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AP Chemistry Comprehensive Syllabus

AP Chemistry Comprehensive Syllabus

Unit 1: Atomic Structure and Properties

Subtopic Number Subtopic Name Outline
1.1 Moles and molar mass The concept of a mole:

  • A mole is a unit of measurement used to express the amount of a substance.
  • Avogadro’s number: Avogadro’s number is the number of atoms or molecules in one mole of a substance.
  • Relationship between mole and number of particles: The number of particles in a substance can be calculated by multiplying the number of moles by Avogadro’s number.

Molar mass:

  • Molar mass is the mass of one mole of a substance.
  • Relationship between mass, moles, and molar mass: Mass can be calculated by multiplying the number of moles by the molar mass, while the number of moles can be calculated by dividing the mass by the molar mass.
  • Calculation of molar mass: Molar mass can be calculated by adding up the atomic masses of all the atoms in a molecule.
  • Units of molar mass: Molar mass is typically expressed in grams per mole (g/mol).

Stoichiometry:

  • Stoichiometry is the calculation of the quantities of reactants and products involved in a chemical reaction.
  • Stoichiometric calculations involving moles and molar mass: Stoichiometric calculations often involve using moles and molar mass to determine the quantities of reactants and products involved in a reaction.
  • Balanced chemical equations: Balanced chemical equations show the relationship between reactants and products in a chemical reaction.
  • Limiting reactants and percent yield: Limiting reactants are the reactants that are completely consumed in a reaction, while percent yield is the ratio of the actual yield to the theoretical yield.

Empirical and molecular formulas:

  • Definition of empirical and molecular formulas: Empirical formulas show the simplest ratio of atoms in a molecule, while molecular formulas show the actual number of atoms in a molecule.
  • Calculation of empirical and molecular formulas from experimental data: Empirical and molecular formulas can be determined from experimental data such as mass and elemental composition.
  • Relationship between empirical and molecular formulas: The molecular formula is a multiple of the empirical formula.

Percent composition:

  • Percent composition is the percentage of each element in a compound.
  • Calculation of percent composition from experimental data: Percent composition can be calculated from experimental data such as mass and elemental composition.
  • Relationship between percent composition and empirical formula: The percent composition can be used to determine the empirical formula.

Gas stoichiometry:

  • Calculation of amount of gas using the ideal gas law: The ideal gas law can be used to calculate the amount of gas involved in a reaction.
  • Stoichiometric calculations involving glasses: Stoichiometric calculations involving gasses often require the use of the ideal gas law.
  • Relationship between volume, pressure, temperature, and amount of gas: The volume, pressure, temperature, and amount of gas are related through the ideal gas law.

Solution stoichiometry:

  • Calculation of concentration of a solution: Concentration of a solution can be expressed in terms of molarity, which is the number of moles of solute per liter of solution.
  • Stoichiometric calculations involving solutions: Stoichiometric calculations involving solutions often require the use of concentration and volume.
  • Dilution calculations: Dilution calculations involve the addition of solvent to a solution to decrease its concentration.

