General chemistry 0600-S1-O-PC

Lectures:
BASIC CHEMICAL LAWS AND DEFINITIONS
Weight, length and temperature. Mol and Avogardo number, accuracy, precision, significant digits. Rounding numbers. Basic units and SI system. Calculation: unit conversion. The law of conservation of mass and the law of constant and multiple proportions. Balancing chemical equations. Chemical symbols. Stoichiometry. Concentration units.
ATOM – NUCLEUS
Dalton's atomic theory. Rutherford's experience. The structure of the atom: the nucleus and electrons. Elementary particles: electrons, neutrons, protons. Atomic number. Mass number and mole of the substance. Nuclear chemistry: transformation of elements. Natural radioactivity: alpha, beta and gamma. Nuclear stability: mass deficit and nucleus binding energy. The origin of chemical elements. Nuclear fusion and division of atomic nuclei. Atom bomb. Nuclear energy.
ATOM – ELECTRONIC STRUCTURE AND CHEMICAL BONDS
Bohr model. Wave-corpuscular dualism. Heisenberg's uncertainty principles. Compton phenomenon. Schroedinger equation. Wave functions and quantum numbers. Shapes of molecular orbitals. Emission spectra. Pauli's prohibition and Hundt’s rule. Energy of atomic orbitals in multi-electron atoms. The periodic table of elements and the electron configuration of elements. Ionic radius. The first and higher ionization potentials. Electron affinity. Octet rule. Ionic bond. Covalent molecules and bonds. The power of covalent bonds. Polarized covalent bonds. Electronegativity. Dot patterns. The theory of valence bonds. Hybridization and hybridized orbitals. Hydrogen molecule and other diatomic molecules. Bond order.
GASES AND LIQUIDS
Gaseous state and pressure of gases. Gas laws. Ideal gas equation. Partial pressure and Dalton's law. Kinetic-molecular theory of gases. Gas diffusion. Behavior of real gases. Critical temperature. Polarized bonds and the dipole moment. Intermolecular interactions. Properties of the liquid. Phase transitions: evaporation-condensation, melting-freezing, sublimation-resublimation. Gibbs phase rule. Phase diagram. Steam pressure and boiling point. Surface tension. Surfactants. Flotation. Solubility of gases in liquids.
SOLIDS
Types of solids: crystalline and non-crystalline. Ionic, covalent, metallic and molecular crystals. Allotropy. Study of the structure of solids: X-ray diffraction. Elementary cells. Liquid crystals. Metallic elements. Semiconductors and their applications.
SOLUTIONS
Electrolytes in aqueous solutions. Acid-base theories: Arrhenius and Broensted-Lowry. The strength of acids and bases. Proton hydration and oxonium ion. Self-ionization of water. pH scale. pH of acid and base solutions. Relationship between Ka and Kb. Acid-base properties of salts. Factors affecting the strength of acids and bases. The effect of a common ion. Buffer solutions. Indicators.
Solubility of ionic compounds. Factors affecting solubility. Precipitation of ionic compounds. Calculation of solubility using Ksp. Ion separation by selective precipitation. Qualitative analysis. Lowering the vapour pressure over solutions: Raoult's law. Boiling point increase and freezing point lowering. Osmosis and osmotic pressure. Dialysis.
CHEMICAL KINETICS
Reaction rate. Kinetic equations and the order of reactions. Half-time and first-order reactions. Radioactive decay. Second-order reactions. Reaction rate and temperature: The Arrhenius equation. Catalysis. Homogeneous and heterogeneous catalysts. Enzymatic catalysis. Equilibrium state. Equilibrium constants Kc and Kp. The use of constant equilibrium. Factors affecting the composition of the reaction mixture: Le Chatelier's rule. Synthesis of ammonia.
REDOX REACTIONS AND ELECTROCHEMISTRY
Galvanic cells. Short notation of galvanic cells. Electrode types. Electromotive force of cells. Standard potentials - Nernst equation. Electrochemical pH determination. Batteries. Fuel cells. Concentration cells. Corrosion. Electrolysis and electrolytic cells. Practical applications of electrolysis. Balancing Redox reaction equations: half reactions.
ELEMENTS OF TERMOCHEMISTRY AND THERMODYNAMICS
Energy and energy conservation. Internal energy and state functions. Energy and enthalpy. Enthalpies of chemical and physical transformations. Hess's Law. Standard creation enthalpies. Introduction to the concept of entropy.

Exercise:
SI base units. Concentration units of solutions and mixtures. Significant numbers. Rounding numbers. The law of conservation of mass. Chemical symbols. Balancing of chemical equations with particular emphasis on redox reaction equations. Computational solving of chemical problems in the field of: the content of elements in compounds and mixtures, concentrations of solutions, gas laws, pH of acid and base solutions, pH of buffer solutions, precipitation and dissolution reactions of chemical compounds, determination of solubility of chemical compounds in solution and product of solubility, conditions of quantitative precipitation of chemical compounds, calculation of electrode potentials and electromotive force of galvanic cells.

Laboratory

PART ONE – PARALLEL WORK ("EQUAL FRONT")

LABORATORY 1: ORGANIZATION OF WORK IN THE LABORATORY. LABORATORY TECHNIQUE.
Order regulations. Safety regulations. Basic laboratory equipment. Measuring liquids. Washing glassware. Heating of liquids. Dissolution and digestion. Evaporation and crystallization.
LABORATORY 2: EQUILIBRIUM IN ELECTROLYTE SOLUTIONS
Chemical equilibrium - the law of mass action. Strong and weak electrolytes. Concentration and activity. The role of the solvent. pH scale. Brönsted's theory. Solutions of acids and bases. Constant and degree of protolysis. Ostwald dilution law. Indicators of pH. pH measurement - indicator and reference electrodes.
LABORATORY 3: CONSTANT AND DEGREE OF DISSOCIATION OF WEAK ELECTROLYTES. BUFFER SOLUTIONS.
Equilibrium in solutions of weak electrolytes. Disturbing the state of equilibrium. The rule of defiance. Buffer solutions: composition, pH value, pH stabilization range, buffer capacity, use of buffer solutions.
LABORATORY 4: PRECIPITATION AND DISSOLUTION OF SEDIMENTS.
Saturated solutions. Solubility: definition and units. Factors affecting solubility. Effect of the common ion on the solubility of sediments. Application of the defiance rule in the process of precipitation and dissolution of sediments. Application of precipitation and dissolution of sediments in qualitative analysis.
LABORATORY 5: COORDINATION COMPOUNDS
Coordination binding. Conditions for the formation of complex ions. Structure of complex ions. Naming of complex compounds. Constants of stability and instability of complex ions. Aqua complexes.
LABORATORY 6: OXIDATION AND REDUCTION (REDOX) PROCESSES. ELEMENTS OF ELECTROCHEMISTRY
Oxidation and reduction reactions as a distinct type of chemical reactions. The concept of the degree of oxidation. Record of redox reaction equations (half reactions in ionic form). Half-cell potential - Nernst equation. Hydrogen electrode standard, its importance for the description of oxidation and reduction processes. Voltage series of metals. Oxidation-reduction series. Influence of temperature and pH on the course of oxidation and reduction reactions. Reaction rate. Passivation

