic.
Examples:
Rules (order of priority) agreedagreed-upon for determ.
determ.
of oxidation state (roman numeral superscript):
0 0
0
0
H
,
Cl
,
Na,
Al
1. Free elements always have oxidation number 0
2
2
+I
+II
-II
Na+, Ca2+, S2-
2. Monoatomic ions: OxNo = charge
3. H OxNo = +I ; O OxNo = -II
-I +I -I
(exception: Hydrides MeHx; H2O2)
-IV+I +IV+IV-II
4. In molecules ΣOxNo=0
OxNo=0 ;
in molecular ions ΣOxNo=
OxNo= charge
CH4, CO2
+VII+VII-II, +VI+VI-II
MnO4-, SO42-
Page 171
Oxidation Numbers
of Selected Compounds
Compounds of chlorine with various oxidation numbers:
OxNo
-I
+I
ClOHClO
Chloride Chlori Hypochlorite
ne
hypochloric acid
Formula Cl-
Name
0
Cl2
+IV
ClO2
ChloroChlorodioxide
+V
ClO3HClO3
Chlorate
+VII
ClO4HClO4
Perchlorate
Chloric acid Perchloric acid
Compounds of sulfur with various oxidation numbers:
OxNo
-II
Formula S2-/H2S
Name
0
S8
Sulfide/Hydro Sulfur
gen sulfide
+III
S2O42Dithionite
+IV
SO2/SO32-/
H2SO3
Sulfur dioxide
sulfurous ac.
+VI
SO3/SO42-/
H2SO4
S.trioxide/Sulfate
S.trioxide/Sulfate
sulfuric acid
Compounds and oxidation numbers of manganese:
OxNo
+II
+IV
+VI
+VII
Formula
Mn2+
MnO2
MnO42MnO4Manganese
Manganese Manganate
Name
Permanganate
dioxide
(II) ion
(blue)
(violet)
Page 172
86
Redox Reactions and pHpH-Value
In some redox reactions (often in aqueous solution)
H3O+ resp. OH- ions for the charge compensation
under consideration of the acidic or alkaline/basic environment
E.g.:
E.g.: Permanganate reduced to manganese(II)
manganese(II) in acidic environment
+VII
+II
For compensation fo charges (left -6;
right +2) 8 H3O+ are introduced
(acidic environment).
2+
2+
MnO4 + 5 e + 8 H3O → Mn + 12 H2O Together with 4 O
(formally) they form 12 H2O
MnO4- + 5 e-
→ Mn2+
E.g.:
E.g.: Hydrogen peroxide is oxidized to oxygen in alkaline.
alkaline.
-I
H2O2
0
→ O2 + 2 e- +2 H+
H2O2 + 2 OH- → O2 + 2 e- +2 H2O
Note: There are
no free protons!
Both examples only partial reaction equations!
Page 173
Redox Titration
Example: Permanganometry
Redox reactions can be used in quantitative analysis (titration) for
determination of an unknown concentration of a solution,
if the equivalence point is detectable, e. g. change of color
Permanganometry:
Permanganometry: Titration of a solution of an oxidizable compound
(e. g. Fe2+) with potassium permanganate solution (KMnO4; violet);
The Fe2+ are oxidized by the drop wise added MnO4- ions gradually
yielding Fe3+ ions. At the same time the violet MnO4- ions are reduced
to colorless Mn2+ ions ⇒ decolorization of the added permanganate
permanganate
solution as long as Fe2+ ions are present.
After equivalency (Fe2+ ions consumed) no reaction of added MnO4ions⇒
ions⇒ violet coloring of the iron ion solution.
From volume and concentration of consumed permanganate solution
the concentration of Fe2+ solution can be calculated.
Account for stoichiometry! 1 mol MnO4- oxidizes 5 mol Fe2+.
