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DISPERSED SYSTEMS
Ingrid Žitňanová
DISPERSED SYSTEMS
Dispersed
phase
(water)Dispersionmedium
(oil)
SOLUTE
(DISPERSED PHASE )
SOLVENT
(DISPERSION MEDIUM )
Solute (NaCl) Solvent (water)
It has a non-uniform composition
There are two or more phases
They can be separated out physically
It has a uniform composition
It has only one phase
It can’t be separated out physically
Classification of the dispersed systems
according to the diameter of dispersed particles
1. Analytical (molecular, true solutions)
2. Colloids
3. Coarse / Crude dispersion (suspension)
< 1 nm
1 – 1000 nm
> 1000 nm
SolutionColloidsTrue solution Coarse dispersion
particle sizeType of dispersion
Properties of the dispersed systems
Dispersion Molecular (true solut.) Colloidal Coarse (crude)
Particles size 1 nm 1 – 1000 nm > 1000 nm (1 μm)
Particles filterability Cannot be separated by
filtration
Can be separated by
semipermeabile
membrane
Can be separated by
filtration
Diffusion of particles rapid slow No diffusion
Visibility of particles Not visible under the
electrone microscope
Can be visible under
the electrone
microscope
Can be seen under
the low power
microscope or eye
Sedimentation of
particles
Particals do not sediment Sediment in the
strong centrifugal
field
Sediment under the
influence of gravity
Optical properties Transparent
No Tyndall effect
Tyndall effect Not transparent
Tyndall´s effect
is due to the scattering of light by colloidal particles, while showing
no light in a true solution.
This effect is used to determine whether a mixture is a true solution
or a colloid.
True
solution
Colloidal
solution
• when light is passed
through a colloidal
solution, the substance
in the dispersed phases
scatters the light in all
directions, making it
readily seen
TRUE SOLUTIONS(Analytical, molecular solutions)
TRUE SOLUTION
• a homogeneous mixture of two or more components
• particle size 1 nm
Liquid (vinegar)
Gas (carbon dioxide)
Solid (sugar)
e.g. water Acetic acid in water
CO2 in water
Sugar water
+
solvent
SolventsPolar
Nonpolar
Solutes
Polar
Nonpolar
• Polar solutes dissolve well in polar solvents
• Nonpolar solutes dissolve well in non-polar solvents
– e.g.water, ethanol, methanol,
– e.g. chloroform, hexane, benzene
- glucose, acetic acid, NaCl
- fats, steroids, waxes
Oil in water
Electrolytes, Nonelectrolytes
In water,
Strong electrolytes separate into ions making solutions that conduct electricity
Weak electrolytes produce a few ions
Nonelectrolytes produce molecules, not ions, do not conduct electricity
Electrolyte – when dissolved in water separates into cations and anions,
which disperse uniformly through the solvent.
Strong electrolytes
are compounds with ionic or very polar covalent bond
strong electrolyte
when dissolved in water, they dissociate 100% . They break up
into positive and negative ions in water
produced ions conduct an electrical current
Examples: KOH, HCl, HNO3, H2SO4...
Solutes that are weak electrolytes
Weak electrolytes
weak electrolyte
dissolve in water forming a few ions
produce solutions that conduct electricity weakly
Examples: HF, acetic acid, lactic acid, ammonia...
Nonelectrolytes
Solutes - nonelectrolytes are covalent compounds which:
nonelectrolyte
do not produce ions in water
form solutions that do not conduct an electrical current
Examples: sucrose, glucose, urea, ethanol, glycerol...
Average ion concentrations in blood plasma, ISF and ICF (mmol/L)
Electrolytes in body fluids
ICF – intracellular fluid
* Most of them are organic phosphates (hexose-P , creatine-P , nucleotides, nucleic acids)
ISF – interstitial fluid - the fluid in spaces between the tissue cells
Ionic composition of body fluids
Blood plasma and ISF (interstitial fluid) have almost identical
composition, ISF does not contain proteins
The main ions of blood plasma are Na+ and Cl-, responsible for
osmotic properties of ECF (extracellular fluid)
The main ions of ICF (intracellular fluid) are K+, organic
phosphates and proteins
Each body fluid is electroneutral total positive charge = total
negative charge
Interstitial
fluid
Water
Intracellular fluid ICF – inside cells – 25 - 30L
Extracellular fluid ECF – 15L - blood plasma, intersticial fluid,
lymph, fluid in gastrointestinal tract, urine...