PART TWO – INDIVIDUAL WORK
ANIONS:
CO32-, C2O42-, CH3COO-, NO3-, NO2-, SO42-, SO32-, S2O32-, S2-, SCN-, Cl-, Br-, I-, Cr2O72-, CrO42-, PO43-, MnO4-, OH-

1. Anions: preliminary tests. pH - measurement with universal paper. Reactions of anions (1-2 drops of the test solution) with AgNO3 (1-2 drops).
2. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with 1M H2SO4 (2-3 drops).
3. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KMnO4 (1 drop) against 1M H2SO4 (1-2 drops).
4. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KI (1 drop) against 1M H2SO4 (1 drop).
5. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with I2 (1 drop) against 1M H2SO4 (1 drop).
6. Anions: characteristic reactions.

CATIONS:
Cations group I: preliminary tests. Precipitation reaction of 1M HCl chlorides and their dissolution with excess conc. HCl. Take about 2-3 drops of cation solutions each.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NaOH. Take about 2 to 3 drops of cation solutions.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NH3. Take about 2 to 3 drops of cation solutions.
Group I cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Precipitation of sulphides with thioacetamide solution in an environment of 0.3M HCl. Dissolution of sulphides in 2M NaOH and in a mixture of 2M NaOH and 3% H2O2 (10:3) . Take about 2 to 3 drops of cation solutions.
Group II cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions. Group III cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with ammonium buffer 2M NH4Cl/2M NH3 (10:1) and attempt to precipitate sulphides with thioacetamide solution. Precipitation sulphides only in the absence of reaction with ammonium buffer. Take about 2 to 3 drops of cation solutions. Group III cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Carbonate precipitation reaction with solution (NH4)2CO3 and their dissolution in 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Chromate precipitation reaction with K2CrO4 solution in an inert and acidified medium of 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Sulphate precipitation reaction with 2M H2SO4 solution Take about 2 to 3 drops of cation solutions each.
Group IV and V cations: Characteristic reactions. Take about 2-3 drops of cation solutions
Self-detection of anions and cations (evaluated) in administered samples (aqueous solutions and powders).

Lectures:
BASIC CHEMICAL LAWS AND DEFINITIONS
Weight, length and temperature. Mol and Avogardo number, accuracy, precision, significant digits. Rounding numbers. Basic units and SI system. Calculation: unit conversion. The law of conservation of mass and the law of constant and multiple proportions. Balancing chemical equations. Chemical symbols. Stoichiometry. Concentration units.
ATOM – NUCLEUS
Dalton's atomic theory. Rutherford's experience. The structure of the atom: the nucleus and electrons. Elementary particles: electrons, neutrons, protons. Atomic number. Mass number and mole of the substance. Nuclear chemistry: transformation of elements. Natural radioactivity: alpha, beta and gamma. Nuclear stability: mass deficit and nucleus binding energy. The origin of chemical elements. Nuclear fusion and division of atomic nuclei. Atom bomb. Nuclear energy.
ATOM – ELECTRONIC STRUCTURE AND CHEMICAL BONDS
Bohr model. Wave-corpuscular dualism. Heisenberg's uncertainty principles. Compton phenomenon. Schroedinger equation. Wave functions and quantum numbers. Shapes of molecular orbitals. Emission spectra. Pauli's prohibition and Hundt’s rule. Energy of atomic orbitals in multi-electron atoms. The periodic table of elements and the electron configuration of elements. Ionic radius. The first and higher ionization potentials. Electron affinity. Octet rule. Ionic bond. Covalent molecules and bonds. The power of covalent bonds. Polarized covalent bonds. Electronegativity. Dot patterns. The theory of valence bonds. Hybridization and hybridized orbitals. Hydrogen molecule and other diatomic molecules. Bond order.
GASES AND LIQUIDS
Gaseous state and pressure of gases. Gas laws. Ideal gas equation. Partial pressure and Dalton's law. Kinetic-molecular theory of gases. Gas diffusion. Behavior of real gases. Critical temperature. Polarized bonds and the dipole moment. Intermolecular interactions. Properties of the liquid. Phase transitions: evaporation-condensation, melting-freezing, sublimation-resublimation. Gibbs phase rule. Phase diagram. Steam pressure and boiling point. Surface tension. Surfactants. Flotation. Solubility of gases in liquids.
SOLIDS
Types of solids: crystalline and non-crystalline. Ionic, covalent, metallic and molecular crystals. Allotropy. Study of the structure of solids: X-ray diffraction. Elementary cells. Liquid crystals. Metallic elements. Semiconductors and their applications.
SOLUTIONS
Electrolytes in aqueous solutions. Acid-base theories: Arrhenius and Broensted-Lowry. The strength of acids and bases. Proton hydration and oxonium ion. Self-ionization of water. pH scale. pH of acid and base solutions. Relationship between Ka and Kb. Acid-base properties of salts. Factors affecting the strength of acids and bases. The effect of a common ion. Buffer solutions. Indicators.
Solubility of ionic compounds. Factors affecting solubility. Precipitation of ionic compounds. Calculation of solubility using Ksp. Ion separation by selective precipitation. Qualitative analysis. Lowering the vapour pressure over solutions: Raoult's law. Boiling point increase and freezing point lowering. Osmosis and osmotic pressure. Dialysis.
CHEMICAL KINETICS
Reaction rate. Kinetic equations and the order of reactions. Half-time and first-order reactions. Radioactive decay. Second-order reactions. Reaction rate and temperature: The Arrhenius equation. Catalysis. Homogeneous and heterogeneous catalysts. Enzymatic catalysis. Equilibrium state. Equilibrium constants Kc and Kp. The use of constant equilibrium. Factors affecting the composition of the reaction mixture: Le Chatelier's rule. Synthesis of ammonia.
REDOX REACTIONS AND ELECTROCHEMISTRY
Galvanic cells. Short notation of galvanic cells. Electrode types. Electromotive force of cells. Standard potentials - Nernst equation. Electrochemical pH determination. Batteries. Fuel cells. Concentration cells. Corrosion. Electrolysis and electrolytic cells. Practical applications of electrolysis. Balancing Redox reaction equations: half reactions.
ELEMENTS OF TERMOCHEMISTRY AND THERMODYNAMICS
Energy and energy conservation. Internal energy and state functions. Energy and enthalpy. Enthalpies of chemical and physical transformations. Hess's Law. Standard creation enthalpies. Introduction to the concept of entropy.

