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DISPERSED SYSTEMS Ingrid Žitňanová

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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+ + SO4

    2-

    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

  • 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– c1Volume 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...