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D. x. n. , professor, head of the department of physical chemistry of the Russian Chemical Technology University named after D. I. Mendeleeva Konyukhov Valery Yurievich [email protected] ru vkontakte. ru
Literature Vishnyakov A. V., Kizim N. F. Physical chemistry. M.: Chemistry, 2012 Physical chemistry // Ed. K. S. Krasnova. M.: Higher school, 2001 Stromberg A. G., Semchenko D. P. Physical chemistry. M.: Higher School, 1999. Fundamentals of Physical Chemistry. Theory and tasks: Textbook. Manual for universities / V. V. Eremin et al. M.: 2005.
Literature Atkins P. Physical chemistry. M.: Peace. 1980. Karapetyants M. Kh. Chemical thermodynamics. M.: Chemistry, 1975.
LOMONOSOV Mikhail Vasilyevich (1711 -65), the first Russian natural scientist of world importance, a poet who laid the foundations of the modern Russian literary language, artist, historian, champion of the development of national education, science and economics. Born on November 8 (19) in the village of Denisovka (now the village of Lomonosovo) in a Pomor family. At the age of 19 he left to study (from 1731 at the Slavic-Greek-Latin Academy in Moscow, from 1735 at the Academic University in St. Petersburg, in 1736-41 in Germany). From 1742 an adjunct, from 1745 an academician of the Petersburg Academy of Sciences.
In 1748 he founded the first chemical laboratory in Russia at the Academy of Sciences. Moscow University was founded on the initiative of Lomonosov (1755). Developed atomic-molecular concepts of the structure of matter. During the period of dominance of the caloric theory, he argued that heat is due to the movement of corpuscles. Formulated the principle of conservation of matter and motion. Eliminated phlogiston from the list of chemical agents. He laid the foundations of physical chemistry.
Investigated atmospheric electricity and gravity. He advanced the doctrine of color. Created a number of optical instruments. Discovered the atmosphere on Venus. He described the structure of the Earth, explained the origin of many minerals and minerals. Published a guide to metallurgy. He emphasized the importance of researching the Northern Sea Route and the development of Siberia. He revived the art of mosaic and smalt production, created mosaic paintings with his students. Member of the Academy of Arts (1763). Buried in St. Petersburg in the necropolis of the 18th century.
Lomonosov's definition: “Physical chemistry is a science that studies, on the basis of the provisions and experiments of physics, what happens in complex bodies during chemical operations…. Physical chemistry can be called chemical philosophy. "
In Western Europe, it is considered to be the year of the creation of physical chemistry in 1888, when W. Ostwald began to read this course, accompanied by practical exercises, and began to publish the journal "Zeitschtift fur physikalische Chemie". In the same year, the Department of Physical Chemistry was organized at the University of Leipzig under the leadership of W. Ostwald.
Born and lived for a long time in the Russian Empire, at 35 he changed his Russian citizenship to German. In Leipzig, he spent most of his life, where he was called “the Russian professor”. At the age of 25 he defended his doctoral dissertation on the topic "Bulk-chemical and optical-chemical research."
In 1887 he accepted an offer to move to Leipzig, where he founded the Physicochemical Institute at the university, which he headed until 1905. In 1888, he occupied a very prestigious department of physical and inorganic chemistry at the University of Leipzig. In this position, he worked for 12 years.
From the "Leipzig School" W. Ostwald left: Nobel laureates S. Arrhenius, J. Van't Hoff, W. Nernst, famous physicochemists G. Tamman and F. Donnan, organic chemist J. Wislicens, famous American chemist G. N. Lewis. Over the years, Ostwald trained Russian chemists: I. A. Kablukov, V. A. Kistyakovsky, L. V. Pisarzhevsky, A. V. Rakovsky, N. A. Shilov and others.
One of the unique features of Ostwald was his many years of active rejection of the atomic-molecular theory (although he proposed the term "mole"). “The chemist does not see any atoms. - He studies only simple and understandable laws, which obey the mass and volume ratios of reagents. "
W. Ostwald contrived to write a voluminous chemistry textbook in which the word "atom" is never mentioned. Speaking on April 19, 1904 in London with a big report to the members of the Chemical Society, Ostwald tried to prove that atoms do not exist, and "what we call matter is just a collection of energies gathered together in a given place."
In honor of V. Ostwald, a memorial plaque with an inscription in Estonian, German and English was installed on the territory of the University of Tartu
predict whether the reaction can proceed spontaneously; if the reaction proceeds, then how deeply (what are the equilibrium concentrations of the reaction products); if there is a reaction, then at what speed.
1. STRUCTURE OF MATTER In this section, on the basis of quantum mechanics (Schrödinger equation), the structure of atoms and molecules (electronic orbitals of atoms and molecules), crystal lattices of solids, etc. are explained, and the states of aggregation of matter are considered.
2. CHEMICAL THERMODYNAMICS on the basis of the laws (principles) of thermodynamics allows: to calculate the thermal effects of chemical reactions and physicochemical processes, to predict the direction of chemical reactions, to calculate the equilibrium concentrations of reagents and reaction products.
3. THERMODYNAMICS OF PHASE EQUILIBRIES Studies the regularities of phase transitions in one-component and multicomponent (solutions) systems. Its main goal is to plot the phase equilibrium diagrams of the indicated systems.