Exercise:
SI base units. Concentration units of solutions and mixtures. Significant numbers. Rounding numbers. The law of conservation of mass. Chemical symbols. Balancing of chemical equations with particular emphasis on redox reaction equations. Computational solving of chemical problems in the field of: the content of elements in compounds and mixtures, concentrations of solutions, gas laws, pH of acid and base solutions, pH of buffer solutions, precipitation and dissolution reactions of chemical compounds, determination of solubility of chemical compounds in solution and product of solubility, conditions of quantitative precipitation of chemical compounds, calculation of electrode potentials and electromotive force of galvanic cells.

Laboratory

PART ONE – PARALLEL WORK ("EQUAL FRONT")

LABORATORY 1: ORGANIZATION OF WORK IN THE LABORATORY. LABORATORY TECHNIQUE.
Order regulations. Safety regulations. Basic laboratory equipment. Measuring liquids. Washing glassware. Heating of liquids. Dissolution and digestion. Evaporation and crystallization.
LABORATORY 2: EQUILIBRIUM IN ELECTROLYTE SOLUTIONS
Chemical equilibrium - the law of mass action. Strong and weak electrolytes. Concentration and activity. The role of the solvent. pH scale. Brönsted's theory. Solutions of acids and bases. Constant and degree of protolysis. Ostwald dilution law. Indicators of pH. pH measurement - indicator and reference electrodes.
LABORATORY 3: CONSTANT AND DEGREE OF DISSOCIATION OF WEAK ELECTROLYTES. BUFFER SOLUTIONS.
Equilibrium in solutions of weak electrolytes. Disturbing the state of equilibrium. The rule of defiance. Buffer solutions: composition, pH value, pH stabilization range, buffer capacity, use of buffer solutions.
LABORATORY 4: PRECIPITATION AND DISSOLUTION OF SEDIMENTS.
Saturated solutions. Solubility: definition and units. Factors affecting solubility. Effect of the common ion on the solubility of sediments. Application of the defiance rule in the process of precipitation and dissolution of sediments. Application of precipitation and dissolution of sediments in qualitative analysis.
LABORATORY 5: COORDINATION COMPOUNDS
Coordination binding. Conditions for the formation of complex ions. Structure of complex ions. Naming of complex compounds. Constants of stability and instability of complex ions. Aqua complexes.
LABORATORY 6: OXIDATION AND REDUCTION (REDOX) PROCESSES. ELEMENTS OF ELECTROCHEMISTRY
Oxidation and reduction reactions as a distinct type of chemical reactions. The concept of the degree of oxidation. Record of redox reaction equations (half reactions in ionic form). Half-cell potential - Nernst equation. Hydrogen electrode standard, its importance for the description of oxidation and reduction processes. Voltage series of metals. Oxidation-reduction series. Influence of temperature and pH on the course of oxidation and reduction reactions. Reaction rate. Passivation

PART TWO – INDIVIDUAL WORK
ANIONS:
CO32-, C2O42-, CH3COO-, NO3-, NO2-, SO42-, SO32-, S2O32-, S2-, SCN-, Cl-, Br-, I-, Cr2O72-, CrO42-, PO43-, MnO4-, OH-

1. Anions: preliminary tests. pH - measurement with universal paper. Reactions of anions (1-2 drops of the test solution) with AgNO3 (1-2 drops).
2. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with 1M H2SO4 (2-3 drops).
3. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KMnO4 (1 drop) against 1M H2SO4 (1-2 drops).
4. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KI (1 drop) against 1M H2SO4 (1 drop).
5. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with I2 (1 drop) against 1M H2SO4 (1 drop).
6. Anions: characteristic reactions.

CATIONS:
Cations group I: preliminary tests. Precipitation reaction of 1M HCl chlorides and their dissolution with excess conc. HCl. Take about 2-3 drops of cation solutions each.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NaOH. Take about 2 to 3 drops of cation solutions.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NH3. Take about 2 to 3 drops of cation solutions.
Group I cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Precipitation of sulphides with thioacetamide solution in an environment of 0.3M HCl. Dissolution of sulphides in 2M NaOH and in a mixture of 2M NaOH and 3% H2O2 (10:3) . Take about 2 to 3 drops of cation solutions.
Group II cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions. Group III cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with ammonium buffer 2M NH4Cl/2M NH3 (10:1) and attempt to precipitate sulphides with thioacetamide solution. Precipitation sulphides only in the absence of reaction with ammonium buffer. Take about 2 to 3 drops of cation solutions. Group III cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Carbonate precipitation reaction with solution (NH4)2CO3 and their dissolution in 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Chromate precipitation reaction with K2CrO4 solution in an inert and acidified medium of 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Sulphate precipitation reaction with 2M H2SO4 solution Take about 2 to 3 drops of cation solutions each.
Group IV and V cations: Characteristic reactions. Take about 2-3 drops of cation solutions
Self-detection of anions and cations (evaluated) in administered samples (aqueous solutions and powders).