Volume of water in body is balanced (intake = output in urine, feces,
sweating, lungs)
Central regulatory organ of water volume – kidneys
Hydrogen
bond
Water – H2O – a polar solvent
O
O
O
O
HH
H
H
H
H
H
H
Average total body water (TBW) as body weight percentage
The water content of the body changes with:
• gender
• age
• body composition (a lean person has a higher TBW than an obese person)
True solutions
Ionic Molecular
• solutions of nonelectrolytes
• contain molecules of compounds in
solution (glucose in water, urea)
• solution of electrolytes in which ions
are present, formed by electrolytic
dissociation of ionic compounds
H2O
H2O
H2O
Hydrated
ions
H2O
NaCl Na+ +
H2OCl
-
Electrolytic
dissociation
Ionic strength ( I )
is the concentration of ions in the solution
i – number of particles
ci - the molar concentration (mol/L)
zi – charge of the particle
Only ionized species contribute to ionic strength in the
solution!!!
Example 2:
Calculate ionic strength of a solution containing 0.02 mol/L
Na2SO4 and 0.1 mol/L glucose.
I1 = 0.5 [(2 x 0.02 × 12 ) + (1 x 0.02 × 22 )] = 0.06 mol/L
1. Na2SO4 = 2Na+ + SO42-
2. Glucose0 no dissociation
I2 = 0.5 x 1 x 0.1 x 02 = 0 mol/L
I = I1 + I2 = 0.06 + 0 = 0.06 mol/L
SO42-2Na+
Solubility
A measure of how much of a solute can be dissolved in a solvent
Saturated Solutions - contain the maximum concentration of a
solute dissolved in the solvent (under a given set of pressure and
temperature). Additional solute will not dissolve in a saturated
solution
Super Saturated Solutions contains more dissolved solute
than could ordinarily dissolve into the solvent. Undissolved
solid remains in the flask.
Unsaturated Solutions – a solution containing less than the
maximum concentration of solute that will dissolve under a given
set of conditions (more solute can dissolve).
Unsaturated Saturated
Super saturated
• The rate of dissolution affects how fast a drug is absorbed in the body.
Clinical significance of solubility
• Aqueous solubility is often considered when formulating drugs.
• Drugs (for oral administration) with low aqueous solubility may have low
bioavailability causing the drug to be not as effective.
Factors affecting solubility
• Temperature
• Pressure
• Polarity
• Concentration of the solute
Solubility
For endothermic reaction (requires energy from its surroundings)
Solubility increases when solution temperature increases
For exothermic reaction (releases energy in the form
of heat)
Solubility is reduced when solution temperature
increases (NH3)
Effect of temperature on solubility
Example:
NH4NO3 used in first-aid cold packs.
Its dissolving in solution is an endothermic reaction, heat
energy is absorbed from the environment. This causes the
surrounding environment to feel cold.
Temperature
For gases
Higher temperature reduces solubility of gases –
it drives gases out of solution
Examples:
Carbonated soft drinks are more bubbly if stored in the
refrigerator (more CO2 is inside the drink)
Warm lakes have less O2 dissolved in them than cool lakes
Higher
temperature
Higher
kinetic E of
gas particles
Breakage of intermolecular
bonds between the gas solute
and solvent
Pressure
• little effect on solubility of solids and liquids
• will greatly increase solubility of gases
• Henry's Law: The solubility of a gas in a liquid is directly proportional
to the pressure of that gas above the surface of the solution.
Clinical significance of pressure on solubility
Decompression Sickness
• scuba divers in deep water → ↑ the pressure in their body → nitrogen in their
body dissolves in their blood
• scuba divers ascend to the surface too quickly → the sudden drop in pressure → nitrogen
bubbles come out of solution → painful and potentially fatal gas embolisms
Polar substances tend to dissolve in polar solvents.