4. ELECTROCHEMISTRY Studies the properties of electrolyte solutions, the peculiarities of their behavior in comparison with molecular solutions, investigates the patterns of interconversion of the energy of chemical reactions and electrical energy during the operation of electrochemical (galvanic) cells and electrolysers.
5. CHEMICAL KINETICS AND CATALYSIS Studying the regularities of the course of chemical reactions in time, investigates the influence on the rate and mechanism of reactions of thermodynamic parameters (pressure, temperature, etc.), the presence of catalysts and inhibitors.
In a separate science, COLLOID CHEMISTRY, a section of physical chemistry is distinguished - the physical chemistry of surface phenomena and dispersed systems.
Classical thermodynamics is a branch of theoretical physics and studies the patterns of interconversions of various types of energy and energy transitions between systems in the form of heat and work (termo - heat, dynamo - motion).
Thermodynamics abstracts from the causes of any process, and the time during which this process occurs, but only operates with the initial and final parameters of the system participating in any physicochemical process. The properties of individual molecules are not taken into account, but the averaged characteristics of systems consisting of many molecules are used.
The tasks of chemical thermodynamics are: measuring and calculating the thermal effects of chemical reactions and physicochemical processes, predicting the direction and depth of reactions, analyzing chemical and phase equilibria, etc.
1. 1. Basic concepts and definitions of TD In thermodynamics, all processes of interest to us occur in thermodynamic systems. A system is a body or a group of bodies, actually or mentally identified by the observer in the environment.
The system is a part of the surrounding world that interests us especially. Everything else in the universe is the environment (environment). It is generally accepted that the environment is so large (has an infinite volume) that the exchange of energy with a thermodynamic system does not change its temperature.
By the nature of the exchange of energy and matter with the environment, systems are classified: isolated - they cannot exchange either matter or energy; closed - they can exchange energy, but they cannot - with matter; open - they can exchange both matter and energy.
According to the number of phases, the systems are subdivided into: homogeneous - consist of one phase (solution of Na. Cl in water); heterogeneous - the system includes several phases, separated from each other by interfaces. An example of heterogeneous systems is ice floating in water, milk (droplets of fat are one phase, aqueous medium is another).
A phase is a set of homogeneous parts of a system that have the same chemical and physical properties, and are separated from other parts of the system by interfaces. Each phase is a homogeneous part of a heterogeneous system
By the number of components, the systems are subdivided into one- two-, three-component and multicomponent. Components are the individual chemicals that make up a system that can be isolated from the system and exist outside of it.
Any thermodynamic system can be characterized by a set of a huge number of physical and chemical properties that take certain values: temperature, pressure, thermal conductivity, heat capacity, component concentrations, dielectric constant, etc.
Chemical thermodynamics deals with those properties that can be unambiguously expressed as a function of temperature, pressure, volume, or concentrations of substances in a system. These properties are called thermodynamic properties.
The state of a thermodynamic system is considered specified if its chemical composition, phase composition and values of independent thermodynamic parameters are indicated. Independent parameters include: pressure (P), volume (V), temperature (T), the amount of substance n in the form of a number of moles or in the form of concentrations (C). These are called state parameters.
According to the current system of units (SI), the main thermodynamic parameters are set in the following units: [m 3] (volume); [Pa] (pressure); [mol] (n); [K] (temperature). As an exception, in chemical thermodynamics, it is allowed to use the off-system unit of pressure, the normal physical atmosphere (atm), equal to 101.325 kPa.
Thermodynamic parameters and properties can be: Intensive - they do not depend on the mass (volume) of the system. These are temperature, pressure, chemical potential, etc. Extensive - they depend on the mass (volume) of the system. These are energy, entropy, enthalpy, etc. During the formation of a complex system, intensive properties are leveled, and extensive ones are summed up.
Any change that occurs in the system and is accompanied by a change in at least one thermodynamic parameter of state (properties of the system) is called a thermodynamic process. If the course of the process changes the chemical composition of the system, then such a process is called a chemical reaction.
Usually, during the course of the process, one (or several) parameters are kept constant. Accordingly, they are distinguished: isothermal process at constant temperature (T = const); isobaric process - at constant pressure (P = const); isochoric process - at constant volume (V = const); adiabatic process in the absence of heat exchange with the environment (Q = 0).
When processes occur in non-insulated systems, heat absorption or release can occur. In accordance with this feature, processes are divided into exothermic (heat is released) endothermic (heat is absorbed).
In the course of the process, the system passes from one equilibrium state to another equilibrium state. Thermodynamic equilibrium is the state of the system in which there is thermal, mechanical and chemical (electrochemical) equilibrium with the environment and between the phases of the system.
Equilibrium states are: stable; metastable. A process is called equilibrium (quasi-static) if it passes infinitely slowly through a continuous sequence of equilibrium states of the system.
Processes that occur by themselves and do not require external energy for their implementation are called spontaneous (positive) processes. when energy is extracted from the environment for the implementation of the process, that is, work is done on the system, then the process is called non-spontaneous (negative).