Lectures:
BASIC CHEMICAL LAWS AND DEFINITIONS
Weight, length and temperature. Mol and Avogardo number, accuracy, precision, significant digits. Rounding numbers. Basic units and SI system. Calculation: unit conversion. The law of conservation of mass and the law of constant and multiple proportions. Balancing chemical equations. Chemical symbols. Stoichiometry. Concentration units.
ATOM – NUCLEUS
Dalton's atomic theory. Rutherford's experience. The structure of the atom: the nucleus and electrons. Elementary particles: electrons, neutrons, protons. Atomic number. Mass number and mole of the substance. Nuclear chemistry: transformation of elements. Natural radioactivity: alpha, beta and gamma. Nuclear stability: mass deficit and nucleus binding energy. The origin of chemical elements. Nuclear fusion and division of atomic nuclei. Atom bomb. Nuclear energy.
ATOM – ELECTRONIC STRUCTURE AND CHEMICAL BONDS
Bohr model. Wave-corpuscular dualism. Heisenberg's uncertainty principles. Compton phenomenon. Schroedinger equation. Wave functions and quantum numbers. Shapes of molecular orbitals. Emission spectra. Pauli's prohibition and Hundt’s rule. Energy of atomic orbitals in multi-electron atoms. The periodic table of elements and the electron configuration of elements. Ionic radius. The first and higher ionization potentials. Electron affinity. Octet rule. Ionic bond. Covalent molecules and bonds. The power of covalent bonds. Polarized covalent bonds. Electronegativity. Dot patterns. The theory of valence bonds. Hybridization and hybridized orbitals. Hydrogen molecule and other diatomic molecules. Bond order.
GASES AND LIQUIDS
Gaseous state and pressure of gases. Gas laws. Ideal gas equation. Partial pressure and Dalton's law. Kinetic-molecular theory of gases. Gas diffusion. Behavior of real gases. Critical temperature. Polarized bonds and the dipole moment. Intermolecular interactions. Properties of the liquid. Phase transitions: evaporation-condensation, melting-freezing, sublimation-resublimation. Gibbs phase rule. Phase diagram. Steam pressure and boiling point. Surface tension. Surfactants. Flotation. Solubility of gases in liquids.
SOLIDS
Types of solids: crystalline and non-crystalline. Ionic, covalent, metallic and molecular crystals. Allotropy. Study of the structure of solids: X-ray diffraction. Elementary cells. Liquid crystals. Metallic elements. Semiconductors and their applications.
SOLUTIONS
Electrolytes in aqueous solutions. Acid-base theories: Arrhenius and Broensted-Lowry. The strength of acids and bases. Proton hydration and oxonium ion. Self-ionization of water. pH scale. pH of acid and base solutions. Relationship between Ka and Kb. Acid-base properties of salts. Factors affecting the strength of acids and bases. The effect of a common ion. Buffer solutions. Indicators.
Solubility of ionic compounds. Factors affecting solubility. Precipitation of ionic compounds. Calculation of solubility using Ksp. Ion separation by selective precipitation. Qualitative analysis. Lowering the vapour pressure over solutions: Raoult's law. Boiling point increase and freezing point lowering. Osmosis and osmotic pressure. Dialysis.
CHEMICAL KINETICS
Reaction rate. Kinetic equations and the order of reactions. Half-time and first-order reactions. Radioactive decay. Second-order reactions. Reaction rate and temperature: The Arrhenius equation. Catalysis. Homogeneous and heterogeneous catalysts. Enzymatic catalysis. Equilibrium state. Equilibrium constants Kc and Kp. The use of constant equilibrium. Factors affecting the composition of the reaction mixture: Le Chatelier's rule. Synthesis of ammonia.
REDOX REACTIONS AND ELECTROCHEMISTRY
Galvanic cells. Short notation of galvanic cells. Electrode types. Electromotive force of cells. Standard potentials - Nernst equation. Electrochemical pH determination. Batteries. Fuel cells. Concentration cells. Corrosion. Electrolysis and electrolytic cells. Practical applications of electrolysis. Balancing Redox reaction equations: half reactions.
ELEMENTS OF TERMOCHEMISTRY AND THERMODYNAMICS
Energy and energy conservation. Internal energy and state functions. Energy and enthalpy. Enthalpies of chemical and physical transformations. Hess's Law. Standard creation enthalpies. Introduction to the concept of entropy.

Exercise:
SI base units. Concentration units of solutions and mixtures. Significant numbers. Rounding numbers. The law of conservation of mass. Chemical symbols. Balancing of chemical equations with particular emphasis on redox reaction equations. Computational solving of chemical problems in the field of: the content of elements in compounds and mixtures, concentrations of solutions, gas laws, pH of acid and base solutions, pH of buffer solutions, precipitation and dissolution reactions of chemical compounds, determination of solubility of chemical compounds in solution and product of solubility, conditions of quantitative precipitation of chemical compounds, calculation of electrode potentials and electromotive force of galvanic cells.

Laboratory

PART ONE – PARALLEL WORK ("EQUAL FRONT")

LABORATORY 1: ORGANIZATION OF WORK IN THE LABORATORY. LABORATORY TECHNIQUE.
Order regulations. Safety regulations. Basic laboratory equipment. Measuring liquids. Washing glassware. Heating of liquids. Dissolution and digestion. Evaporation and crystallization.
LABORATORY 2: EQUILIBRIUM IN ELECTROLYTE SOLUTIONS
Chemical equilibrium - the law of mass action. Strong and weak electrolytes. Concentration and activity. The role of the solvent. pH scale. Brönsted's theory. Solutions of acids and bases. Constant and degree of protolysis. Ostwald dilution law. Indicators of pH. pH measurement - indicator and reference electrodes.
LABORATORY 3: CONSTANT AND DEGREE OF DISSOCIATION OF WEAK ELECTROLYTES. BUFFER SOLUTIONS.
Equilibrium in solutions of weak electrolytes. Disturbing the state of equilibrium. The rule of defiance. Buffer solutions: composition, pH value, pH stabilization range, buffer capacity, use of buffer solutions.
LABORATORY 4: PRECIPITATION AND DISSOLUTION OF SEDIMENTS.
Saturated solutions. Solubility: definition and units. Factors affecting solubility. Effect of the common ion on the solubility of sediments. Application of the defiance rule in the process of precipitation and dissolution of sediments. Application of precipitation and dissolution of sediments in qualitative analysis.
LABORATORY 5: COORDINATION COMPOUNDS
Coordination binding. Conditions for the formation of complex ions. Structure of complex ions. Naming of complex compounds. Constants of stability and instability of complex ions. Aqua complexes.
LABORATORY 6: OXIDATION AND REDUCTION (REDOX) PROCESSES. ELEMENTS OF ELECTROCHEMISTRY
Oxidation and reduction reactions as a distinct type of chemical reactions. The concept of the degree of oxidation. Record of redox reaction equations (half reactions in ionic form). Half-cell potential - Nernst equation. Hydrogen electrode standard, its importance for the description of oxidation and reduction processes. Voltage series of metals. Oxidation-reduction series. Influence of temperature and pH on the course of oxidation and reduction reactions. Reaction rate. Passivation

PART TWO – INDIVIDUAL WORK
ANIONS:
CO32-, C2O42-, CH3COO-, NO3-, NO2-, SO42-, SO32-, S2O32-, S2-, SCN-, Cl-, Br-, I-, Cr2O72-, CrO42-, PO43-, MnO4-, OH-

1. Anions: preliminary tests. pH - measurement with universal paper. Reactions of anions (1-2 drops of the test solution) with AgNO3 (1-2 drops).
2. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with 1M H2SO4 (2-3 drops).
3. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KMnO4 (1 drop) against 1M H2SO4 (1-2 drops).
4. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KI (1 drop) against 1M H2SO4 (1 drop).
5. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with I2 (1 drop) against 1M H2SO4 (1 drop).
6. Anions: characteristic reactions.