 Limitations of the mole concept

1.2 Mass spectroscopy of elements
  • Mass Spectrometry is a technique used to identify the chemical composition of a sample based on the mass-to-charge ratio of its ions. It involves ionizing a sample, separating the resulting ions based on their mass-to-charge ratio, and detecting and measuring the abundance of each ion.
  • Isotopes: Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei, and thus different mass numbers. Mass Spectrometry can be used to distinguish between isotopes based on their mass-to-charge ratios, which allows for the determination of the isotopic composition of a sample.
  • Molecular Ion: A molecular ion is an ion formed by the removal of one or more electrons from a molecule, resulting in a positive charge. Mass Spectrometry can be used to identify molecular ions based on their mass-to-charge ratio, which can provide information about the molecular formula and structure of a compound.
  • Fragmentation: Fragmentation is the process by which a molecular ion breaks down into smaller fragments, often due to the loss of a neutral molecule such as water or carbon dioxide. The resulting fragment ions can be used to identify the functional groups and connectivity of atoms within the original molecule.
  • Mass Spectral Interpretation: Mass Spectral Interpretation is the process of using Mass Spectrometry data to identify and characterize chemical compounds. This involves analyzing the mass-to-charge ratio and abundance of ions in the mass spectrum, as well as considering the possible fragmentation pathways that could lead to the observed ions.
1.3 Elemental composition of pure substances
  • Atomic Structure: Atomic structure refers to the basic structure of an atom, which is made up of a nucleus containing protons and neutrons, surrounded by electrons in various energy levels or orbitals.
  • Isotopes: Isotopes are atoms of the same element that have different numbers of neutrons in their nuclei, resulting in different atomic masses.
  • Atomic Mass: Atomic mass is the mass of an atom of an element, expressed in atomic mass units (amu), which is the weighted average of the masses of all the naturally occurring isotopes of that element.
  • Mass Percent Composition: Mass percent composition is the mass of each element in a compound divided by the total mass of the compound, expressed as a percentage.
  • Empirical and Molecular Formulas: The empirical formula of a compound gives the simplest whole-number ratio of the atoms in the compound, while the molecular formula gives the actual number of atoms of each element in a molecule of the compound.
  • Stoichiometry: Stoichiometry involves using the mole concept and balanced chemical equations to determine the quantities of reactants and products in a chemical reaction.
  • Limiting and Excess Reactants: The limiting reactant is the reactant that is completely consumed in a reaction, while the excess reactant is the one that is not completely consumed.
  • Percent Yield: Percent yield is the actual yield of a reaction, expressed as a percentage of the theoretical yield, which is the maximum amount of product that could be obtained based on the amount of limiting reactant present.
1.4 Composition of mixtures
  • Mixture: A combination of two or more substances that are not chemically bonded to each other.
  • Homogeneous mixture: A mixture in which the composition is uniform throughout. Also called a solution.
  • Heterogeneous mixture: A mixture in which the composition is not uniform throughout.
  • Solute: The substance that is dissolved in a solution.
  • Solvent: The substance that dissolves the solute in a solution.
  • Concentration: A measure of the amount of solute dissolved in a given amount of solvent or solution.
  • Molarity: The number of moles of solute per liter of solution.
  • Molality: The number of moles of solute per kilogram of solvent.
  • Mass Percent: The mass of the solute divided by the mass of the solution, multiplied by 100%.
  • Parts per million (ppm): The number of parts of solute per million parts of solution.
  • Parts per billion (ppb): The number of parts of solute per billion parts of solution.
  • Colligative Properties: Properties of solutions that depend on the concentration of solute particles in solution, but not on their chemical identity. Examples include boiling point elevation and freezing point depression.
1.5 Atomic structure and electron configuration
  • Atomic Structure: Refers to the composition of an atom, which includes the protons, neutrons, and electrons. Protons are positively charged particles found in the nucleus of an atom, while neutrons are neutral particles also found in the nucleus. Electrons are negatively charged particles that orbit the nucleus.
  • Electron Configuration: Refers to the arrangement of electrons in an atom. The electrons occupy various energy levels, also known as electron shells, and each shell can only hold a specific number of electrons. The electron configuration of an atom can be represented using various notations, such as the orbital notation or the electron configuration notation.
  • Orbital Notation: A representation of the electron configuration of an atom that uses arrows pointing up or down to represent electrons in different orbitals. Each orbital can hold a maximum of two electrons, with opposite spins.