Nonpolar substances tend to dissolve in nonpolar solvents.
Examples
Polarity
Vitamin A is soluble in nonpolar compounds (e.g. fats)
Vitamin C is soluble in water
Vitamin A Vitamin C
Properties of true solutions
Colligative properties don´t depend on the chemical composition of a
solute, but depend only on the number of solute particles (molecules or
ions).
The processes based on colligative properties are:
• Diffusion
• Dialysis
• Osmosis
• Freezing point depression
• Boiling point elevation
Diffusion
is a process of spontaneous movement of particles of a dissolved
compound from a region of higher concentration to a region of lower
concentration, to distribute themselves uniformly = movement of a
substance down a concentration gradient
The rate of diffusion depends on the concentration gradient
Particles move until equilibrium is reached
Diffusion usually happens in a solution in gas or in a liquid.
Examples of diffusion:
A sugar cube is left in a beaker of water for a while.
The smell of food spread in the whole house
Biomedical importance of diffusion
Exchange of O2 and CO2 in lungs and in tissues
Certain nutrients are absorbed by diffusion in the gastrointestinal
tract e.g. water soluble vitamins, minerals...
Dialysis
Water and low molecular weight (LMW) compounds (not
macromolecules) are transported across a semipermeable
membrane. LMW compounds go from the more concentrated
solution to the less concentrated solution till equilibrium is reached.
Dialysis
Concentrated
sugar solutionDiluted
sugar solution
Movement of LMW solute
to equal concentrations
Semipermeable
membrane
Water and low molecular weight (LMW) compounds (not
macromolecules) are transported across a semipermeable
membrane. LMW compounds go from the more concentrated
solution to the less concentrated solution till equilibrium is reached.
Biomedical importance of dialysis
Biological ultrafiltrates
Many extracellular fluids like interstitial fluids, cerebrospinal fluid,
glomerular filtrate of kidneys are formed by ultrafiltration. Proteins do not
appear in ultrafiltrates.
Hemodialysis - Blood dialysis
- in patients with acute kidney injury blood is dialyzed in artificial
kidney to eliminate waste products (e.g. urea or creatinine) or toxins
Filtered blood
returning to body Blood flows to
dialyzer
Hemodialyzer
machine
Hemodialyzer
(where filtering takes place)
Biomedical importance of dialysis
Dialyzing
membrane
Dialysate
- solution isotonic with blood,
- it has the same concentrations of all the
essential substances that should be left in blood
Dialysate
Osmosis
Osmosis
Osmosis is the flow of solvent (e.g. water) across a semipermeable
membrane (with smaller pores than dialyzing membrane) from a
lower solute concentration to a higher solute concentration
semipermeable membrane is permeable only to solvent molecules,
not to solute molecules
Concentrated
solution
Diluted
solution
Semi-permeable
membrane
Osmotic pressure (π)
- external pressure that has to be applied on the more
concentrated solution to prevent osmosis
i – number of solute particles in solution to which the compound dissociates
c – amount of substance concentration (mol/L)
R – gas constant – 8.314 J K-1 mol-1
T – temperature in Kelvins (0 °C = -273.15 K)
π = i . c . R . T
π of blood - 780 kPa
Movement of solvent (water)
to equal concentrations
π
Osmolarity (cosm)
molar concentration of all osmotically active particles of solutes in
solution
cosm = i . c
cosm - osmolarity mol/L
i – number of solute particles in solution to which the
compound dissociates
c – amount of substance concentration (mol/L)
Example 2:
Calculate osmolarity of the solution containing 0.2 mol/L CaCl2
and 0.1 mol/L glucose.
1. CaCl2 = Ca2+ + 2Cl-
2. Glucose no dissociation
cosm = i1 . c1 + i2 . c2
cosm = 3 x 0.2 + 1 x 0.1 = 0.7 mol/L
i1 = 1Ca2+ + 2Cl
- = 3
i2 = 1glucose
Osmolarity (cosm)
Blood serum osmolarity:
πblood = i . c . R . T
cosm
πblood
Blood cosm = = 0.3 mol/L R . T
.
= 780 kPa
Isotonic /isoosmotic solutions
Isotonic solutions are two solutions of equal osmolarity.