CATIONS:
Cations group I: preliminary tests. Precipitation reaction of 1M HCl chlorides and their dissolution with excess conc. HCl. Take about 2-3 drops of cation solutions each.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NaOH. Take about 2 to 3 drops of cation solutions.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NH3. Take about 2 to 3 drops of cation solutions.
Group I cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Precipitation of sulphides with thioacetamide solution in an environment of 0.3M HCl. Dissolution of sulphides in 2M NaOH and in a mixture of 2M NaOH and 3% H2O2 (10:3) . Take about 2 to 3 drops of cation solutions.
Group II cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions. Group III cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with ammonium buffer 2M NH4Cl/2M NH3 (10:1) and attempt to precipitate sulphides with thioacetamide solution. Precipitation sulphides only in the absence of reaction with ammonium buffer. Take about 2 to 3 drops of cation solutions. Group III cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Carbonate precipitation reaction with solution (NH4)2CO3 and their dissolution in 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Chromate precipitation reaction with K2CrO4 solution in an inert and acidified medium of 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Sulphate precipitation reaction with 2M H2SO4 solution Take about 2 to 3 drops of cation solutions each.
Group IV and V cations: Characteristic reactions. Take about 2-3 drops of cation solutions
Self-detection of anions and cations (evaluated) in administered samples (aqueous solutions and powders).

Lectures:
BASIC CHEMICAL LAWS AND DEFINITIONS
Weight, length and temperature. Mol and Avogardo number, accuracy, precision, significant digits. Rounding numbers. Basic units and SI system. Calculation: unit conversion. The law of conservation of mass and the law of constant and multiple proportions. Balancing chemical equations. Chemical symbols. Stoichiometry. Concentration units.
ATOM – NUCLEUS
Dalton's atomic theory. Rutherford's experience. The structure of the atom: the nucleus and electrons. Elementary particles: electrons, neutrons, protons. Atomic number. Mass number and mole of the substance. Nuclear chemistry: transformation of elements. Natural radioactivity: alpha, beta and gamma. Nuclear stability: mass deficit and nucleus binding energy. The origin of chemical elements. Nuclear fusion and division of atomic nuclei. Atom bomb. Nuclear energy.
ATOM – ELECTRONIC STRUCTURE AND CHEMICAL BONDS
Bohr model. Wave-corpuscular dualism. Heisenberg's uncertainty principles. Compton phenomenon. Schroedinger equation. Wave functions and quantum numbers. Shapes of molecular orbitals. Emission spectra. Pauli's prohibition and Hundt’s rule. Energy of atomic orbitals in multi-electron atoms. The periodic table of elements and the electron configuration of elements. Ionic radius. The first and higher ionization potentials. Electron affinity. Octet rule. Ionic bond. Covalent molecules and bonds. The power of covalent bonds. Polarized covalent bonds. Electronegativity. Dot patterns. The theory of valence bonds. Hybridization and hybridized orbitals. Hydrogen molecule and other diatomic molecules. Bond order.
GASES AND LIQUIDS
Gaseous state and pressure of gases. Gas laws. Ideal gas equation. Partial pressure and Dalton's law. Kinetic-molecular theory of gases. Gas diffusion. Behavior of real gases. Critical temperature. Polarized bonds and the dipole moment. Intermolecular interactions. Properties of the liquid. Phase transitions: evaporation-condensation, melting-freezing, sublimation-resublimation. Gibbs phase rule. Phase diagram. Steam pressure and boiling point. Surface tension. Surfactants. Flotation. Solubility of gases in liquids.
SOLIDS
Types of solids: crystalline and non-crystalline. Ionic, covalent, metallic and molecular crystals. Allotropy. Study of the structure of solids: X-ray diffraction. Elementary cells. Liquid crystals. Metallic elements. Semiconductors and their applications.
SOLUTIONS
Electrolytes in aqueous solutions. Acid-base theories: Arrhenius and Broensted-Lowry. The strength of acids and bases. Proton hydration and oxonium ion. Self-ionization of water. pH scale. pH of acid and base solutions. Relationship between Ka and Kb. Acid-base properties of salts. Factors affecting the strength of acids and bases. The effect of a common ion. Buffer solutions. Indicators.
Solubility of ionic compounds. Factors affecting solubility. Precipitation of ionic compounds. Calculation of solubility using Ksp. Ion separation by selective precipitation. Qualitative analysis. Lowering the vapour pressure over solutions: Raoult's law. Boiling point increase and freezing point lowering. Osmosis and osmotic pressure. Dialysis.
CHEMICAL KINETICS
Reaction rate. Kinetic equations and the order of reactions. Half-time and first-order reactions. Radioactive decay. Second-order reactions. Reaction rate and temperature: The Arrhenius equation. Catalysis. Homogeneous and heterogeneous catalysts. Enzymatic catalysis. Equilibrium state. Equilibrium constants Kc and Kp. The use of constant equilibrium. Factors affecting the composition of the reaction mixture: Le Chatelier's rule. Synthesis of ammonia.
REDOX REACTIONS AND ELECTROCHEMISTRY
Galvanic cells. Short notation of galvanic cells. Electrode types. Electromotive force of cells. Standard potentials - Nernst equation. Electrochemical pH determination. Batteries. Fuel cells. Concentration cells. Corrosion. Electrolysis and electrolytic cells. Practical applications of electrolysis. Balancing Redox reaction equations: half reactions.
ELEMENTS OF TERMOCHEMISTRY AND THERMODYNAMICS
Energy and energy conservation. Internal energy and state functions. Energy and enthalpy. Enthalpies of chemical and physical transformations. Hess's Law. Standard creation enthalpies. Introduction to the concept of entropy.