  • Electron Configuration Notation: A shorthand representation of the electron configuration of an atom that uses numbers and letters to represent the different electron shells and subshells. For example, the electron configuration of carbon can be written as 1s2 2s2 2p2, which indicates that carbon has two electrons in its first energy level, two electrons in its second energy level (2s subshell), and two electrons in its third energy level (2p subshell).
  • Valence Electrons: Refers to the electrons in the outermost energy level of an atom, also known as the valence shell. These electrons are involved in chemical bonding and determine the reactivity of an element.
  • Ionization Energy: Refers to the amount of energy required to remove an electron from an atom. The ionization energy of an element increases as you move from left to right across a period and decreases as you move from top to bottom down a group on the periodic table.
  • Electronegativity: Refers to the ability of an atom to attract electrons towards itself in a chemical bond. The electronegativity of an element also increases as you move from left to right across a period and decreases as you move from top to bottom down a group on the periodic table.
1.6 Photoelectron spectroscopy
  • Photoelectron Spectroscopy (PES): PES is a technique used to determine the electronic structure of atoms and molecules. In PES, a photon of high energy is directed at a sample, which then causes an electron to be ejected from an inner shell of the atom. The energy required to remove the electron is measured and used to determine the binding energy of the electron to the atom.
  • Binding Energy: Binding energy is the energy required to remove an electron from an atom or molecule. It is also known as ionization energy.
  • Energy Levels and Subshells: Energy levels are the regions of space around an atomic nucleus where electrons are most likely to be found. Subshells are the regions within energy levels where electrons are most likely to be found. The subshells are labeled s, p, d, and f.
  • Orbital Diagrams: Orbital diagrams are diagrams that show the distribution of electrons in the energy levels and subshells of an atom or molecule.
  • Molecular Orbitals: Molecular orbitals are the result of combining atomic orbitals of atoms in a molecule. They describe the distribution of electrons in a molecule and are used to determine its properties.
1.7 Periodic trends
  • Periodic Table: A table that organizes elements based on their atomic number, electron configuration, and chemical properties. Elements in the same group share similar chemical properties.
  • Atomic Radius: Atomic radius decreases from left to right across a period and increases from top to bottom within a group.
  • Electronegativity:. Electronegativity increases from left to right across a period and decreases from top to bottom within a group.
  • Ionization Energy: Ionization energy increases from left to right across a period and decreases from top to bottom within a group.
  • Electron Affinity: The energy change that occurs when an electron is added to a neutral atom or ion in the gas phase. Electron affinity increases from left to right across a period and decreases from top to bottom within a group.
  • Metallic Character: A measure of how easily an atom can lose an electron. Metallic character increases from right to left across a period and from bottom to top within a group.
  • Trends in Group Properties: Elements in the same group exhibit similar chemical properties due to their similar valence electron configuration. Groups 1, 2, 13, and 18 have distinctive trends in properties such as reactivity, ionization energy, and electronegativity.
  • Exceptions to Periodic Trends: Some elements exhibit deviations from periodic trends due to their unique electronic configurations, such as the noble gasses and the transition metals.
1.8 Valence electrons and ionic compounds
  • Valence Electrons: These are the electrons present in the outermost energy level of an atom that are involved in chemical reactions.
  • Ionic Compounds: These are compounds formed by the transfer of electrons between atoms. They consist of positively charged cations and negatively charged anions, held together by electrostatic attractions.
  • Ionic Bonds: These are the bonds formed by the electrostatic attraction between oppositely charged ions in an ionic compound.
  • Lattice Energy: This is the energy required to completely separate one mole of a solid ionic compound into its gaseous ions.
  • Properties of Ionic Compounds: Ionic compounds have high melting and boiling points, are usually crystalline solids at room temperature, and are often soluble in polar solvents.
  • Lewis Dot Structures: These are diagrams that show the bonding between atoms in a molecule or ion and the lone pairs of electrons that may exist in the molecule.
  • Octet Rule: This is the tendency of atoms to gain, lose, or share electrons in order to have a full outer shell of eight electrons.
  • Electronegativity: This is a measure of the ability of an atom to attract electrons towards itself in a covalent bond.
  • Ionic Radius: This is the distance between the nucleus and the outermost electron in an ion.
  • Cation: This is a positively charged ion that is formed by the loss of one or more electrons from an atom.
  • Anion: This is a negatively charged ion that is formed by the gain of one or more electrons by an atom.