Page 174
87
Chromatometry, Iodometry
Two other analytic methods like permanganometry (titrations)
but with different stoichiometry:
Chromatometry (dichromate titration):
titration):
22+
Cr2O7 + 6 Fe + 14 H3O+ → 6 Fe3+ + 2 Cr3+ + 21 H2O
Dichromate
Iodometry:
Iodometry: determination of concentration by means of redox system
2 I- ⇒ I2 + 2 e-: Elemental iodine forms with starch solution a
deep blue complex.
complex.
Reversed titration method:
Example:
Example: determination c(Cu2+):
- Measure off defined volume (e. g. 20 ml)
- Add excess of KI solution ⇒ 2 Cu2+ + 2I- → 2 Cu+ + I2
- 2 mol Cu2+ produce 1 mol I2, which dyes the solution blue
- Determination of the molar of produced iodine by redox
titration with reductant solution of known concentration.
- I2+2S2O32- → 2I-+S4O62- (2 mol thiosulfate reduces 1 mol iodine)
- From vol. and conc. of used S2O32- ⇒n(I2)⇒n(Cu2+)⇒c(Cu
c(Cu2+)
Page 175
Noble/Base Metals
Why does iron rust, but gold does not? Does aluminum oxidize?
What is the difference between noble and base metals?
In general all metals can react by loosing electrons,
forming metal kations, i. e. they oxidize. But:
The various metals have a different oxidizability:
Noble metals:
metals: resistant to oxidation, low reactivity, occur in
elemental form. Gold, platinum, silver, (copper)
Base metals:
metals: readily oxidizable, reactive, occur in compounds
all mainmain-group metals:
(Alkali, alkaline earth
H
metals, aluminum)
Li Be
B
Transition metals:
Iron, Zinc, etc.
Na Mg
Al
K Ca Sc
Ti V Cr Mn Fe Co Ni Cu Zn Ga
Rb Sr Y
Zr Nb Mo Tc Ru Rh Pd Ag Cd In
Cs Ba La * Hf Ta W Re Os Ir Pt Au Hg Tl
Page 176
88
Chemical Reactivity
(Oxidizability) of Alkali Metals
Alkali metals have highest reactivity.
Storage in inert media (ether) or
paraffin oil. Otherwise intense
reactions, like
4 Na + O2 → 2 Na2O
Sodium under paraffin oil
Water not feasible as storage medium,
since sodium (like all alkali metals) reacts with water.
2 Na + 2 H 2O → 2 Na(+aq ) + 2OH (−aq ) + H 2 ↑
Sodium hydroxide is produced
and hydrogen, which is inflamed.
The sodium melts (Tm=98°
=98°C)
due to reaction heat
Page 177
Redox Reactions
of Noble and Base Metals
Cu2+
Zn+2 HCl →
Zn2++2Cl-+H2↑
Cu+2HCl →
Zn2+
Zn + Cu2+→Zn2++Cu
Why?
Why?
Page 178
89
The oxidizability of metals, i. e. their tendency
to loose electrons, can be compared by direct
reaction with each other ⇒ ranking
+
+
Examples: Ag + Au+ → Ag 2++ Au
increasing reducing reaction
Gold
Platinum
Palladium
Mercury
Silver
Copper
Lead
Tin
Nickel
Cobalt
Cadmium
Iron
Chromium
Zinc
Aluminum
Magnesium
Sodium
Calcium
Potassium
Lithium
decreasing oxidizability
reactive
noble
Table of Noble/Base Metals
Cu + 2 Ag → Cu + 2 Ag
Ni + Cu2+→ Ni2++Cu
Fe + Ni2+ → Fe2+ + Ni
Zn + Fe2+→ Zn2+ + Fe
but: Cu + Zn2+ →
Driving force for the oxidation of base metals
by noble metal kations is: ∆G<O
How can the tendency of electron
attraction or repulsion be quantified?
How can the energy released in redox
reactions be utilized?
Page 179
Conversion of Energy
Chemical energy
Heat
z.B.:
z.B.: C + O2 → CO2
Oder: Zn + Cu2+→Zn2++Cu
Electric motor
Electric energy
Mechanical energy
Dynamo
Battery
Chemical energy
Electrical energy
Spatial separation of oxidation (Zn → Zn2++ 2e-) and
reduction (Cu
(Cu2++ 2e- → Cu) reaction. Instead of direct transfer from
oxidation to reduction site. electrons have to flow through conductor.
conductor.