Exercise:
SI base units. Concentration units of solutions and mixtures. Significant numbers. Rounding numbers. The law of conservation of mass. Chemical symbols. Balancing of chemical equations with particular emphasis on redox reaction equations. Computational solving of chemical problems in the field of: the content of elements in compounds and mixtures, concentrations of solutions, gas laws, pH of acid and base solutions, pH of buffer solutions, precipitation and dissolution reactions of chemical compounds, determination of solubility of chemical compounds in solution and product of solubility, conditions of quantitative precipitation of chemical compounds, calculation of electrode potentials and electromotive force of galvanic cells.

Laboratory

PART ONE – PARALLEL WORK ("EQUAL FRONT")

LABORATORY 1: ORGANIZATION OF WORK IN THE LABORATORY. LABORATORY TECHNIQUE.
Order regulations. Safety regulations. Basic laboratory equipment. Measuring liquids. Washing glassware. Heating of liquids. Dissolution and digestion. Evaporation and crystallization.
LABORATORY 2: EQUILIBRIUM IN ELECTROLYTE SOLUTIONS
Chemical equilibrium - the law of mass action. Strong and weak electrolytes. Concentration and activity. The role of the solvent. pH scale. Brönsted's theory. Solutions of acids and bases. Constant and degree of protolysis. Ostwald dilution law. Indicators of pH. pH measurement - indicator and reference electrodes.
LABORATORY 3: CONSTANT AND DEGREE OF DISSOCIATION OF WEAK ELECTROLYTES. BUFFER SOLUTIONS.
Equilibrium in solutions of weak electrolytes. Disturbing the state of equilibrium. The rule of defiance. Buffer solutions: composition, pH value, pH stabilization range, buffer capacity, use of buffer solutions.
LABORATORY 4: PRECIPITATION AND DISSOLUTION OF SEDIMENTS.
Saturated solutions. Solubility: definition and units. Factors affecting solubility. Effect of the common ion on the solubility of sediments. Application of the defiance rule in the process of precipitation and dissolution of sediments. Application of precipitation and dissolution of sediments in qualitative analysis.
LABORATORY 5: COORDINATION COMPOUNDS
Coordination binding. Conditions for the formation of complex ions. Structure of complex ions. Naming of complex compounds. Constants of stability and instability of complex ions. Aqua complexes.
LABORATORY 6: OXIDATION AND REDUCTION (REDOX) PROCESSES. ELEMENTS OF ELECTROCHEMISTRY
Oxidation and reduction reactions as a distinct type of chemical reactions. The concept of the degree of oxidation. Record of redox reaction equations (half reactions in ionic form). Half-cell potential - Nernst equation. Hydrogen electrode standard, its importance for the description of oxidation and reduction processes. Voltage series of metals. Oxidation-reduction series. Influence of temperature and pH on the course of oxidation and reduction reactions. Reaction rate. Passivation

PART TWO – INDIVIDUAL WORK
ANIONS:
CO32-, C2O42-, CH3COO-, NO3-, NO2-, SO42-, SO32-, S2O32-, S2-, SCN-, Cl-, Br-, I-, Cr2O72-, CrO42-, PO43-, MnO4-, OH-

1. Anions: preliminary tests. pH - measurement with universal paper. Reactions of anions (1-2 drops of the test solution) with AgNO3 (1-2 drops).
2. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with 1M H2SO4 (2-3 drops).
3. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KMnO4 (1 drop) against 1M H2SO4 (1-2 drops).
4. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with KI (1 drop) against 1M H2SO4 (1 drop).
5. Anions: preliminary tests. Reactions of anions (approx. 0.5 cm3 of solution) with I2 (1 drop) against 1M H2SO4 (1 drop).
6. Anions: characteristic reactions.

CATIONS:
Cations group I: preliminary tests. Precipitation reaction of 1M HCl chlorides and their dissolution with excess conc. HCl. Take about 2-3 drops of cation solutions each.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NaOH. Take about 2 to 3 drops of cation solutions.
Cations group I: preliminary tests. The reaction of precipitation of 1M HCl chlorides and their dissolution with an excess of 2M NH3. Take about 2 to 3 drops of cation solutions.
Group I cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions. Take about 2 to 3 drops of cation solutions.
Group II cations: preliminary tests. Precipitation of sulphides with thioacetamide solution in an environment of 0.3M HCl. Dissolution of sulphides in 2M NaOH and in a mixture of 2M NaOH and 3% H2O2 (10:3) . Take about 2 to 3 drops of cation solutions.
Group II cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with 2M NaOH and excess reagent. Take about 2 to 3 drops of cation solutions. Group III cations: preliminary tests. Reaction with 2M NH3 and excess reagent. Take about 2 to 3 drops of cation solutions.
Group III cations: preliminary tests. Reaction with ammonium buffer 2M NH4Cl/2M NH3 (10:1) and attempt to precipitate sulphides with thioacetamide solution. Precipitation sulphides only in the absence of reaction with ammonium buffer. Take about 2 to 3 drops of cation solutions. Group III cations: Characteristic reactions. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Carbonate precipitation reaction with solution (NH4)2CO3 and their dissolution in 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Chromate precipitation reaction with K2CrO4 solution in an inert and acidified medium of 2M CH3COOH. Take about 2 to 3 drops of cation solutions.
Group IV cations: preliminary tests. Sulphate precipitation reaction with 2M H2SO4 solution Take about 2 to 3 drops of cation solutions each.
Group IV and V cations: Characteristic reactions. Take about 2-3 drops of cation solutions
Self-detection of anions and cations (evaluated) in administered samples (aqueous solutions and powders).