Unit 2: Molecular and Ionic Compound Structure and Properties

Subtopic Number Subtopic Name Outline
2.1 Types of chemical bonds
  • Ionic Bonds: Ionic bonds are formed between atoms with very different electronegativity values. An ionic bond occurs when one atom gives up one or more electrons to another atom. This results in the formation of ions, which are atoms with an electrical charge.
  • Covalent Bonds: Covalent bonds are formed when two atoms share electrons. The atoms share electrons in order to achieve a stable electron configuration. This type of bond is typically formed between atoms that have similar electronegativity values.
  • Metallic Bonds: Metallic bonds are formed between metal atoms. In a metallic bond, the valence electrons of metal atoms are shared between all the atoms in the metal. This results in a sea of electrons that move freely through the metal.
  • Hydrogen Bonds: Hydrogen bonds are a type of intermolecular force that occurs when hydrogen atoms are covalently bonded to a highly electronegative atom such as oxygen or nitrogen. The hydrogen atom in the bond is attracted to another electronegative atom nearby.
  • Van der Waals Forces: Van der Waals forces are weak, temporary attractions that occur between atoms and molecules. They are caused by temporary dipoles in the electron distribution of the atoms or molecules.
2.2 Intramolecular force and potential energy
  • Potential Energy: The energy that is stored in an object due to its position, composition, or state. The potential energy of a system depends on the position and interactions between its components. The closer the components are, the more potential energy they have. When the components move apart, the potential energy decreases, and the system releases energy. Potential energy can be calculated using the formula PE = mgh, where m is the mass, g is the acceleration due to gravity, and h is the height.
2.3 Structure of ionic solids
  • Ionic solids are a type of crystalline solid that consists of ions held together by electrostatic forces. Some important subtopics related to the structure of ionic solids are:
  • Crystal Lattice: Ionic solids have a repeating three-dimensional pattern of ions called a crystal lattice. The structure of the crystal lattice is determined by the relative sizes and charges of the ions, as well as by the strength of the electrostatic forces between them.
  • Unit Cell: The smallest repeating unit of the crystal lattice is called the unit cell. There are several types of unit cells, including simple cubic, body-centered cubic, and face-centered cubic. The type of unit cell used depends on the size and arrangement of the ions in the crystal lattice.
  • Coordination Number: The coordination number of an ion in an ionic solid is the number of ions of the opposite charge that surround it in the crystal lattice. The coordination number is determined by the size and charge of the ions.
  • Ionic Radii: The size of an ion affects its position in the crystal lattice and the overall structure of the ionic solid. Ionic radii can be determined experimentally and vary depending on the charge and electron configuration of the ion.
  • Defects in Ionic Solids: Ionic solids can have defects in their crystal lattice due to missing or extra ions. These defects can affect the physical and chemical properties of the ionic solid, such as its electrical conductivity.
  • Applications of Ionic Solids: Ionic solids have many practical applications, including in batteries, fuel cells, and electronic devices. Understanding their structure is important in designing and improving these technologies.
2.4 Structure of metals and alloys
  • Metallic bonding: This type of bonding occurs between metal atoms, where they share their valence electrons and form a lattice structure. The valence electrons are delocalized and not associated with any particular atom.
  • Crystal lattice: The structure of metals and alloys is a three-dimensional lattice structure that consists of cations surrounded by a sea of valence electrons. The arrangement of atoms in the lattice determines the physical properties of the metal.
  • Close-packed structures: There are two types of close-packed structures: hexagonal close-packed (HCP) and cubic close-packed (CCP). These structures are the most efficient way of packing spheres together.