Hypertonic solution
Hypertonic solution is one of two solutions that has a higher
osmolarity.
Hypotonic solution
Hypotonic solution is one of two solutions that has a lower
osmolarity.
hemolysis
Crenation
Cells shrink
Solution of NaCl with concentration of 0.15 mol/L
Solution of NaCl with osmolarity of 0.3 mol/L
0.9% NaCl solution (9 g NaCl/L)
Physiological solution
Solution which osmotic pressure corresponds to blood plasma:
Any solution added in large quantity into the bloodstream has
to be isotonic with blood!!
Oncotic pressure
The capillary wall is permeable for small molecules and water but not
permeable for proteins
protein
Oncotic pressure
Oncotic pressure, or colloid osmotic pressure, is a form
of osmotic pressure exerted by proteins (e.g. albumin) in a blood
that usually tends to pull water into the circulatory system.
Water flow driven by
oncotic pressure
diference
Capilary
lumen
Interstitial
space
Interstitial
space
• Within the extracellular fluid, the distribution of H2O between blood
and ISF (interstitial fluid) depends on the plasma protein concentration.
• The capillary wall, which separates plasma from the ISF is freely
permeable to H2O and electrolytes, but restricts the flow of proteins.
• Albumin makes about 80% of oncotic pressure.
Osmotic pressure of blood plasma: 780 – 795 kPa
Oncotic pressure of blood plasma: 2.7 -3.3 kPa
• 2.7 -1.4 kPa ......sizable edemas
• <1.4 kPa...........unless albumin is given i.v., survival is hardly possible
6
The significance of oncotic pressure
Force of pumping blood
from heart pushes fluids
from blood into
interstitial fluid
Proteins that remain in
blood attract interstitial
fluid back into
bloodstream
Filtration No net movement Reabsorption
Fluid exits capillary since
capillary hydrostatic
pressure (35 mm Hg) is
greater than blood oncotic
pressure (25 mm Hg)
No net movement of fluid
since capillary hydrostatic
pressure (25 mm Hg) = blood
oncotic pressure (25 mm Hg)
Fluid re-enters capillary since
capillary hydrostatic pressure
(15 mm Hg) is less than blood
oncotic pressure (25 mm Hg)
Arterial end
net filtration pressure
= +10 mm Hg
Mid capillary
net filtration pressure
= 0 mm Hg
Venous end
net filtration pressure
= -7 mm Hg
Small molecules and ions can be dialyzed in both directions between
blood and the interstitial compartment
Large protein molecules do not have this ability – their presence
produces excess osmotic pressure of blood (oncotic pressure)
compared to the interstitial fluid.
The hydrostatic pressure of the blood (at the arterial end of
capillary) tends to push water out of the capillary – filtration.
The oncotic pressure (at the venous end of capillary) pulls the
water from the interstitial space back into the capillary –
reabsorption.
Important function of oncotic pressure:to maintain water in capillaries
If capillaries become more permeable for proteins
(surgical procedures or extensive burns)
proteins migrate from blood
loss of blood oncotic pressure
total blood volume decreases
reduces the ability of blood to transfer oxygen and to eliminate CO2
Decrease of blood volume associated with insufficient brain oxygen supply
leads to shock
Biomedical importance of oncotic pressure
If plasma oncotic pressure is reduced (starvation, kidney disease)
Reduced force drawing water back into capillary from interstitium
Biomedical importance of oncotic pressure
Edema
Accumulation of excess fluid in tissue spaces
• What is edema and what is the general cause?
Accumulation of water in extracellular space.
General cause: filtration of blood is much higher than reabsorption of
water back to the bloodstream
Examples: decreased plasma protein concentration, increased capillary
permeability to proteins, conditions which elevate venous blood pressure.
• What are some examples that can cause edema?
Colloidal dispersions
Colloidal dispersion
size of particles 1 – 1000 nm
almost all reactions in the organism proceed in colloid environment
True
solution
Colloid
High–molecular weight
(macromolecular) compounds
(e.g. proteins, polysaccharides)
Colloidal dispersion
Low–molecular weight compounds
by clustering of molecules into
aggregates – micelles
(e.g. soap solutions).
