Physical chemistry. Lecture notes
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Gas is a state of aggregation of a substance in which molecules move randomly, 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 close to each other and can move relative to each other; A liquid, like a gas, has no fixed shape. Plasma is a highly rarefied gas in which randomly moving electrically charged particles are electrons and positively charged nuclei of atoms or ions.).
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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 substances are located at unequal distances from each other, due to which the forces of attraction acting between them manifest themselves to an unequal degree
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Basic Provisions of MKT All substances - liquid, solid and gaseous - are formed from the smallest particles - molecules, which themselves consist of atoms ("elementary molecules"). Molecules of a chemical substance 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. Particles interact with each other by forces that are electrical in 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 state - condensed T-L melting L-T solidification (freezing) Phase transitions are accompanied by absorption or release of heat
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1. The doctrine of aggregate states 1.2. Gaseous state of matter Gas is the state of aggregation of matter in which its constituent particles (atoms, molecules, ions) are not bound or very weakly bound by interaction forces, move freely, filling the entire volume provided to them. The main characteristics of gases: they have a low density, because particles are far apart and have neither their own shape nor their own volume; they completely fill the vessel in which they are located, and take its form and are easily compressed.
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The ideal gas equation of state An ideal gas is a theoretical gas model in which the size and interaction of gas particles are neglected and only their elastic collisions are taken into account. 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 (that is, they have no volume) there are no forces of mutual attraction between the particles (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 gas particles are interconnected by interaction forces that 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. The 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 of an ideal gas (the Mendeleev-Clapeyron equation) is a relation that relates the values of pressure, volume and temperature: where n is the number of moles of gas, R = 8.31431 J / mol.K) - 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 constant volume is directly proportional to the absolute temperature Boltzmann's constant: k=R/NA=1.38 10-23 J/K
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Ideal gases have the same molar volume. at n. y. = 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 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. The coefficients a and b for different gases are different, so the van der Waals equation is not universal. At low pressures and high temperatures, the van der Waals equation becomes the equation of state for an ideal gas.
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The main property of a liquid, which distinguishes it from other states of aggregation, is the ability to change its shape indefinitely under the action of tangential mechanical stresses, even arbitrarily small, while practically maintaining volume. The liquid state is usually considered intermediate between a solid and a gas: a gas retains neither volume nor shape, while a solid retains both. Liquid state of matter
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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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D. x. n. , Professor, Head of the Department of Physical Chemistry, Russian Chemical Technical University named after. D. I. Mendeleeva Konyukhov Valery Yurievich [email protected] ru vkontakte. en
Literature Vishnyakov A.V., Kizim N.F. Physical chemistry. Moscow: 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: Proc. Allowance for universities / V. V. Eremin et al. M. : 2005.
Literature Atkins P. Physical chemistry. M. : Mir. 1980. Karapetyants M.Kh. Chemical thermodynamics. Moscow: Chemistry, 1975.
LOMONOSOV Mikhail Vasilievich (1711-65), the first Russian natural scientist of world significance, a poet who laid the foundations of the modern Russian literary language, an artist, a historian, an advocate 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 (since 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 adjunct, from 1745 academician of the St. Petersburg Academy of Sciences.
In 1748 he founded the first chemical laboratory in Russia at the Academy of Sciences. On the initiative of Lomonosov, Moscow University was founded (1755). He developed atomic and molecular ideas about the structure of matter. During the dominance of the theory of caloric, he argued that heat is due to the movement of corpuscles. He formulated the principle of conservation of matter and motion. Excluded phlogiston from the number of chemical agents. Laid the foundations of physical chemistry.
Investigated atmospheric electricity and gravity. Proposed the doctrine of color. Created a number of optical instruments. Discovered the atmosphere on Venus. Described the structure of the Earth, explained the origin of many minerals and minerals. Published a guide to metallurgy. He emphasized the importance of exploring the Northern Sea Route, the development of Siberia. He revived the art of mosaics and the production of smalt, created mosaic paintings with his students. Member of the Academy of Arts (1763). He was 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 may be called chemical philosophy.
In Western Europe, it is customary to consider 1888 the year of the creation of physical chemistry, 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 the age of 35 he changed 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 "Volume-chemical and optochemical research".