  • Alloys: Alloys are mixtures of two or more metals that are combined to form a material with enhanced properties. The properties of the alloy depend on the types and amounts of metals that are mixed.
  • Substitutional alloys: In substitutional alloys, the atoms of one metal are replaced by atoms of another metal of similar size. This type of alloy is formed when the two metals have similar atomic radii.
  • Interstitial alloys: In interstitial alloys, the smaller atoms of one metal occupy interstitial sites in the crystal lattice of the larger metal. This type of alloy is formed when the two metals have significantly different atomic radii.
  • Solid solutions: A solid solution is a homogenous mixture of two or more substances, where one substance is dissolved in the other. In alloys, the metals form a solid solution with each other.
  • Intermetallic compounds: Intermetallic compounds are formed when two or more metals combine to form a compound with a specific chemical formula. The properties of these compounds can be quite different from those of the constituent metals.
  • Properties of metals and alloys: The properties of metals and alloys depend on their crystal structure, atomic arrangement, and bonding. These properties include conductivity, ductility, malleability, strength, and corrosion resistance.
2.5 Lewis diagrams
  • Lewis Diagrams, also known as Lewis Structures, are diagrams that show the bonding between atoms in a molecule and the placement of lone pairs of electrons.
  • Valence Electrons: The electrons in the outermost shell of an atom that participate in chemical bonding. Lewis diagrams are based on the valence electrons of atoms.
  • Octet Rule: The idea that atoms tend to form bonds in such a way that each atom has a full outer shell of eight electrons (except for hydrogen and helium, which only need two electrons to fill their outer shells).
  • Single, Double, and Triple Bonds: The different types of covalent bonds that can be formed between atoms, depending on the number of electrons that are shared between them.
  • Formal Charge: A method for determining the distribution of electrons in a molecule, and for predicting the most stable Lewis structure for that molecule.
  • Resonance Structures: When there are multiple valid Lewis structures that can be drawn for a molecule, these are called resonance structures. The true structure of the molecule is a hybrid of all the resonance structures.
  • VSEPR Theory: The Valence Shell Electron Pair Repulsion theory predicts the shapes of molecules based on the number of valence electrons and lone pairs of electrons around the central atom.
  • Polarity: The distribution of electric charge across a molecule. Polarity is determined by the electronegativity difference between the atoms in a bond, and the shape of the molecule.
2.6 Resonance and formal charge
  • Resonance: It is a concept in chemistry that describes the distribution of electrons in a molecule or ion when more than one valid Lewis structure can be drawn for that molecule or ion. Resonance structures are imaginary structures that are used to represent the delocalized electrons in the molecule or ion.
  • Formal Charge: It is the charge assigned to an atom in a molecule or ion assuming that electrons in a chemical bond are shared equally between the atoms. The formal charge is calculated by subtracting the number of nonbonding electrons and one-half of the bonding electrons from the total number of valence electrons of the atom.
2.7 VSEPR and bond hybridization
  • VSEPR Theory: The Valence Shell Electron Pair Repulsion (VSEPR) theory is a model that is used to predict the shape of molecules based on the arrangement of their valence electron pairs. According to the theory, electron pairs around an atom repel each other and adopt a geometry that minimizes this repulsion.
  • Bond Hybridization: Bond hybridization refers to the process by which atomic orbitals combine to form hybrid orbitals that are involved in bonding. Hybrid orbitals have shapes and energies that are intermediate between the atomic orbitals from which they were formed.
  • The type of hybridization that occurs in a molecule is determined by the number and geometry of the electron domains (bonds and lone pairs) around the central atom.