Classification of colloids
Colloids are classified based on the following criteria:
Physical state of dispersed phase and dispersion medium
Affinity of dispersed molecules with dispersing media
Classification of colloids
1. Based on physical state of dispersed phase and
dispersion medium
• Sol – colloids with solid particles dispersed in a liquid
• Emulsion - liquid dispersed in liquid (immiscible liquids)
• Foam – gas particles dispersed in a liquid or solid
• Aerosol - small liquid particles or solid dispersed in a gas
Sauces and dressingsclouds
• Gel - liquid particles dispersed in a solid
gelatin
Sols
Are colloidal solutions made of globular proteins with normal
viscosity
Sol - a colloidal solution appears as fluid
Gels
.
they arise by swelling macromolecular compounds (e.g.proteins) in
solvent – acceptation of water by solid polymers
are formed from fibrous proteins (gelatin from collagen), polysaccharides
(gels – dextran, sephadex).
Gels - a colloidal solution appears as solid
Gels undergo aging - particles coagulate, gel volume diminishes and
water is displaced
Emulsions
are colloidal dispersions of two immiscible liquids (e.g. oil in water, or
water in oil) when are shaken together.
usually are not stable (e.g. the oil soon separates from the aqueous layer).
can be stabilized by a third component called emulsifying agent
(emulsifiers).
Biologically important emulsifying agents are salts of bile acids.
Emulsifiers
Hydrophilic
water-loving head Hydrophobic
water-fearing tail
• All emulsifiers have 2 components: hydrophilic head
hydrophobic tail
• enable fat to be uniformly dispersed in water as an emulsion (they
stabilize emulsions).
• Their action is similar to soap in washing
emulsifier
Emulsifiers
Oil droplets
Hydrophilic head will associate with water and its hydrophobic tail with oil droplets.
This prevents separation of two layers and thus stabilizes the emulsion.
emulsifier emulsifier
Hydrophobic groups
(nonpolar)
Hydrophilic groups
(polar)
Bile acids as emulsifiers
Fat
Fat
Fat
Fat
Fat Fat
Step 1: Emulsification of fat droplets
Step 2: Hydrolysis of triacylglycerols in emulsified fat droplets into fatty acid and monoacylglycerols
Sterp 3: Dissolving of fatty acids and monoacylglycerols into micelles to produce „mixed micelles“
Bile acids as emulsifiers
Foam
is composed of small bubbles of gas (usally air) dispersed in a liquid (e.g.
egg white foam)
As liquid egg white is whisked, air bubbles are incorporated.
If egg white is heated, protein coagulates and moisture is driven off. This
forms a solid foam, e.g. a meringue
Colloids
LyophilicLiquid-loving Lyophobic
Liquid-hating
Micelles
2. Classification of colloids according to the affinity
of dispersed molecules with dispersing media
Macromolecular
compounds
Low-molecular
weight compounds
Low-molecular
weight amphipatic
compounds (soaps)
1. Lyophilic colloids
• If water is the solvent (dispersing medium), it is known as a
hydrosol or hydrophilic colloids
• particles of a lyophilic colloid are stabilized in solution
(prevention of aggregation) by solvation (hydration) shell, i.e.
oriented solvent molecules
• are formed by spontaneous dissolving of macromolecular substances
(e.g. solutions of proteins, starch...)
1. Lyophilic colloids
The loss of hydration shell after excess of neutral salt (electrolyte) is
added into solution results in irreversible salting out (precipitation)
of particles from solution.
The living cells represent solutions of lyophilic colloids (as well as
coarse dispersions)
• solvent hating colloids, have no affinity for the dispersion medium
2. Lyophobic colloids
• unstable colloid systems in which the dispersed particles:
- tend to repel liquids,
- are easily precipitated
• require protective colloids (lyophilic colloids – gums, gelatin...) to
stabilize in water
Lyophobic soll particle
(particle being protected)
Lyophilic colloidal particle
(protecting particle)
Explanation: The particles of the hydrophobic sol adsorb the particles
of the lyophilic particles. The hydrophobic colloid, therefore, behaves
as a hydrophilic sol and is precipitated less easily by electrolytes.