State functions State functions are properties of the system (internal energy U, enthalpy H, entropy S, etc.), they characterize the given state of the system. Their changes in the course of the process do not depend on its path and are determined only by the initial and final states of the system.
The infinitesimal change in this function is the total differential d. U, d. S, etc.:
Process (transition) functions Process functions (heat Q, work W) - they are not properties of the system (they are not in the system), they arise in the course of the process in which the system participates.
If there is no heat and work in the system, then it makes no sense to talk about their change, we can only talk about their quantity Q or W in a particular process. Their quantities depend on the way the process is carried out. Infinitesimal quantities represent Q, W.
Movement is an attribute of matter. Energy is a measure of movement, that is, a quantitative and qualitative characteristic. Energy is a function of the state of the system. Its change in a particular process does not depend on the path of the process and is determined only by the initial and final states of the system.
Many different types of energy are known: mechanical, electrical, chemical, etc., but from system to system energy can only be transferred in two forms: in the form of heat or work.
Heat (Q) is a form of energy transfer from system to system due to the chaotic movement of particles (molecules, atoms, ions, etc.) of contacting systems.
In thermodynamics, the heat supplied to the system is assumed to be positive (for example, the heat of an endothermic reaction), and the heat removed from the system is assumed to be negative (the heat of an exothermic reaction). In thermochemistry, the opposite is true.
Work is a form of energy transfer from system to system due to the directed movement of micro- or macro-objects. In the literature, work is denoted either W (from the English "work") or A (from German "arbait").
There are different types of work: mechanical, electrical, magnetic, surface changes, etc. An infinitesimal work of any kind can be represented as the product of a generalized force and a change in a generalized coordinate, for example:
The sum of all types of work with the exception of work against external pressure forces P - expansion - compression work is called useful work W ':
In thermodynamics, work is considered positive if it is performed by the system itself and negative if it is performed on the system. According to IUPAC recommendations, it is considered positive to work done on the system ("egoistic" principle - positive is that which increases internal energy)
Work of expansion of an ideal gas in various processes 1. Expansion into vacuum: W = 0. 2. Isochoric reversible expansion: d. V = 0 W = 0
The conclusions and relationships of thermodynamics are formulated on the basis of two postulates and three laws. Any isolated system over time comes to an equilibrium state and cannot spontaneously leave it (the first postulate) That is, thermodynamics does not describe systems of an astronomical scale and microsystems with a small number of particles (
A spontaneous transition from a nonequilibrium state to an equilibrium state is called relaxation. That is, an equilibrium state will necessarily be achieved, but the duration of such a process is not determined, since there is no concept of time.
The second postulate If system A is in thermal equilibrium with system B, and that - with system C, then systems A and C are also in thermal equilibrium
The internal energy of any thermodynamic system U consists of the kinetic (energy of motion) and potential (interaction energy) energies of all particles (molecules, nuclei, electrons, quarks, etc.) that make up the system, including unknown types of energy.
The internal energy of a system depends on its mass (extensive property), on the nature of the system's substance and thermodynamic parameters: U = f (V, T) or U = (P, T) is measured in J / mol or J / kg. U is a state function, therefore U does not depend on the process path, but is determined by the initial and final state of the system. d. U - full differential.
The internal energy of the system can change as a result of the exchange of energy with the environment only in the form of heat or work.
This fact, which is a generalization of the practical experience of mankind, conveys the first law (beginning) of thermodynamics: U = Q - W In differential form (for an infinitely small part of the process): d. U = Q W
"The heat supplied to the system is used to increase the internal energy of the system and to perform work by the system."
For an isolated system, Q = 0 and W = 0, that is, U = 0 and U = const. The internal energy of an isolated system is constant
In the formulation of Clausius: "The energy of the world is constant." A perpetual motion machine of the first kind (perpetum mobile) is impossible. Different forms of energy pass into each other in strictly equivalent amounts. Energy does not arise and is not destroyed, but only passes from system to system.
The U function is additive. This means that if two systems characterized by the values of U 1 and U 2 are combined into one single system, then the resulting internal energy U 1 + 2 will be equal to the sum of the energies of its constituent parts: U 1 + 2 = U 1 + U 2
In the general case, heat Q is a function of the process, that is, its amount depends on the path of the process, but in two cases important for practice, heat acquires the properties of a state function, that is, the value of Q ceases to depend on the path of the process, but is determined only initial and final states of the system.
We will assume that during the process only work can be performed against the forces of external pressure, and useful work W = 0: Q = d. U + P d. V, and since V = const, then P d. V = 0: QV = d. U or in integral form: QV = Uк - Uн
Again we will assume that the useful work W = 0, then: Q = d. U + P d. V, Since Р = const, it is possible to write down: QР = d. U + d (PV), QP = d (U + P V). Let's designate: Н U + P V (enthalpy) QР = d. H or: QP = Hк - Hн
Thus, the thermal effect of a chemical reaction acquires properties as a function of state at P = const: QP = H; at V = const: QV = U.
Since chemical reactions and physicochemical processes are often carried out at constant pressure (in the open air, that is, at P = const = 1 atm), in practice the concept of enthalpy is often used for calculations, rather than internal energy. Sometimes the word "heat" of the process is replaced without additional explanation by "enthalpy", and vice versa. For example, they say "the warmth of education", but write f. N.