Unit 3: Intermolecular Forces and Properties

Subtopic Number Subtopic Name Outline
3.1 Intermolecular forces
  • Intermolecular forces are the forces of attraction or repulsion between molecules. These forces are responsible for determining the physical properties of materials, such as melting and boiling points, viscosity, and solubility.

Types of Intermolecular Forces:

  • London Dispersion Forces: These are the weakest intermolecular forces and exist between all molecules, regardless of polarity. They occur due to the temporary fluctuations in electron density in a molecule.
  • Dipole-Dipole Forces: These are stronger than London Dispersion Forces and occur between polar molecules due to the attraction between their positive and negative ends.
  • Hydrogen Bonding: This is a special type of dipole-dipole force that occurs between a hydrogen atom bonded to an electronegative atom (such as N, O, or F) and a lone pair of electrons on another electronegative atom in a nearby molecule.
  • Ion-Dipole Forces: These occur between an ion and a polar molecule due to the attraction between the ion’s charge and the opposite charge on the polar molecule.
3.2 Properties of solids 
  • Crystalline vs. Amorphous Solids: Crystalline solids are solids that have a regular, repeating arrangement of atoms, ions, or molecules in three dimensions. Amorphous solids, on the other hand, lack this long-range order and have a more random arrangement.
  • Lattice Structures: The lattice structure of a solid is the arrangement of atoms, ions, or molecules in a crystalline solid. Different types of lattice structures include simple cubic, face-centered cubic, and body-centered cubic.
  • Intermolecular Forces: The intermolecular forces between atoms, ions, or molecules in a solid affect its physical properties such as melting point, boiling point, and hardness. Examples of intermolecular forces include London dispersion forces, dipole-dipole interactions, and hydrogen bonding.
  • Thermal Properties: The thermal properties of a solid include its melting point, boiling point, and heat of fusion. The heat of fusion is the amount of energy required to melt a solid. The specific heat capacity is the amount of energy required to raise the temperature of a unit mass of the solid by one degree Celsius.
  • Mechanical Properties: The mechanical properties of a solid include its hardness, ductility, and elasticity. Hardness is the resistance of a solid to indentation or scratching. Ductility is the ability of a solid to be stretched into a wire. Elasticity is the ability of a solid to deform and return to its original shape when a force is applied and then removed.
  • Electrical Properties: The electrical properties of a solid include its conductivity, resistivity, and dielectric constant. Conductivity is the ability of a solid to conduct an electrical current. Resistivity is the resistance of a solid to the flow of an electrical current. Dielectric constant is a measure of a solid’s ability to store electrical energy in an electric field.
3.3 Solids, liquids, and gasses 
  • Kinetic Molecular Theory: The kinetic molecular theory is a model that explains the behavior of matter in different states based on the motion of its constituent particles.
  • Intermolecular Forces: Intermolecular forces are the attractive or repulsive forces that exist between the particles in matter. These forces determine the physical properties of matter, such as melting and boiling points, and the behavior of matter in different states.
  • Gas Laws: Gas laws describe the behavior of gasses under different conditions, such as pressure, temperature, and volume. Some important gas laws include Boyle’s law, Charles’s law, and the ideal gas law.
  • Phase Changes: Phase changes refer to the processes in which matter changes from one state to another, such as melting, boiling, and condensation.
  • Phase Diagrams: Phase diagrams are graphical representations of the behavior of matter at different temperatures and pressures. They show the conditions under which a substance exists as a solid, liquid, or gas.
3.4 Ideal gas law
  • The Ideal Gas Law is a fundamental concept in thermodynamics and chemistry that describes the behavior of gasses under different conditions. It is a mathematical formula that relates the pressure, volume, temperature, and number of particles of a gas. The formula is: PV = nRT
  • Boyle’s Law: describes the relationship between the pressure and volume of a gas at a constant temperature
  • Charles’s Law: describes the relationship between the volume and temperature of a gas at a constant pressure
  • Avogadro’s Law: describes the relationship between the volume and number of particles of a gas at a constant temperature and pressure
  • Dalton’s Law of Partial Pressures: describes the relationship between the total pressure of a gas mixture and the partial pressures of its components
  • Ideal Gas Law Problems: involves applying the Ideal Gas Law to solve problems related to gasses at different conditions
3.5 Kinetic molecular theory
  • The Nature of Gases: The kinetic molecular theory describes the behavior of gasses in terms of the motion of their molecules. The theory explains why glasses have no definite shape or volume and can be compressed easily.
  • Properties of Gases: The kinetic molecular theory explains the properties of gasses, such as their pressure, volume, temperature, and number of molecules. The theory also explains why gasses mix easily and why their pressure is proportional to the number of molecules and their average kinetic energy.