2. Lyophobic colloids
• are made artificially by aggregation of low molecular weight substances
• Examples: sols of metals and their insoluble compounds like sulphides and
oxides (e.g. gold, silver, platinum in water, cluster of inorganic molecules,
e.g. As2S3)
• Aplication in therapy: colloidal systems are used as therapeutic agents
Silver colloid – germicidal effects
Copper colloid – anticancer effects
Mercury colloid - antisyphilis
Colloidal gold Colloidal silver
3. Association colloids – micelles
are formed by dissolving of low-molecular weight amphipathic compounds
Amphipathic compounds contain both polar (hydrophilic) and nonpolar
hydrophobic regions (e.g. fatty acids, phospholipids)
Polar part
Nonpolar part
when mixed with water, amphipatic compounds form colloidal particles –
micelles (e.g. soap, detergents)
Hydrophilic headHydrophobic tail
Biological importance of colloids
Biological compounds as colloidal particles: high-molecular weight proteins,
complex lipids and polysaccharides
Blood coagulation: when blood clotting occurs, the sol is converted finally
into the gel.
Biological fluids as colloids: these include blood, milk and cerebrospinal
fluid, lymph, mucus, cytosol, nucleus, cell membranes
Colloidal state is one of the most widespread in nature:
Reaction kinetics
Chemical reaction
Reaction means a change
Chemical reaction is a conversion of reactants to products
A + B C + DReactants Products
Reagents
Irreversible reactions
Reversible reactions
A + B C + D ReactantsProducts
Chemical kinetics
Kinetics of a chemical reaction can tell us:
how fast the concentration of A or B decreases
how fast the concentration of product C increases
A + B C
Rate equation(Guldberg Waage rate law)
The rate of a given chemical reaction (at constant temperature and
pressure) is proportional to product of reactants concentration.
Rate: v = k . [A]a . [B] b
k = rate constant
[A], [B]= molar concentrations of reactants (mol/L)
For the general reaction:
aA + bB cC
Rate constant
k = rate constant
A = Arrhenius constant for each chemical reaction (total number of collisions)-frequency factor
Ea = activation energy
R = gas constant (8.314 J K-1 mol-1)
T = Temperature in Kelvins
e = euler number (2.71828...)
Temperature has a dramatic effect on reaction rate.
For many reactions, an increase of 10°C will double the rate.
Arrhenius equation
Effective collisions
For reactants to make products
They must collide in the correct orientation and with sufficient energy
The correct orientation of collisions
A B C D
A B A B
Effective collisions
For reactants to make products
They must collide in the correct orientation and with sufficient energy
The energy of collision must be greater than the bond energy between
the atoms
Activation energy
The minimum amount of energy required to start a chemical reaction
Activation energy
Transition state
(activated complex)
Activation energy
Reactants
Products
Factors which affect the rate of chemical
reactions
Rate of
reaction
The nature of
reactants
Temperature
Concentration
of reactants
Catalysts
Natu
reo
fre
acta
nts Number of bonds
• fewer bonds per reactant - faster reaction
Strength of bonds• Breaking of weaker bonds - a faster rate
(-C-C- / -C=C-)
The size and shape of a molecule• Complicated molecules or complex ions
are often less reactive
Less particles, less frequent
and successful collision
More particles, more frequent
and successful collision
Concentration of reactants
As the concentration of reactants increases, so does the likelihood that reactant
molecules will collide - the reaction rate will increase
A temperature increase of about 10°C will often double the rate of a reaction
Higher
temperatureHigher
speedMore high-energy
collisionsMore collisions
that break bonds
Faster
reaction
Temperature
Kinetic energy
Catalysts Catalysts speed up reactions by changing the mechanism of the reaction –
they reduce activation energy of reaction
Catalysts are not consumed during the course of the reaction
EaEa
Does a catalyst speed up the reaction in only one direction or both?
Does a catalyst shift the location of the equilibrium position?
No! Catalysts do not affect the amounts of reactants and products present at
equilibrium, just the time it takes to establish equilibrium.