But if the process of interest to us occurs at V = const (in an autoclave), then the expression should be used: QV = U.
Let's differentiate the expression: H = U + P V d. H = d. U + Pd. V + Vd. P, at constant pressure V d. P = 0 and d. H = d. U + P d. V In integral form: H = U + P V
For an ideal gas, the Clapeyron-Mendeleev equation is valid: Р V = n R T, where n is the number of moles of the gas, R 8, 314 J / mol K is the universal gas constant. Then (at T = const) P V = n R T. Finally we have: H = U + n R T n - the change in the number of moles of gaseous substances during the reaction.
For example, for the reaction: N 2 (g) + 3 H 2 (g) = 2 NH 3 (g) n = -2, and for the reaction: 2 H 2 O (g) 2 H 2 (g) + O 2 ( d) n = 3.
The differences between QV and QP are significant only when gaseous substances participate in the reaction. If there are none, or if n = 0, then QV = QP.
The thermal effect of the reaction is understood as the amount of energy released or absorbed in the course of the reaction in the form of heat, provided: that P = const or V = const; that the temperature of the starting materials is equal to the temperature of the reaction products; that no other (useful) work is performed in the system, except for the work of the expansion and compression.
Enthalpy change in the course of various processes Process Measurement conditions Hо, k J / mol C 2 H 6 O (l) + 3 O 2 (g) → 2 CO 2 (g) + 3 H 2 O (l) P = 1 atm T = 298 K - 1 370.68 Heat of dissociation: H 2 O (l) → H + + OH- P = 1 atm T = 298 K +57. 26 Heat of neutralization: H + + OH- → H 2 O (l) P = 1 atm T = 298 K - 57.26 Heat of vaporization: H 2 O (l) → H 2 O (g) P = 1 atm T = 373 K +40. 67 Heat of fusion: H 2 O (cr) → H 2 O (l) P = 1 atm T = 273 K +6. 02
The fact of the constancy of QV or QP, long before the formation of chemical thermodynamics as a science, was experimentally established by G.I. ways to transform them into each other.
German Ivanovich Hess (1802 - 1850) - one of the largest Russian scientists, professor at the Technological Institute in St. Petersburg. Born in Geneva, and raised from an early age in St. Petersburg. He received his medical education in Yuryev, after graduation he worked in Stockholm under J. Berzelius. Hess tried in his experiments to establish the law of multiple thermal ratios (similar to D. Dalton's law of multiple ratios). This he did not succeed (there is no such law in nature), but as a result of experimental studies, Hess derived the law of constancy of the sums of heat (Hess's law). This work, published in 1842, is an anticipation of the first law of thermodynamics.
H 1 = H 2 + H 3 = H 4 + H 5 + H 6
CO 2 C + O 2 = CO 2 CO + 1/2 O 2 = CO 2 C + 1/2 O 2 = CO H 2 H 1 C CO H 3 H 1 = H 2 + H 3
Heat of formation - the heat effect of the formation of 1 mol of a given substance from simple substances: f. H. Substances are called simple if they are made up of atoms of the same kind. These are, for example, nitrogen N 2, oxygen O 2, graphite C, etc.
It follows from the definition that the heat of formation of water is equal in magnitude to the heat effect of the reaction: Н 2 + 1/2 О 2 = Н 2 О QP = f. N
If the reaction is carried out at P = 1 atm, then the measured heat of reaction will be equal to f. But - the standard heat of formation of water. Usually f. No tabulated at 298 K for almost all substances used in practice: f. No 298 (H 2 O).
Reaction products H prod f r H Initial substances H Ref. c-c f Simple substances
The heat effect of a chemical reaction: a 1 A 1 + a 2 A 2 + = b 1 B 1 + b 2 B 2 + is equal to the sum of the heats of formation of the reaction products minus the sum of the heats of formation of the starting materials (taking into account the stoichiometric coefficients ai and bj):
Example 1: Calculate the heat effect of the hydrogenation reaction of benzene vapor (this reaction is carried out on the surface of heterogeneous catalysts - platinum metals): C 6 H 6 + 3 H 2 = C 6 H 12 at 298 K and P = 1 atm:
C 6 H 6 (g) f. Ho 298, c.J / mol 82, 93 C 6 H 6 (g) 49, 04 C 6 H 12 (g) H 2 -123, 10 0 Substance r. H 0298 = -123, 10 - (82, 93 +3 0) = -206, 03 k.J r. N 0298 = -123, 10– (49, 04 + 3 0) = -72, 14 K. J isp. H 0 = 82, 93 - 49, 04 = +33, 89 k J / mol
The heat of combustion is the heat effect of the deep oxidation (combustion) reaction of a substance (to higher oxides). In the case of hydrocarbons, the higher oxides are H 2 O (l) and CO 2. In this case, the heat of combustion, for example, of methane is equal to the heat effect of the reaction: CH 4 + 2 O 2 = CO 2 + 2 H 2 O (l) QP = ox ... H
The ox. Ho 298 are called standard heats of combustion, they are tabulated at 298 K. Here, the "o" index indicates that the heats are determined at a standard state (P = 1 atm), the "ox" index comes from the English - oxidation - oxidation.