Accumulator
Chemical energy
Electrical energy
Page 180
90
4.4 Electrochemistry:
Redox Potential; Galvanic (Voltaic) Cell
Between a metal bar (electrode) and the surrounding metal ion solution
solution
(electrolyte) charges are exchanged by chemical reactions building
building up a
difference in the chemical potential. More noble metals have a lower
lower
tendency to loose electrons. If two such halfhalf-cell are connected by a
conductor the potential difference ∆E can be measure as voltage.
voltage.
∆E=1,1 V
Half-cells
more
noble
less
noble
⊕
⊕
⊕
⊕
⊕
⊕
⊕
⊕
⊕
⊕
Page 181
Table of Standard Electrode Potenials
FaradayFaraday-Konst.
Konst. F
increasing oxidation activity
∆ R G 0 = − ne ⋅ F ⋅ ∆ E 0
Base metals oxidized by acid (H+⇒H2
noble metals and copper are not.
∆E0
F2
ne
+ 2e-
Red.
Fluor (F)
↔ 2F-
+2,87 V
Gold (Au)
Au+
+ e-
↔ Au
+1,69 V
Platinum (Pt)
Pt2+
+ 2e-
↔ Pt
+1,20 V
Cl2
+ 2e-
↔ 2Cl-
+1,31 V
Ox.
Chlorine (Cl)
oxygen (O2)
O2+4H+ + 4e- ↔2 H2O +0,85 V
increasing reducing activity
Strength oxidizing/reducing agents
depends on change of Gibbs free enerenergy ∆G during electron loss or gain:
Each redox system (electrode/
electrolyte ⇒ redox potential;
potential;
listed in Table of Standard Electrode
Potential.
Potential.
Isolated potential (half(half-cell)= not
obtained emp. ⇒ only galvanic cell
potial between pairs of electrodes ∆E.
E0 Hydrogen Electrode defined 0
Stronger oxidizers (Ox)
⇒ positive potential ∆E0
Stronger reduction agents (Red)
⇒ negative potential ∆E0
Stand.cond.(25°
(25°C,1013 hPa,1 mol/L).
mol/L).
Silver (Ag)
Ag+
+ e-
↔ Ag
+0,80 V
Sulfur (S)
S
+ 2e-
↔ S2-
+0,48 V
Copper (Cu)
Cu2+
+ 2e-
↔ Cu
+0,35 V
hydrogen (H2)
2H+
+ 2e-
↔ H2
0
Lead (Pb)
Pb2+
+ 2e-
↔ Pb
-0,13 V
Tin (Sn)
Sn2+
+ 2e-
↔ Sn
-0,14 V
Nickel (Ni)
Ni2+
+ 2e-
↔ Ni
-0,23 V
Cadmium (Cd)
Cd2+
+ 2e-
↔ Cd
-0,40 V
Iron (Fe)
Fe2+
+ 2e-
↔ Fe
-0,41 V
Zinc (Zn)
Zn2+
+ 2e-
↔ Zn
-0,76 V
Aluminum (Al)
Al3+
+ 3e-
↔ Al
-1,66 V
+ 2e-
↔ Mg
-1,66 V
+ 1e-
↔ Na
-2,71 V
+ 1e-
↔ Li
-3,02 V Page 182
Magnesium (Mg) Mg2+
Sodium (Al)
Na
Lithium (Li)
Li+
+
91
Standard Hydrogen Electrode (SHE)
Determination of standard electrode potentials in comparison
to Standard Hydrogen Electrode (set to ∆E0=0 V)
∆E0=0 Volt
Salt bridge
Page 183
Galvanic Cell / Voltaic Element
Electromotive Force (EMF)
Combination of two halbhalb-cells is a Galvanic Cell
∆E0= ∆E0 (cathode) -∆E0 (Anode) = EMF (Electromotive Force)
Daniell cell
Anode
cathode
0.35-(-0.76)V
metallic
copper
Zink electrode
dissolved
precipitated
Page 184
92
Nernst Equation
Concentration Dependence of Electrode Potential
The potential of a real electrode depends on the standard potential
potential of the
redox system (electrode/electrolyte) an on the conditions (concentration,
(concentration,
pressure, temperature).