In 1887, he accepted an offer to move to Leipzig, where he founded the Institute of Physics and Chemistry at the university, which he directed until 1905. In 1888, he occupied the very prestigious Department of Physical and Inorganic Chemistry at the University of Leipzig. He worked in this position for 12 years.
From the "Leipzig School" of W. Ostwald came: Nobel laureates S. Arrhenius, J. Van't Hoff, W. Nernst, famous physicochemists G. Tamman and F. Donnan, organic chemist J. Wislicens, the famous American chemist G. N. Lewis. Over the years, Russian chemists trained at Ostwald: 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 explores only simple and understandable laws that govern 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 members of the Chemical Society, Ostwald tried to prove that atoms do not exist, and "what we call matter is only 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 a reaction can proceed spontaneously; if the reaction proceeds, then how deep (what are the equilibrium concentrations of the reaction products); If the reaction proceeds, then at what rate.
1. STRUCTURE OF SUBSTANCE In this section, on the basis of quantum mechanics (Schrödinger's equation), the structure of atoms and molecules (electronic orbitals of atoms and molecules), crystal lattices of solids, etc. is explained, aggregate states of matter are considered.
2. CHEMICAL THERMODYNAMICS based on the laws (beginnings) of thermodynamics allows: to calculate the thermal effects of chemical reactions and physico-chemical processes, to predict the direction of chemical reactions, to calculate the equilibrium concentrations of reactants and reaction products.
3. THERMODYNAMICS OF PHASE EQUILIBRIUM He studies the regularities of phase transitions in one-component and multicomponent (solution) systems. Its main purpose is to construct phase equilibrium diagrams for these systems.
4. ELECTROCHEMISTRY It studies the properties of electrolyte solutions, the features of their behavior in comparison with molecular solutions, explores the patterns of interconversion of the energy of chemical reactions and electrical energy during the operation of electrochemical (galvanic) cells and electrolyzers.
5. CHEMICAL KINETICS AND CATALYSIS Investigates the regularities of the course of chemical reactions in time, investigates the effect of thermodynamic parameters (pressure, temperature, etc.), the presence of catalysts and inhibitors on the rate and mechanism of reactions.
In a separate science, COLLOID CHEMISTRY is distinguished by a section of physical chemistry - 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 - movement).
Thermodynamics abstracts from the causes that cause any process, and the time during which this process occurs, but only operates with the initial and final parameters of the system involved in any physical and chemical 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: measurement and calculation of thermal effects of chemical reactions and physicochemical processes, prediction of the direction and depth of reactions, analysis of 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. System - a body or a group of bodies, actually or mentally identified by an observer in the environment.
The system is the part of the surrounding world that we are particularly interested in. 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.
According to the nature of the exchange of energy and matter with the environment, systems are classified: isolated - they cannot exchange either matter or energy; closed - can exchange energy, but cannot - matter; open - can exchange both matter and energy.
According to the number of phases, the systems are divided into: homogeneous - consist of one phase (Na. Cl solution 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 (fat droplets - one phase, the aquatic environment - 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 phase interfaces. Each phase is a homogeneous part of a heterogeneous system
According to the number of components, the systems are divided into one-two-, three-component and multi-component. 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 on certain values: temperature, pressure, thermal conductivity, heat capacity, component concentrations, dielectric constant, etc.
In chemical thermodynamics, one deals with those properties that can be unambiguously expressed as functions 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 given if its chemical composition, phase composition, and values of independent thermodynamic parameters are indicated. The 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). They 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 an off-system unit of pressure, the normal physical atmosphere (atm), equal to 101. 325 k. Pa
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. When a complex system is formed, intensive properties are aligned, and extensive ones are summed up.
Any change that occurs in the system and is accompanied by a change in at least one thermodynamic state parameter (system properties) 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, any one (or several) parameters are kept constant. Accordingly, they distinguish: an isothermal process at a constant temperature (T = const); isobaric process - at constant pressure (P = const); isochoric process - at a constant volume (V = const); adiabatic process in the absence of heat exchange with the environment (Q = 0).
When processes occur in non-isolated systems, heat can be absorbed or released. In accordance with this characteristic, processes are divided into exothermic (heat is released) and endothermic (heat is absorbed).