Combustion products (СО 2, Н 2 О) och. H Ref. in-in oh. H prod Reaction products r. H Starting materials
The heat effect of a chemical reaction: a 1 A 1 + a 2 A 2 + = b 1 B 1 + b 2 B 2 + is equal to the sum of the heats of combustion of the starting materials minus the sum of heats of combustion of the reaction products (taking into account the stoichiometric coefficients ai and bj):
Example 2: Using the heats of combustion of substances, calculate the heat effect of the reaction of obtaining ethanol (wine alcohol) by fermentation of glucose. C 6 H 12 O 6 = 2 C 2 H 5 OH + 2 CO 2 r. Н 0298 = 2815, 8 - 2 1366, 91 2 ∙ 0 = 81, 98 kJ The heat of combustion of СО 2 is zero.
The specific heat depends on the temperature. Therefore, a distinction is made between average and true heat capacities. The average heat capacity of the system in the temperature range T 1 - T 2 is equal to the ratio of the amount of heat supplied to the system Q to the value of this interval:
The true heat capacity is determined by the equation: The relationship between the true and average heat capacity is expressed by the equation:
The heat capacity of a system depends on its mass (or amount of matter), that is, it is an extensive property of the system. If the heat capacity is related to a unit mass, then an intensive value is obtained - the specific heat capacity of the vessel [J / kg K]. If we refer C to the amount of substance in the system, the molar heat capacity is obtained cm [J / mol K].
Distinguish: heat capacity at constant pressure Cp heat capacity at constant volume Cv. In the case of an ideal gas, the indicated heat capacities are related to each other by the equation: Ср = С v + R
The heat capacity of substances depends on temperature. For example, the heat capacity of ice varies from 34.70 J / mol K at 250 K to 37.78 J / mol K at 273 K. For solids, Debye derived an equation that, for temperatures close to 0 K, gives: CV = a T 3 (Debye's law of T-cubes), and for high ones: CV = 3 R.
Usually, the dependence of heat capacity on temperature is conveyed using empirical equations of the form: where a, b and c are const, they are given in the reference books of the physical and chemical properties of substances.
If the mathematical dependence r. CP versus T is unknown, but there are experimental values of the heat capacity of the reaction participants at different temperatures, then a graph is plotted in r coordinates. Co. P = f (T) and graphically calculate the area under the curve within 298 - T 2, it is equal to the integral:
If one or several phase transitions occur in the temperature range under consideration, then their thermal effects should be taken into account when calculating r. H:
Computation scheme for r. H reactions at an arbitrary temperature T are as follows. First, r is calculated from the standard heats of formation or heats of combustion of substances. H 298 reaction (as described above). Further, according to the Kirchhoff equation, the thermal effect is calculated at any temperature T:
The tables show the standard heats (enthalpies) of formation f for almost all substances. Ho 0 at 0 K and values: at temperature T (they are given with an interval of 100 K).
The thermal effect of a chemical reaction is calculated by the equation: r. H 0 T = r. H 00 +
r. H 00 is calculated in the same way as r. H 0298 i.e. as the difference between the sums of the heats of formation of products and initial substances (but at 0 K):
The values are calculated: = prod ref. c-c taking into account the stoichiometric coefficients of the reaction.
Physical chemistry studies the relationship
chemical processes and physical phenomena,
who accompany them, sets
patterns between chemical
composition, structure of substances and their
properties, explores the mechanism and speed
chemical reactions depending on
conditions of their course.
Physical chemistry arose and developed on the basis of the application
physical research methods to study chemical properties
substances, as well as studying the effect of the chemical composition of substances and their
structures for physical properties.
The emergence of physical chemistry as an independent
science belongs to the middle of the XVIII century.
In 1752 - 1754 - the world's first course in physical chemistry
(Lomonosov M.V.)
End of the 18th century - studies of heat capacities and thermal
reaction effects carried out by Lavoisier and Laplace (1779
- 1784)
In 1800 Berthelot introduced the concept of chemical equilibrium and
the value of the concentration of reactants.
In the first half of the XIX century. - developed atomistic
representations of Lomonosov in the works of Dalton, Gay Lussac and Avogadro
1830 - the laws of electrolysis are found (research by Devi,
Faraday, Berzelius)
1840 - Russian scientist Hess discovered the main
the law of thermochemistry. 1865 - Beketov reintroduced the teaching of the course
physical chemistry at Kharkov University.
XIX century:
Mendeleev (periodic law of 1869, as well as
gas pressure study - equation of state
ideal gas);
Guldberg and Vaage - the law of mass action;
Van't - Hoff - mathematical expression
kinetic patterns;
Menshutkin - kinetics of chemical
reactions in solutions and clarified the role of the solvent (1887
G.);
Arrhenius - the theory of electrolytic
dissociation (1887) and investigated the effect of temperature
on the rate of chemical reactions (1889).
J. Gibbs (1873 - 1878) - thermodynamic
equilibrium theory.
Le Chatelier in 1881 - 1885 formulated
usually created a quantitative theory
electrolytic dissociation. XX century:
Rutherford (1911) - nuclear model
atom.
Bohr (1913) - quantitative theory
a hydrogen atom.