  • Ideal Gas Law: The ideal gas law is a mathematical relationship that describes the behavior of ideal gasses. It relates the pressure, volume, temperature, and number of molecules of a gas. The ideal gas law can be used to predict the behavior of gasses under different conditions.
  • Real Gasses: Real gasses do not behave according to the ideal gas law. The deviations from the ideal gas law are due to the interactions between gas molecules. The behavior of real gasses can be described using equations of state that take into account the attractive and repulsive forces between molecules.
  • Effusion and Diffusion: Effusion is the process by which a gas escapes through a small hole into a vacuum. Diffusion is the process by which two gasses mix and become evenly distributed over time. The rates of effusion and diffusion can be explained by the kinetic molecular theory.
3.6 Deviation from ideal gas law 
  • Real Gasses: In reality, gasses do not always behave ideally, and their behavior can deviate from the predictions of the ideal gas law. The factors that contribute to this deviation include intermolecular forces, gas volume, and gas pressure.
  • Van der Waals Equation: This equation was developed to account for the non-ideal behavior of real gasses by incorporating corrections for the intermolecular forces and the volume of the gas particles. The equation is a modification of the ideal gas law and is used to calculate the behavior of real gasses.
  • Compressibility Factor: The compressibility factor is a measure of how much a gas deviates from ideal behavior. It is defined as the ratio of the actual volume of the gas to the volume that would be predicted by the ideal gas law. The compressibility factor can be used to compare the behavior of different glasses.
  • Critical Temperature and Pressure: The critical temperature and pressure of a gas are the temperature and pressure at which the gas can no longer be liquefied by increasing pressure. Above the critical temperature and pressure, the gas is said to be in a supercritical state, where it exhibits properties of both a gas and a liquid.
  • Liquefaction of Gases: Although gasses generally expand to fill the container they are in, they can be compressed to a smaller volume under the right conditions. At low temperatures and high pressures, gasses can be liquefied. The liquefaction of gasses is an important process in many industrial applications.
  • Phase Diagrams: A phase diagram is a graphical representation of the relationship between the temperature, pressure, and phase of a substance. For gasses, the phase diagram shows the conditions under which the gas can exist as a solid, liquid, or gas. The phase diagram is a useful tool for understanding the behavior of real gasses and for predicting their properties under different conditions.
3.7 Solutions and mixtures
  • Solution: A solution is a homogeneous mixture composed of two or more substances. In a solution, the solute is the substance that is dissolved, while the solvent is the substance that does the dissolving.
  • Solubility: Solubility is the maximum amount of a solute that can be dissolved in a given amount of solvent at a specific temperature and pressure.
  • Concentration: Concentration is the amount of solute present in a given amount of solution. It is usually expressed in moles per liter (M) or as a percentage by weight.
  • Molarity: Molarity is a measure of concentration that represents the number of moles of solute per liter of solution.
  • Molality: Molality is a measure of concentration that represents the number of moles of solute per kilogram of solvent.
  • Colligative Properties: Colligative properties are physical properties of a solution that depend on the concentration of solute particles in the solution. Examples of colligative properties include freezing point depression, boiling point elevation, and vapor pressure lowering.
  • Raoult’s Law: Raoult’s law is a law that describes the vapor pressure of a solution as a function of the vapor pressures and mole fractions of the components of the solution.
  • Henry’s Law: Henry’s law is a law that describes the solubility of a gas in a liquid at a given temperature and pressure.
  • Osmosis: Osmosis is the diffusion of solvent molecules across a semipermeable membrane from a region of lower solute concentration to a region of higher solute concentration.
  • Osmotic pressure: Osmotic pressure is the pressure required to prevent osmosis from occurring between two solutions separated by a semipermeable membrane.
3.8 Representations of solutions
  • Ideal solutions: Solutions that obey Raoult’s law, meaning that the vapor pressure of each component in the solution is proportional to its mole fraction.
  • Non-ideal solutions: Solutions that do not obey Raoult’s law due to interactions between solute particles or between solute and solvent particles.
  • Solubility: The maximum amount of solute that can dissolve in a given amount of solvent at a given temperature and pressure.
  • Saturation: A condition in which a solution contains the maximum amount of solute that can dissolve at a given temperature and pressure.
  • Supersaturation: A condition in which a solution contains more solute than it can hold at a given temperature and pressure.
3.9 Separation of solutions and mixtures chromatography 
  • Chromatography: Chromatography is a technique used to separate the components of a mixture based on their differential distribution between a mobile phase and a stationary phase. There are several types of chromatography techniques, including gas chromatography, liquid chromatography, and paper chromatography.
3.10 Solubility
  • Solubility: Solubility refers to the ability of a substance to dissolve in a solvent and form a homogenous solution.