1. L. Jones, P. Atkins, L. Leroy, Chemistry, Molecules, Matter, and Change, 2000, W.H. Freeman & Company, New York
2. L. Pajdowski, Chemia ogólna, 2002, PWN, Warszawa
3. A. Bielański, Podstawy chemii nieorganicznej tom 1 (wyd. 6), 2012, PWN, Warszawa
4. F. A. Cotton, G. Wilkinson, P. L. Gaus, Chemia nieorganiczna Podstawy, 2002, PWN, Warszawa
5. M. J. Sienko, R. A. Plane, Chemia, podstawy i zastosowania, (wyd. 6), 2002, WNT, Warszawa
6. J. D. Lee, Zwięzła chemia nieorganiczna, 1998, PWN, Warszawa
7. L. Pauling, P. Pauling, Chemia, 1997, PWN, Warszawa
8. A. Hulanicki, Reakcje kwasów i zasad w chemii analitycznej (wyd.4), 2021, PWN, Warszawa
9. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część I, 2008, UMK, Toruń
10. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część II, 2008, UMK, Toruń
11. Z.S. Szmal, T. Lipiec, Chemia analityczna z elementami analizy instrumentalnej, 1997, PZWL, Warszawa
12. A. Reizer , Ćwiczenia z podstaw chemii i analizy jakościowej (wyd.2), 2000, Wydawnictwo UJ, Kraków
13. Z. Hubicki, Ćwiczenia laboratoryjne z klasycznej analizy jakościowej nieorganicznej, 2017, Wydawnictwo UMCS, Lublin
14. L. Chmurzyński, Ćwiczenia laboratoryjne z chemii ogólnej I. Część teoretyczna, Wydawnictwo UG, Gdańsk 2011
15. E. Schweda, Chemia nieorganiczna Tom I Wprowadzenie i analiza jakościowa, 2014, Medpharm , Wrocław
16. R. Kocjan, Chemia analityczna tom 1 (wyd.2), PZWL, Warszawa 2014
17. J. Minczewski, Z. Marczenko, Chemia analityczna t. I, Podstawy teoretyczne i analiza jakościowa, 2017, Warszawa
18. J. Minczewski, Z. Marczenko, Chemia analityczna t. II, Analiza ilościowa, 2017, PWN, Warszawa
19. A. Śliwa, Obliczenia chemiczne, zbiór zadań z chemii ogólnej i analitycznej nieorganicznej. 1987, PWN, Warszawa
20. H. Kowalczyk-Dembińska, Ćwiczenia rachunkowe z podstaw chemii, 2001, Wydawnictwo UMK, Toruń
21. Z. Warnke, Obliczenia z chemii ogólnej, Wydawnictwo UG, 2015, Gdańsk
22. A. Persona, J. Reszko-Zygmunt, T. Gęca, Zbiór zadań z chemii ogólnej i analitycznej z pełnymi rozwiązaniami, 2020, Wydawnictwo Medyk, Warszawa
23. J. Kalembkiewicz, B. Papciak, Chemia ogólna i nieorganiczna. Podstawy chemii. Roztwory i procesy w roztworach. Obliczenia chemiczne i problemy, 2021, Wydawnictwo Politechniki Rzeszowskiej, Rzeszów
24. Z. Galus, Ćwiczenia rachunkowe z chemii analitycznej, 1993, PWN, Warszawa

1. L. Jones, P. Atkins, L. Leroy, Chemistry, Molecules, Matter, and Change, 2000, W.H. Freeman & Company, New York
2. L. Pajdowski, Chemia ogólna, 2002, PWN, Warszawa
3. A. Bielański, Podstawy chemii nieorganicznej tom 1 (wyd. 6), 2012, PWN, Warszawa
4. F. A. Cotton, G. Wilkinson, P. L. Gaus, Chemia nieorganiczna Podstawy, 2002, PWN, Warszawa
5. M. J. Sienko, R. A. Plane, Chemia, podstawy i zastosowania, (wyd. 6), 2002, WNT, Warszawa
6. J. D. Lee, Zwięzła chemia nieorganiczna, 1998, PWN, Warszawa
7. L. Pauling, P. Pauling, Chemia, 1997, PWN, Warszawa
8. A. Hulanicki, Reakcje kwasów i zasad w chemii analitycznej (wyd.4), 2021, PWN, Warszawa
9. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część I, 2008, UMK, Toruń
10. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część II, 2008, UMK, Toruń
11. Z.S. Szmal, T. Lipiec, Chemia analityczna z elementami analizy instrumentalnej, 1997, PZWL, Warszawa
12. A. Reizer , Ćwiczenia z podstaw chemii i analizy jakościowej (wyd.2), 2000, Wydawnictwo UJ, Kraków
13. Z. Hubicki, Ćwiczenia laboratoryjne z klasycznej analizy jakościowej nieorganicznej, 2017, Wydawnictwo UMCS, Lublin
14. L. Chmurzyński, Ćwiczenia laboratoryjne z chemii ogólnej I. Część teoretyczna, Wydawnictwo UG, Gdańsk 2011
15. E. Schweda, Chemia nieorganiczna Tom I Wprowadzenie i analiza jakościowa, 2014, Medpharm , Wrocław
16. R. Kocjan, Chemia analityczna tom 1 (wyd.2), PZWL, Warszawa 2014
17. J. Minczewski, Z. Marczenko, Chemia analityczna t. I, Podstawy teoretyczne i analiza jakościowa, 2017, Warszawa
18. J. Minczewski, Z. Marczenko, Chemia analityczna t. II, Analiza ilościowa, 2017, PWN, Warszawa
19. A. Śliwa, Obliczenia chemiczne, zbiór zadań z chemii ogólnej i analitycznej nieorganicznej. 1987, PWN, Warszawa
20. H. Kowalczyk-Dembińska, Ćwiczenia rachunkowe z podstaw chemii, 2001, Wydawnictwo UMK, Toruń
21. Z. Warnke, Obliczenia z chemii ogólnej, Wydawnictwo UG, 2015, Gdańsk
22. A. Persona, J. Reszko-Zygmunt, T. Gęca, Zbiór zadań z chemii ogólnej i analitycznej z pełnymi rozwiązaniami, 2020, Wydawnictwo Medyk, Warszawa
23. J. Kalembkiewicz, B. Papciak, Chemia ogólna i nieorganiczna. Podstawy chemii. Roztwory i procesy w roztworach. Obliczenia chemiczne i problemy, 2021, Wydawnictwo Politechniki Rzeszowskiej, Rzeszów
24. Z. Galus, Ćwiczenia rachunkowe z chemii analitycznej, 1993, PWN, Warszawa