Under standard condition (c = 1mol/L solution,1013 hPa,
hPa, 25°
25°C)
the standard potential E0 is listed in the table.
For higher or lower concentration of the electrolyte the electrode
electrode potential
of the halfhalf-cell E can be calculated by the Nernst equation:
E = E0 +
0,059
⋅ log c( Me n + )
n
Nernst
equation
Example:
Example: What is the potential of a copper electrode
in a Cu2+-ion solution of c = 0,1 mol/L?
E = 0,35 +
Walter Hermann Nernst
0,059
0,059
⋅ log 10 −1 = 0,35 +
⋅ (−1) = 0,3205V
2
2
Page 185
Corrosion,
Corrosion, Passivation
Corrosion:
Corrosion: Disintegration of an engineered material
due to chemical reaction with its surrounding (oxygen, acid)
2 Me + O2 + 2 H 2O → 2 Me 2+ + 4OH −
2 Me + 2 H 3O + → 2 Me 2+ + 2 H 2O + H 2 ↑
Basic metals oxidized by atmospheric oxygen ⇒ few nm oxide layer.
Formation of hard, nonnon-reactive surface film inhibits further corrosion
(e.g. Al2O3 on Al, ZnO on Zn, TiO2 on Ti),⇒
Ti),⇒ Passivation)
Passivation)
Iron, however, forms a porous oxide layer (rust). oxygen can penetrate
penetrate
rust layer until complete corrosion (Fe2O3, Fe3O4) (Pitting
).
(Pitting).
In hothot-dip galvanizing process iron or steel are coated with a zinc layer
layer
by passing metal through molten zinc bath (460°
(460°C). Exposed to atmoatmosphere, zinc reacts with O2 to form ZnO and further
with CO2 to ZnCO3, a dull grey, fairly strong material
protecting the iron below against further corrosion
(passivation). If the coating is damaged, zinc works
as a galvanic anode so that iron is not oxidized.
Page 186
93
Galvanic Corrosion
Galvanic (Sacrificial)
Sacrificial) Anode
Galvanic Corrosion:
Corrosion:
ShortShort-circuited galvanic element;
Direct connection with a more noble
metal enhances oxidation (corrosion)
of the more basic metal, since electrons
are deflected. Contact between gold and
amalgam fillings should be avoided,
otherwise mercury and tin are dissolved.
Hg,Ag,Sn
Galvanic anode:
anode: Basic metal
(Zn, Mg, Al, alloy) in direct contact
protects metal (Fe) from corrosion;
itself being dissolved (sacrificed)
In ships, pipes, etc.
Iron nail in saline with indicator
(with Fe3+ green dye):
Galvanic corrosion (contact with nobler
Cu) enhances iron dissolution,
Galvanic anode from Zn stops
corrosion. Zn2+(aq) colorless
Fe/Cu
Fe
Fe/Zn
Page 187
Batteries:
Batteries:
ZincZinc-Carbon Battery (Dry Cell)
Cell)
Anode (Minus pole, Zinc case):
Zn → Zn2+ + 2eCathode (Plus pole, carbon rod):
+III
MnO2+ H++ 1e- → MnO(OH)
MnO(OH)
Manganese dioxide → ManganoxidManganoxidhydroxid
Regeneration of H+ from the
Ammonium chloride electrolyte:
NH4+ → NH3+ H+
Zinc ions bind produced ammonia
as complex:
Zn2++ 2NH3 → [Zn(NH3)2]2+
This complex forms with the chloride
ions (ammonium chloride) a salt
[Zn(NH3)2]2+ + 2Cl- →[Zn(NH3)2]Cl2(s)
electrolyte
paste
Manganese
dioxide
Cathode
Graphite
rod
+IV
zinc case
anode
Total reaction:
reaction:
Zn + 2MnO
2MnO2 +2NH
+2NH4Cl → 2MnO(OH) + [Zn(NH3)2]Cl2
all solid!