A catalyst speeds up the forward and reverse reactions exactly the same
Oxidation – reduction reactions
(redox reactions)
Oxidation is the loss of electrons (or hydrogen), the species which loses
the electrons is oxidized, it becomes more positive
Reduction is the gain of electrons (hydrogen), the species which gains
electrons is reduced, becomes less positive.
Na0 → Na+ + 1e-
Cl20 + 2e- → 2Cl-
Oxidation and reduction reactions occur simultaneously
chemical reactions where one of the reactants is oxidized and one of the
reactants is reduced
Biological oxidation-reduction reactions
In biological systems, oxidation is often synonymous with dehydrogenation
Many enzymes that catalyze oxidation reactions are oxidoreductases, called
dehydrogenases.
O : H ratio1 : 6
O : H ratio1 : 4
O : H ratio1 : 2
More reduced compounds are richer in hydrogen
The oxidation states of carbon in biomolecules
Most oxidized
Most reduced
Oxidizing agent – oxidant - is the chemical species causing the
oxidation. This species is reduced and can also be called the
electron acceptor.
2Na0 + Cl20 2Na+Cl-
oxidant
Reducing agent – reductant- is the species causing the
reduction. This species is oxidized and can be called the electron
donor.
reductant
The number of electrons lost by the reductant must be equal to the
number of electrons gained by the oxidant.
e-
Compounds can be oxidized by one of four different ways:
1. Direct transfer of electrons Fe2+ + Cu2+ Fe3+ + Cu+
2. Transfer of hydrogen atoms
H = H+ + 1e- AH2+ B ↔ A + BH2
Hydrogen/electron donor
Reduced
3. Transfer of hydride ion (Hˉ), which has two electrons (H+ + 2e-)
This occurs in the case of NAD+-linked dehydrogenases
4. Through the direct combination with oxygen
R−CH3+ ½O2 R−CH2OH
Dismutation (disproportionation)
The special case of oxidation – reduction reaction
a compound of intermediate oxidation state converts to two different
compounds, one of higher and one of lower oxidation states.
Examples:
The dismutation of superoxide free radical to hydrogen peroxide and oxygen,
catalysed in living systems by the enzyme superoxide dismutase
2 O2. −1 + 2 H+1 → H2
+1 O2-2 + O2
0SOD
Potassium chlorate decomposes at elevated temperature into perchlorate
and potassium chloride4 KClO3 = 3 KClO4 + KCl
4 ClIV = 3 ClVII + Cl-1
Oxidation-reduction reactions
Oxidation – reduction reactions occur together
Fe2+ + Cu2+ Fe3+ + Cu+
This reaction can be described in two half-reactions:
(1) Fe2+ Fe3+ + 1e-
(2) Cu2+ + 1e- Cu+
Reductant – donates electrons
Oxidant – accepts electrons
Electron donor e- + electron acceptor
Conjugate redox pair
Reduction potentials
When two conjugate redox pairs are together in solution, electron transfer from the electron donor of one pair to the electron acceptor of the other may occur spontaneously.
The tendency for a reaction depends on the relative affinity of the electron
acceptor of each redox pair for electrons.
The standard reduction potential (E0) is the tendency for a chemical
species to be reduced, and is measured in volts at standard conditions
2H+ + 2eˉ H2 E0 = 0 V
The electrode at which this half-reaction occurs is arbitrarily assigned a
standard reduction potential of 0.00V.
Fe2+ + Cu2+ Fe3+ + Cu+
Element with the more positive redox potential has a higher
affinity towards electrons – it has an oxidizing property
Element with the more negative redox potential has a lower
affinity towards electrons – it can easily donate electrons – it has
an reducing property
You can tell how likely a compound is to be oxidized from the
reduction potential
-0.77 Fe2+ (aq) - 1e- Fe3+ (aq)
+0.161 Cu2+ (aq) + 1e- Cu1+ (aq)
Standard reduction potentials at 25oC
R is gas constant (8.314 JKˉ1molˉ1
T is temperature (in Kelvins)
n is the number of electrons transferred per molecule
F is the Faraday constant (9.68 . 104 Cmolˉ1).