Kurnakov - a new direction in
studies of multicomponent systems:
development of physical and chemical analysis - the doctrine of
dependence of the properties of physicochemical systems on
composition.
Debye and Hückel (1923) - theory
solutions of strong electrolytes.
Shilov and Semenov - chain theory
reactions and the theory of catalysis.
The main sections of physical chemistry. Their importance for pharmacy
Chemical thermodynamicsPhase equilibrium
Solutions
Electrochemistry
Kinetics and catalysis
Chemical thermodynamics. Basic concepts
Chemical thermodynamics considersenergy aspects (i.e. mutual transformations
energies associated with the transfer of energy between bodies in
the form of heat and work) of various processes and
determines the conditions for their spontaneous flow.
The subject of classical thermodynamics is the study of laws
mutual transformations of various types of energy associated with transitions
energy between bodies in the form of heat and work.
The subject of chemical thermodynamics is the application of laws
classical thermodynamics to chemical and physicochemical phenomena;
she considers the thermal effects of chemical reactions, phase
transitions of individual substances and mixtures, chemical equilibria. Object of study in thermodynamics
is a thermodynamic system.
A system is a separate body or
a group of bodies, actually or mentally
separated from the environment.
The environment is all that
is in direct or indirect
contact with the system. The system is called thermodynamic,
if between the bodies that make up it,
heat exchange can take place and
substance, and if the system is completely
described by thermodynamic
parameters.
Depending on the nature of the interaction
systems are distinguished with the environment:
An open system is ...
etc. (on one's own) The totality of all physical and chemical
properties of the system are called the state of the system.
It is characterized by thermodynamic
parameters that are:
Intense are properties that are not
depend on the mass and which equalize at
contact of systems (temperature, pressure,
density, concentration, chemical potential).
The properties of the system that depend on the mass are called
extensive (volume, mass, heat capacity,
internal energy, enthalpy, entropy,
thermodynamic potentials). Extensive
the property of the system as a whole is equal to the sum
corresponding extensive properties of individual
components included in this system
(additivity property). Those physical quantities, the value of which
completely determines the state of the system and
which lend themselves directly
measurement are called parameters
states.
The functions of these parameters are called
state functions (not amenable to
direct measurement).
State function properties:
1. The infinitesimal change in the function f is complete
differential (denoted by df).
2. Change in f during the transition of the system from state 1 to
state 2 does not depend on path df
process,
f 2 f1 a is determined
only its initial and final states:
3. As a result of the cyclical
df 0 process status function not
changes:
Thermodynamic processes and their classification
On one's own!Internal energy
Internal energy (U) characterizes the total stocksystem energy. It includes all types of energy
motion and interaction of particles that make up
system: kinetic energy of molecular motion
(translational and rotational); intermolecular energy
attraction and repulsion of particles; intramolecular or
chemical energy; energy of electronic excitation;
intranuclear and radiant energy.
The amount of internal energy depends on the nature
substance, its mass and temperature.
It is impossible to measure the full stock U (no point
reference), therefore, use a change in the internal
energy (dU or U):
U = Ufin-Uinit, J / mol.
Internal energy is a function of state, extensive
magnitude.
Enthalpy
Enthalpy is the energy possessed bysystem at constant
pressure;
enthalpy is numerically equal to the sum
internal energy and potential
system energy.
H = U + pV.
ΔН = ΔU + pΔV.
Warmth and work
The transfer of energy from the system to the environment and vice versa is carried outonly in the form of heat (Q) and work (W) -
two forms of transmission
energy.
The form of energy transfer from one part of the system to
another due to disordered
(chaotic) movement of molecules is called
warmth, but by orderly
(organized) movement of molecules under
by the action of a certain force - by work.
Work and heat are related to the process and are
functions of a process, not a state.
Measured in J / mol.
The first law of thermodynamics
Formulations:1. Energy of an isolated system
constant.
2. Energy does not disappear completely and does not
arises from nothing, its transition from one
kind in the other happens in strictly
equivalent amounts.
3. Perpetual motion machine of the first kind
impossible, which means
machine that does work without cost
energy. 4. The amount of heat supplied to the system
or taken away from her, goes to change
internal energy and work done
system or above the system.
Mathematical expression:
For final changes: Q = U + W
For infinitesimal elementary processes:
δQ = dU + δW = dU + pdV + δW ',
where δW is the sum of all types of work, pdV is mechanical work, δW ’is useful work (all,
except mechanical). Considering that δW ’0, then
pdV> δW ':
δQ = dU + pdV.
The first law of thermodynamics as applied to certain processes
1. Isothermal processes. T = const.δQ = dU + δW.
Because U = 3/2 nRT, then dU = 0 and U = 0 too.
Then: δQ = δW; δW = pdV; W = pdV.
nRT
p
From the Mendeleev - Cliperon equation V
Because
V2
nRT
W
dV nRT ln
V
V1
p
p1V1 = p2V2, then W nRT ln 1.
p2
V2
QT = WT nRT ln
V1
p1
nRT ln
p2
.
.2. Isochoric processes. V = const.
δQ = dU + δW.
δW = pdV; and since V = const, then dV = 0 and V = 0.
Then δW = pdV = 0,
and for final changes W = p V = 0.