Page 188
94
LeadLead-Acid Battery
Lead
Lead electr.
electr.
electrode
Lead
(minus pole) dioxide
coating
(plus pole)
DISCHARGING
dissociated sulfuric acid
The chemical reactions are (charged to discharged):
Minus pol:
Plus pol:
0
+IV
discharge+II
= comproportionation
charging = disproportionation
For car: 12 Volt
stacking!
Page 189
Fuel Cell
Idea:
Idea: 1838 Christian Schö
Schönbein; since 1950 in space technology.
Electrodes separated by
Theor. max.1,23 V: pract. 0,5-1 V ⇒ Stacking
membrane or electrolyte.
electrolyte.
Anode flushed with fuel
(H2, methane,
methane, ethanol,
ethanol,
DC
glucose solution),
solution),
cathode with oxidizing
agent (O2, H2O2).
Electrodes made of
carbon nano tubes
air
with PtPt- or PdPd-coating
as catalyst.
catalyst.
Typs:
Alkaline Fuel Cell
H2/ O2/OH- (100kW)
Direct Methanol Fuel Cell
MeOH/Luft/Polymer/H+ (500 kW)
Membrane
Cathode
Page 190
95
Electrolysis
Processes in voltaic cells can be
reversible: Applying DCDC-voltage at
electrodes leads to electrolysis:
electrolysis:
Elements in the electolyte are separated
by enforced redox reaction.
reaction.
Voltaic cell
spontaneous
Electrolysis
nonnon-spontaneous
Ion migration due to electrostatic attraction:
attraction:
Positive Cations (Cu2+)→minus pole (cathode
(cathode))
Negative Anions (Cl-) → plus pole (anode
(anode))
Cathode: Cu2+ + 2 e– → Cu
Reduction
Oxidation
→ Cl2 + 2
Note:
Note: Polarity of cathode/anode reversed with
→ minus pole
respect to voltaic element:
element: Plus ←
Anode:
e–
2 Cl–
Oxidation always at the anode!
anode!
Electrolysis of copper chloride solution
Page 191
Electrolysis of Water
faucet
faucet
oxygen
Hoffman
electrolysis
apparatus
hydrogen
acidic
2 H2O→ 1O2 + 4 H+ + 4 eanode
alkaline
4 H2O + 4 e- → 2 H2 + 4 OHplatinum electrode (cathode)
cathode)
15 Volt DC !
Page 192
96
Electrolytic Copper Purification
Crude copper anode
(99,5% pure) dissolved
at 0.2-0.3 V at anode.
Highly pure copper
(99,95%) precipitated
at cathode.
Impurities of base
metals (Fe, Zn) remain
in solution (as ions);
noble metal particles
are deposited on the
ground (anode sludge).
Anode
Cathode
Crude copper Pure copper
Copper sulfate sol.
Anode sludge
Page 193
Aluminum Production
HallHall-Heroult Process
Aluminum ore (bauxite) heated in
conc. NaOH sol. 200°
200°C (Bayer proc.).
Aluminum hydroxide ⇒ sol. Aluminate
Al(OH)3 + NaOH Na+ + [Al(OH)4]Filtering of solid impurities;
Neutralisation → pure Al(OH)3 precip.;
precip.;
Calcination (1000°
(1000°C) ⇒ Alumina.
2 Al(OH)3 Al2O3 + 3 H2O
Alumina dissolved in molten cryolite
(Na3AlF6) at 1000°
1000°C in a ceramic vat;
electroelectro-chem.
chem. reduction to Aluminum
5 Volt DC, some 100,000 A current!
Anodes of fused coke are consumed
by oxidation:
2Al2O3 + 3C → 4Al + 3CO2;
∆H = + 0,835 MegaJ/mol
MegaJ/mol Al
Extreme energy consumption!
Page 194
97