The Nerst – Peterson equation:
The reduction potential of a half-cell depends on:
the chemical species present
on their concentrations
Application of reduction potentials
Known oxidation-reduction potentials of biological redox systems allow to
determine the direction and sequence of oxidation-reduction reactions in
biological systems.
The strict sequence of enzymatic reactions in “respiratory chain” allows a
gradual release of energy during biological oxidation.
Electron transport chain
Mixing two or more solutions having the same
solute, but different concentrations:
Solute A
concentration of the solution 1– c1
Volume of the solution 1 – V1
Solute A
c2 - concentration of the solution 2
V2 – volume of the solution 2
Solute A
c3 - concentration of the solution 3
V3 – volume of the final solution 3
c1V1 + c2V2 = c3V3
n1 + n2 = n3
12
3
Example:
400 mL of a 0.1 mol/L NaCl solution is mixed with 100 mL of a
0.2 mol/L NaCl solution. What is the concentration of the final
solution?
c1V1 + c2V2 = c3V3
NaCl NaCl+Final
NaCl
c1 = 0.1 mol/L
V1 = 400ml=0.4 L
c3 = ?
V3 = 0.4 + 0.1 = 0.5 L
c2 = 0.2 mol/L
V2 = 100ml= 0.1 L
c1V1 + c2V2 0.1*0.4 + 0.2*0.1
c3 = = = 0.12 mol/LV3 0.5
Ways to dermine concentration of
solutions
To determine how much solute is dissolved in a unit amount of
solution
Ways to determine concentration of solution:
Molar concentration – molarity – amount of substance concentration
Molal concentration – molality
Mass concentration (density)
Weight fraction / Volume fraction
Weight / volume percent
Amount of substance concentration, molar
concentration, molarity ( c )
parameter unit
c – molar concentration mol/L
n – moles of the solute mol
V – volume of the solution liter (L)
m – mass of the solute gram (g)
Mw – molecular weight gram/mol (g/mol)
n m/Mw m
c = = =
V V Mw . V
Example
Calculate amount of substance concentration of a solution which has
18 g of glucose in 2 liters of water (Mwglucose = 180 g/mol)?
m = 18g Mw glc = 180 g/mol V = 2 L
c = ?
18 g
c = = 0.05 mol/L
180 g/mol * 2L
n m
c = =
V Mw . V
Molal concentration, Molality
molality (mol / kg)
n – moles of the solute (mol)
m – mass of the solvent (kg)
n
msolvent
Example
mMgCl2 =45.7g MwMgCl2 = 95.21g/mol msolvent = 2.4 kg
n m/ Mw m 45.7g
Molality = = = =
msolvent msolvent Mw . msolvent 95.21 g/mol . 2.4 kg
Molality = 0.2 mol/kg
45.7 g of magnesium chloride (MwMgCl2 = 95.21g/mol) is dissolved in
2.40 kg of water. What is the molality of the solution?
Weight (mass) concentration
– weight concentration g/L
m – mass of the solute g
V – volume of the solution L
Example
Glucose concentration in blood is 5 mmol/L. What is its mass concentration?
(Mwglucose = 180 g/mol)
c = 5 mmol/L = 0.005 mol/L Mw glucose = 180 g/mol
= ?
m
c =
Mw . V
m
= = c . Mw= 0.005mol/L . 180 g/mol = 0.9 g/L
V
Weight fraction (w)
mB
wB =
msolution
wB - weight fraction (g/g)
mB – mass of the solute B (g)
Weight percent (w%)
mB
w% = . 100%
msolution
w% - weight percent
mB – mass of the solute B (g)
Example
w = 0.5 % mNaCl = 2 g
msolution = ?
Calculate the mass of NaCl solution, when its mass percentage is
w = 0.5 % and it contains 2 g of NaCl.
2
0.5 = . 100%
msolution
mB
wB = . 100%
msolution
200
msolution= = 400 g
0.5
Volume fraction (vB)
VB
vB =
Vsolution
v B - volume fraction (mL/mL, or L/L)
VB – volume of the solute B (mL or L)
VB
v% = . 100%
Vsolution
V% - volume percent
VB – volume of the solute B (mL or L)
Volume percent (v%)
Thank you for your attention...