1. L. Jones, P. Atkins, L. Leroy, Chemistry, Molecules, Matter, and Change, 2000, W.H. Freeman & Company, New York
2. L. Pajdowski, Chemia ogólna, 2002, PWN, Warszawa
3. A. Bielański, Podstawy chemii nieorganicznej tom 1 (wyd. 6), 2012, PWN, Warszawa
4. F. A. Cotton, G. Wilkinson, P. L. Gaus, Chemia nieorganiczna Podstawy, 2002, PWN, Warszawa
5. M. J. Sienko, R. A. Plane, Chemia, podstawy i zastosowania, (wyd. 6), 2002, WNT, Warszawa
6. J. D. Lee, Zwięzła chemia nieorganiczna, 1998, PWN, Warszawa
7. L. Pauling, P. Pauling, Chemia, 1997, PWN, Warszawa
8. A. Hulanicki, Reakcje kwasów i zasad w chemii analitycznej (wyd.4), 2021, PWN, Warszawa
9. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część I, 2008, UMK, Toruń
10. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część II, 2008, UMK, Toruń
11. Z.S. Szmal, T. Lipiec, Chemia analityczna z elementami analizy instrumentalnej, 1997, PZWL, Warszawa
12. A. Reizer , Ćwiczenia z podstaw chemii i analizy jakościowej (wyd.2), 2000, Wydawnictwo UJ, Kraków
13. Z. Hubicki, Ćwiczenia laboratoryjne z klasycznej analizy jakościowej nieorganicznej, 2017, Wydawnictwo UMCS, Lublin
14. L. Chmurzyński, Ćwiczenia laboratoryjne z chemii ogólnej I. Część teoretyczna, Wydawnictwo UG, Gdańsk 2011
15. E. Schweda, Chemia nieorganiczna Tom I Wprowadzenie i analiza jakościowa, 2014, Medpharm , Wrocław
16. R. Kocjan, Chemia analityczna tom 1 (wyd.2), PZWL, Warszawa 2014
17. J. Minczewski, Z. Marczenko, Chemia analityczna t. I, Podstawy teoretyczne i analiza jakościowa, 2017, Warszawa
18. J. Minczewski, Z. Marczenko, Chemia analityczna t. II, Analiza ilościowa, 2017, PWN, Warszawa
19. A. Śliwa, Obliczenia chemiczne, zbiór zadań z chemii ogólnej i analitycznej nieorganicznej. 1987, PWN, Warszawa
20. H. Kowalczyk-Dembińska, Ćwiczenia rachunkowe z podstaw chemii, 2001, Wydawnictwo UMK, Toruń
21. Z. Warnke, Obliczenia z chemii ogólnej, Wydawnictwo UG, 2015, Gdańsk
22. A. Persona, J. Reszko-Zygmunt, T. Gęca, Zbiór zadań z chemii ogólnej i analitycznej z pełnymi rozwiązaniami, 2020, Wydawnictwo Medyk, Warszawa
23. J. Kalembkiewicz, B. Papciak, Chemia ogólna i nieorganiczna. Podstawy chemii. Roztwory i procesy w roztworach. Obliczenia chemiczne i problemy, 2021, Wydawnictwo Politechniki Rzeszowskiej, Rzeszów
24. Z. Galus, Ćwiczenia rachunkowe z chemii analitycznej, 1993, PWN, Warszawa

1. L. Jones, P. Atkins, L. Leroy, Chemistry, Molecules, Matter, and Change, 2000, W.H. Freeman & Company, New York
2. L. Pajdowski, Chemia ogólna, 2002, PWN, Warszawa
3. A. Bielański, Podstawy chemii nieorganicznej tom 1 (wyd. 6), 2012, PWN, Warszawa
4. F. A. Cotton, G. Wilkinson, P. L. Gaus, Chemia nieorganiczna Podstawy, 2002, PWN, Warszawa
5. M. J. Sienko, R. A. Plane, Chemia, podstawy i zastosowania, (wyd. 6), 2002, WNT, Warszawa
6. J. D. Lee, Zwięzła chemia nieorganiczna, 1998, PWN, Warszawa
7. L. Pauling, P. Pauling, Chemia, 1997, PWN, Warszawa
8. A. Hulanicki, Reakcje kwasów i zasad w chemii analitycznej (wyd.4), 2021, PWN, Warszawa
9. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część I, 2008, UMK, Toruń
10. H. Kowalczyk-Dembińska, J. Łukaszewicz, Chemia ogólna i jakościowa analiza chemiczna. Ćwiczenia laboratoryjne – część II, 2008, UMK, Toruń
11. Z.S. Szmal, T. Lipiec, Chemia analityczna z elementami analizy instrumentalnej, 1997, PZWL, Warszawa
12. A. Reizer , Ćwiczenia z podstaw chemii i analizy jakościowej (wyd.2), 2000, Wydawnictwo UJ, Kraków
13. Z. Hubicki, Ćwiczenia laboratoryjne z klasycznej analizy jakościowej nieorganicznej, 2017, Wydawnictwo UMCS, Lublin
14. L. Chmurzyński, Ćwiczenia laboratoryjne z chemii ogólnej I. Część teoretyczna, Wydawnictwo UG, Gdańsk 2011
15. E. Schweda, Chemia nieorganiczna Tom I Wprowadzenie i analiza jakościowa, 2014, Medpharm , Wrocław
16. R. Kocjan, Chemia analityczna tom 1 (wyd.2), PZWL, Warszawa 2014
17. J. Minczewski, Z. Marczenko, Chemia analityczna t. I, Podstawy teoretyczne i analiza jakościowa, 2017, Warszawa
18. J. Minczewski, Z. Marczenko, Chemia analityczna t. II, Analiza ilościowa, 2017, PWN, Warszawa
19. A. Śliwa, Obliczenia chemiczne, zbiór zadań z chemii ogólnej i analitycznej nieorganicznej. 1987, PWN, Warszawa
20. H. Kowalczyk-Dembińska, Ćwiczenia rachunkowe z podstaw chemii, 2001, Wydawnictwo UMK, Toruń
21. Z. Warnke, Obliczenia z chemii ogólnej, Wydawnictwo UG, 2015, Gdańsk
22. A. Persona, J. Reszko-Zygmunt, T. Gęca, Zbiór zadań z chemii ogólnej i analitycznej z pełnymi rozwiązaniami, 2020, Wydawnictwo Medyk, Warszawa
23. J. Kalembkiewicz, B. Papciak, Chemia ogólna i nieorganiczna. Podstawy chemii. Roztwory i procesy w roztworach. Obliczenia chemiczne i problemy, 2021, Wydawnictwo Politechniki Rzeszowskiej, Rzeszów
24. Z. Galus, Ćwiczenia rachunkowe z chemii analitycznej, 1993, PWN, Warszawa