During the process, the system passes from one equilibrium state to another equilibrium state. Thermodynamic equilibrium is the state of the system in which thermal, mechanical and chemical (electrochemical) equilibrium with the environment and between the phases of the system is observed.
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 system properties (internal energy U, enthalpy H, entropy S, etc.), they characterize the given state of the system. Their changes during the process do not depend on its path and are determined only by the initial and final states of the system.
An infinitesimal change in this function is the total differential of 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 during the process in which the system participates.
If there is no heat and work in the system, then it is meaningless 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. Infinitely small quantities are denoted by Q, W.
Movement is an attribute of matter. The measure of movement, i.e., the quantitative and qualitative characteristic, is energy. 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 energy can pass from system to system in only 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 taken as positive (for example, the heat of an endothermic reaction), and the heat removed from the system is 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-bodies. In the literature, work is denoted either by W (from the English “work”) or A (from the German “arbait”).
There are different types of work: mechanical, electrical, magnetic, surface changes, etc. An infinitely small 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 the forces of external pressure P - work of expansion - compression, 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 the IUPAC recommendations, it is customary to consider the work done on the system as positive (the “egoistic” principle is positive that 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 relations of thermodynamics are formulated on the basis of two postulates and three laws. Any isolated system eventually 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 (
The spontaneous transition from a non-equilibrium state to an equilibrium state is called relaxation. That is, the equilibrium state will necessarily be achieved, but the duration of such a process is not defined, and there is no concept of time.
The second postulate If system A is in thermal equilibrium with system B, and system B is in thermal equilibrium with system C, then systems A and C are also in thermal equilibrium
The internal energy of any thermodynamic system U is the sum of the kinetic (motion energy) 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 substance of the system and thermodynamic parameters: U = f(V, T) or U = (P, T) is measured in J/mol or J/kg. U is a state function, so U does not depend on the path of the process, but is determined by the initial and final state of the system. d. U is the total 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 infinitesimal part of the process): d. U = QW
"The heat supplied to the system goes to increase the internal energy of the system and the performance of work by the system."
For an isolated system, Q = 0 and W = 0, i.e., 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 quantities. Energy does not arise and is not destroyed, but only passes from system to system.
The function U is additive. This means that if two systems characterized by the values 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, i.e., its amount depends on the path of the process, but in two cases important for practice, heat acquires the properties of a state function, i.e., the value of Q ceases to depend on the path of the process, and is determined only initial and final states of the system.
We assume that in the course of the process only work against the forces of external pressure can be performed, 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 \u003d Uk - Un
Again we assume that the useful work W = 0, then: Q = d. U+P d. V, Since P = const, we can write: QP = d. U + d(PV), QP = d(U + P V). Denote: Н U + P V (enthalpy) QР = d. H or: QP \u003d Hk - Hn
Thus, the thermal effect of a chemical reaction acquires the properties of a state function at P = const: QP = H; for V = const: QV = U.
Since chemical reactions and physicochemical processes are more often carried out at a constant pressure (in the open air, that is, at P = const = 1 atm), in practice, the concept of enthalpy is more often used for calculations, rather than internal energy. Sometimes the word "heat" of the process is replaced without further explanation by "enthalpy", and vice versa. For example, they say “heat of formation”, but write f. N.
But if the process of interest to us proceeds at V = const (in an autoclave), then the expression should be used: QV = U.
Let's differentiate the expression: Н = 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: P V \u003d n R T, where n is the number of moles of 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 is 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) \u003d 2 NH 3 (g) n \u003d -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.
Under the thermal effect of the reaction understand the amount of energy released or absorbed during 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 work (useful) is done in the system, except for the work of expanding contraction.
Enthalpy change during various processes Process Measurement conditions Ho, 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 evaporation: 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. Hess (the law of constancy of heat sums or Hess's law): The thermal effect of a chemical reaction depends on the type and state of the starting substances and reaction products and does not depend on 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 brought up from an early age in St. Petersburg. He received his medical education in Yuriev, after graduating from the university he worked in Stockholm with J. Berzelius. Hess tried in his experiments to establish the law of multiple thermal ratios (similar to the law of multiple ratios of D. Dalton). He did not succeed in this (there is no such law in nature), but as a result of experimental studies, Hess deduced the law of constancy of heat sums (Hess's law). This work, published in 1842, is an anticipation of the first law of thermodynamics.