The first law of thermodynamics in isochoric
processes will look like this:
δQV = dU
for final changes:
QV = U. 3. Isobaric processes. p = const.
δQ = dU + δW;
δW = d (pV);
δQ = dU + d (рV) or δQ = d (U + pV) = dH,
since H = U + pV.
For final changes:
QP = U + p V = N.
In the case of ideal gas, work
calculated: W = p V = nR T.
Title: Physical chemistry. Lecture notes.
This textbook is intended for students of the chemical faculties of higher educational institutions of pedagogical and technical direction. The basic concepts and processes that make up modern physical chemistry are outlined. The material complies with the state standard. The manual is recommended to help students prepare for exams.
Physical chemistry is the science that explains chemical phenomena and establishes their laws on the basis of general principles of physics.
The general task of physical chemistry is to predict the time course of a chemical process and the final result based on data on the structure and properties of molecules.
The term "physical chemistry" was proposed by MV Lomonosov. He also read the first course on his own book "Introduction to Physical Chemistry". In 1860, NN Beketov for the first time introduced physical chemistry as a special educational discipline, gave a course of lectures at Kharkov University, created the department of physical chemistry. In 1887 W. Ostwald organized the Department of Physical Chemistry at the University of Leipzig. He also publishes the first periodical on physical chemistry. A year earlier, I. A. Kablukov taught a course at Moscow University. By the end of the XIX century. defined three main sections of physical chemistry: chemical thermodynamics, chemical kinetics and electrochemistry.
At present, physical chemistry has fully developed as a science, which includes chemical thermodynamics (thermochemistry, phase equilibrium), complementing chemical kinetics with catalysis, and has also created a variety of physicochemical methods of analysis.
Table of contents
Introduction
LECTURE No. 1. Ideal gas. Real gas equation of state
1. Elements of molecular kinetic theory
2. Equation of state of ideal gas
3. Kinetic theory of gases
4. Equation of state of real gas
LECTURE No. 2. Chemical thermodynamics
1. Systems and their classification
2. Thermodynamic parameters. Thermodynamic indicators. Voltage balance
3. The first law of thermodynamics. Caloric coefficients. Relationship between CP and Cv functions
4. Isoprocesses in thermodynamics. Helmholtz energy
5. Processes. The second law of thermodynamics
6. Carnot cycle
7. Impossibility of a perpetual motion machine
LECTURE No. 3. Solutions
1. General characteristics of solutions
2. Concentration and ways of expressing it
3. Solubility of gases in liquids
4. Solutions of non-electrolytes. Raoult's law and its consequences
5. Osmosis
6. Fugacity
7. Henry's Law
LECTURE No. 4. Catalysis
1. History of the discovery of the phenomenon of catalysis
2. The mechanism of catalytic interaction. Types of catalysts
LECTURE No. 5. Chemical equilibrium
1. The concept of chemical equilibrium. Mass action law
2. Equation of the isotherm of a chemical reaction
3. Equations of isochores, isobars of a chemical reaction
4. Calculation of KP (Temkin-Shvartsman method)
5. Calculation of the equilibrium composition of chemical equilibrium
LECTURE No. 6. Chemical kinetics
1. The concept of chemical kinetics
2. Factors affecting the rate of chemical reaction
LECTURE No. 7. Corrosion of metals
1. Basic concepts and terminology
2. Classification of metal corrosion processes
3. Types of corrosion damage
4. Methods of protection against corrosion
LECTURE No. 8. Physical and chemical analysis
1. The essence of physical and chemical analysis
2. One-component systems
3. Physicochemical methods for analyzing the composition of alloys
LECTURE No. 9. Thermochemistry
1. The concept of thermochemistry
2. Hess's law
3. Kirchhoff's law. Integral form of Kirchhoff equations
LECTURE No. 10. Galvanic cells
1. The concept of a galvanic cell
2. Chemical power sources
3. Regeneration and utilization of CPS
LECTURE No. 11. Electrochemistry
1. The concept of electrochemistry
2. Electrode processes
3. Cathodic and anodic processes in electroplating
4. Modern trends in the development of thermodynamic and applied electrochemistry
LECTURE No. 12. Theoretical electrochemistry
1. Associations in electrolyte solutions. The concept of the theory of strong electrolytes. Activity
2. Thermodynamics of electrolyte solutions. DES types
3. Modern approaches to the description of the thermodynamic properties of electrolyte solutions
4. Thermodynamic characteristics of ions in electrolyte solutions
5. Nonequilibrium phenomena in the ionic system
6. Equilibrium in the liquid-liquid system
7. The concept of DES. Model concepts of the structure of DES at the interface
8. Conductors of the first and second kind
9. Reference electrodes
LECTURE No. 13. Electrochemical kinetics
1. Basic kinetic characteristics and methods of their calculation
2. Equations of electrochemical kinetics, limits of their applicability
3. Kinetic features of electrodeposition of metals and alloys
4. Influence of the nature of the solvent on the rate of electrochemical reactions
5. Electroosmosis
6. Electrocapillary curves
7. Electrochemical overvoltage (charge transfer overvoltage)
8. Factors affecting hydrogen overvoltage. Oxygen overvoltage
LECTURE No. 14. Application of theoretical and applied electrochemistry
1. Applied electrochemistry
2. Electrochemistry of carbon
3. Bioelectrochemistry
4. Stochastic processes and self-organizing systems
5. Investigation of the phenomenon of high-temperature superconductivity in oxides of complex composition
6. Simulation of electrochemical processes
7. Method of galvanostatic curves
LECTURE No. 15. The third law of thermodynamics
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Description of the presentation for individual slides:
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Slide Description:
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Gas is an aggregate state of matter, in which molecules move chaotically, located at a great distance from each other. In solids, the distances between particles are small, the force of attraction corresponds to the force of repulsion. Liquid is a state of aggregation, intermediate between solid and gaseous. In a liquid, particles are located close to one another and can move relative to each other; a liquid, like a gas, has no definite shape. Plasma is a highly discharged gas in which electrically charged particles are chaotically moving - electrons and positively charged nuclei of atoms or ions.).