H 1 \u003d H 2 + H 3 \u003d H 4 + H 5 + H 6
CO 2 C + O 2 \u003d CO 2 CO + 1/2 O 2 \u003d CO 2 C + 1/2 O 2 \u003d CO H 2 H 1 C CO H 3 H 1 \u003d 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. Simple substances are called substances consisting of atoms of the same type. This, for example, is nitrogen N 2, oxygen O 2, graphite C, etc.
It follows from the definition that the heat of water formation is equal in magnitude to the thermal effect of the reaction: H 2 + 1/2 O 2 = H 2 O QP = f. H
If the reaction is carried out at P = 1 atm, then the measured heat of reaction will be equal to f. Ho is the standard heat of formation of water. Usually the values of f. But tabulated at 298 K for almost all substances used in practice: f. Ho 298(H 2 O).
Reaction products H prod f r H Starting materials H Ref. c-c f Simple substances
The thermal 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 substances (taking into account the stoichiometric coefficients ai and bj):
Example 1: Calculate the heat effect of the benzene vapor hydrogenation reaction (this reaction is carried out on the surface of heterogeneous catalysts - platinum metals): C 6 H 6 + 3 H 2 \u003d C 6 H 12 at 298 K and P \u003d 1 atm:
C 6 H 6(g) f. Ho 298, q. 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 \u003d -123.10 - (82.93 +3 0) \u003d -206.03 k. J r. H 0298 \u003d -123, 10– (49, 04 + 3 0) \u003d -72, 14 k. J isp. H 0 \u003d 82.93 - 49.04 \u003d +33.89 k. J / mol
The heat of combustion is the thermal effect of the reaction of deep oxidation (combustion) 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 calorific value of, for example, methane is equal to the thermal effect of the reaction: CH 4 + 2 O 2 \u003d CO 2 + 2 H 2 O (l) QP \u003d ox . H
ox values. Ho 298 is called the standard heat of combustion, they are tabulated at 298 K. Here the index "o" indicates that the heats are determined at the standard state (P \u003d 1 atm), the index "oh" comes from English - oxidation - oxidation.
Combustion products (CO 2, H 2 O) oh. H Ref. in-in oh. N prod Reaction products r. H Starting materials
The thermal 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 substances minus the sum of the heats of combustion of the reaction products (taking into account the stoichiometric coefficients ai and bj):
Example 2: Using the heat of combustion of substances, calculate the heat effect of the reaction for producing ethanol (wine alcohol) by fermenting glucose. C 6 H 12 O 6 \u003d 2 C 2 H 5 OH + 2 CO 2 r. H 0298 \u003d 2815.8 - 2 1366.91 2 ∙ 0 \u003d 81.98 kJ The heat of combustion of CO 2 is zero.
The heat capacity depends on the temperature. Therefore, a distinction is made between the 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 capacities is expressed by the equation:
The heat capacity of a system depends on its mass (or amount of matter), i.e., this is an extensive property of the system. If the heat capacity is attributed to a mass unit, then an intensive value is obtained - the specific heat capacity court [J / kg K]. If, however, we attribute C to the amount of the substance of the system, we get the molar heat capacity cm [J / mol K].
There are: heat capacity at constant pressure Cp; heat capacity at constant volume Cv. In the case of an ideal gas, these heat capacities are interconnected 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: СV= a T 3 (Debye's law of T-cubes), and for high ones: СV=3 R.
Usually, the dependence of heat capacity on temperature is transmitted using empirical equations of the form: where a, b and c are const, they are given in reference books of the physicochemical properties of substances.
If the mathematical dependence r. CP from 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 the coordinates r. Co. P \u003d f (T) and graphically calculate the area under the curve within 298 - T 2, it is equal to the integral:
If one or more phase transitions occur in the temperature range under consideration, then their thermal effects should be taken into account when calculating r. H:
Calculation scheme r. H reactions at an arbitrary temperature T is 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. H00+
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 starting materials (but at 0 K):
The values are calculated: = prod ref. in-in, taking into account the stoichiometric coefficients of the reaction.