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The aggregate states of the same substance do not differ in chemical properties and composition, and their physical properties are not the same. An example is H2O (water). Differences in physical properties are due to the fact that particles in gaseous, liquid and solid things are located at an unequal distance from each other, due to which the forces of attraction acting between them are manifested in an unequal degree
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The main provisions of the ICT All substances - liquid, solid and gaseous - are formed from the smallest particles - molecules, which themselves consist of atoms ("elementary molecules"). The molecules of a chemical can be simple or complex and consist of one or more atoms. Molecules and atoms are electrically neutral particles. Under certain conditions, molecules and atoms can acquire an additional electrical charge and turn into positive or negative ions. Atoms and molecules are in continuous chaotic motion. The particles interact with each other by forces of an electrical nature. The gravitational interaction between particles is negligible.
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1. The doctrine of aggregate states 1.1 Introduction Phase transition - the transition of a substance from one state of aggregation to another L-G vaporization (evaporation) T-G sublimation (sublimation) G-F liquefaction (condensation) solid and liquid G-T desublimation (condensation) state - condensed T-W melting W-T solidification (freezing) Phase transitions are accompanied by absorption or release of heat
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1. The doctrine of states of aggregation 1.2. Gaseous state of matter Gas is the aggregate state of matter in which its constituent particles (atoms, molecules, ions) are not bound or are very weakly bound by the forces of interaction, move freely, filling the entire volume provided to them. The main characteristics of gases: have a low density, because particles far apart from each other have neither their own shape nor their own volume; they completely fill the vessel in which they are located, and take its shape and are easily compressed.
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Equation of state of an ideal gas Ideal gas is a theoretical model of a gas in which the size and interaction of gas particles are neglected and only their elastic collisions are taken into account or Ideal gas is a gas in which there are no forces of attraction between molecules.
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gas particles (atoms, molecules, ions) are taken as material points (i.e., have no volume) between the particles there are no forces of mutual attraction (intermolecular forces), the interaction between molecules is reduced to absolutely elastic impacts (i.e., impacts in which the kinetic energy is completely transferred from one object to another) gas particles (atoms, molecules, ions) have a volume of gas particles are interconnected by forces of interaction, which decrease with increasing distance between particles, collisions between molecules are not absolutely elastic Ideal gas Real gas 1. The doctrine of aggregate states 1.2. Gaseous state of matter A real gas is similar to an ideal gas under strong rarefaction and at ordinary temperatures
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The equation of state for an ideal gas (Mendeleev-Clapeyron equation) is a relationship between the values of pressure, volume and temperature: where n is the number of moles of gas, R = 8.31431 J / mol.K) is the gas constant Gas obeying this law, called ideal. Gas laws
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Gas laws At constant temperature and mass, the volume of a gas is inversely proportional to its pressure The volume of a given mass of gas at constant pressure is directly proportional to the absolute temperature The pressure of a given mass of gas at a constant volume is directly proportional to the absolute temperature Boltzmann constant: k = R / NA = 1.38 10-23 J / C
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Ideal gases have the same molar volume. at n. at. = 22.4140 dm3 (l) At other temperatures and pressures, this value will be different! Gas laws
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They do not obey the laws of ideal gases. The main reasons for the deviations are the mutual attraction of gas molecules and the presence of their own volume.The molar volume can serve as a characteristic of deviations. Real gases
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Real gases Real gases do not obey the Mendeleev – Clapeyron equation. The equation of state of a real gas (van der Waals equation) for one mole for n moles a - takes into account intermolecular interactions; b - takes into account the intrinsic volume of molecules. Coefficients a and b are different for different gases, so the van der Waals equation is not universal. At low pressures and high temperatures, the van der Waals equation transforms into the equation of state for an ideal gas.
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The main property of a liquid that distinguishes it from other aggregate states is the ability to change its shape indefinitely under the action of tangential mechanical stresses, even arbitrarily small, while practically maintaining the volume. A liquid state is usually considered intermediate between a solid and a gas: a gas does not retain either volume or shape, while a solid retains both. The liquid state of the thing
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vibrational-translational motion of molecules, incompressibility due to internal pressure, association (in the case of polar molecules), the presence of short-range order in the absence of long-range order, surface tension, viscosity. Properties of liquids:
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