Factors Affecting Solubility:

  • Temperature: In general, solubility increases with increasing temperature for solids in liquids, and decreases with increasing temperature for gasses in liquids.
  • Pressure: Solubility of gasses in liquids increases with increasing pressure.
  • Nature of Solvent and Solute: The nature of the solvent and solute can also affect solubility. Polar solvents dissolve polar solutes and nonpolar solvents dissolve nonpolar solutes.
  • Units of Solubility: Solubility is usually expressed in terms of the mass of solute that dissolves in a given volume of solvent to form a saturated solution at a specified temperature. The common units of solubility are g/L and mol/L.
  • Solubility Product Constant (Ksp): The solubility product constant is the equilibrium constant for the dissolution of a sparingly soluble salt in water. It is denoted by Ksp and is equal to the product of the ion concentrations raised to their stoichiometric coefficients in the balanced equation for the dissolution of the salt.
  • Common Ion Effect: The common ion effect refers to the decrease in the solubility of an ionic compound caused by the presence of a common ion in the solution.

Applications of Solubility:

  • Determination of the concentration of ions in a solution
  • Purification of substances using recrystallization
  • Separation of ions using fractional precipitation
  • Design of pharmaceuticals
3.11 Spectroscopy and the electromagnetic spectrum
  • Spectroscopy: The study of the interaction between matter and electromagnetic radiation. It involves the use of spectrometers to separate and measure the different wavelengths of light emitted or absorbed by a substance.
  • Electromagnetic Spectrum: The range of all types of electromagnetic radiation. The spectrum includes radio waves, microwaves, infrared radiation, visible light, ultraviolet radiation, X-rays, and gamma rays.
  • Absorption Spectrum: The pattern of absorption of light by a substance as a function of wavelength or frequency. Each substance has a unique absorption spectrum, which can be used to identify the substance or measure its concentration.
  • Emission Spectrum: The pattern of emission of light by a substance as a function of wavelength or frequency. Each substance has a unique emission spectrum, which can be used to identify the substance or measure its concentration.
  • Flame Test: A test that uses the characteristic color of the flame produced by a substance when heated to identify the substance.
  • Mass Spectrometry: A technique that uses an instrument called a mass spectrometer to measure the mass-to-charge ratio of ions. It can be used to identify unknown substances and determine the molecular structure of a compound.
  • Infrared Spectroscopy: A technique that uses infrared radiation to measure the absorption or transmission of light by a substance. It can be used to identify functional groups in organic compounds.
  • Ultraviolet-Visible Spectroscopy: A technique that uses ultraviolet and visible light to measure the absorption or transmission of light by a substance. It can be used to determine the concentration of a substance or identify the presence of certain functional groups.
  • Nuclear Magnetic Resonance (NMR) Spectroscopy: A technique that uses the magnetic properties of atomic nuclei to determine the molecular structure of a compound. It is commonly used in organic chemistry.
3.12 Photoelectric effect
  • The photoelectric effect is the phenomenon where electrons are emitted from a metal surface when it is exposed to electromagnetic radiation of sufficient frequency. 
  • Work function: The minimum energy required to remove an electron from a metal surface is called the work function of that metal. The work function depends on the type of metal and is a fundamental property of that metal.
  • Threshold frequency: The minimum frequency of electromagnetic radiation required to produce the photoelectric effect is called the threshold frequency. The threshold frequency is directly proportional to the work function of the metal surface.
  • Photons: Electromagnetic radiation is quantized in terms of photons. Each photon carries a specific amount of energy and frequency, and this energy is absorbed by the metal surface, leading to the ejection of electrons.
  • Photoelectric current: The number of electrons emitted from the metal surface is proportional to the intensity of the electromagnetic radiation. The flow of these electrons is called the photoelectric current.
3.13 Beer-Lambert Law
  • The Beer-Lambert Law is a fundamental concept in analytical chemistry that describes the relationship between the concentration of a solution and the amount of light that is absorbed by that solution. Here are some subtopics and descriptions related to the Beer-Lambert Law:
  • Absorbance: Absorbance is a measure of how much light is absorbed by a sample at a specific wavelength. The amount of light absorbed is proportional to the concentration of the sample and the thickness of the sample.
  • Transmittance: Transmittance is a measure of the amount of light that passes through a sample at a specific wavelength. It is the complement of absorbance, meaning that 100% transmittance corresponds to 0 absorbance.
  • Beer-Lambert Law equation: The Beer-Lambert Law equation is A = εlc, where A is the absorbance of the sample, ε is the molar absorptivity (a constant that describes how strongly a molecule absorbs light), l is the path length (the distance the light travels through the sample), and c is the concentration of the solution.
  • UV-Visible spectroscopy: UV-Visible spectroscopy is a technique that uses light in the ultraviolet and visible regions of the electromagnetic spectrum to analyze samples. This technique is based on the Beer-Lambert Law and can be used to determine the concentration of a solution or to identify the presence of a specific molecule.
  • Limitations of the Beer-Lambert Law: The Beer-Lambert Law assumes that the sample is homogeneous and that the molecules in the sample do not interact with each other. However, these assumptions may not be valid for complex samples, such as those containing multiple chromophores or in cases where the sample undergoes a chemical reaction.

Unit 4: Chemical Reactions

Subtopic Number Subtopic Name Outline
4.1 Introduction for reactions
  • Chemical reactions involve the breaking and formation of chemical bonds resulting in the conversion of reactants to products with the release or absorption of energy. In order to understand reactions, it is important to understand the nature of chemical bonds and their behavior during chemical reactions.
  • Chemical Bonding: Chemical bonding refers to the forces that hold atoms together in molecules or ions. The nature of chemical bonds is determined by the electron configuration of the atoms involved and their relative electronegativities. The different types of chemical bonds include covalent bonds, ionic bonds, and metallic bonds.
  • Types of Reactions: There are several types of chemical reactions, including synthesis reactions, decomposition reactions, single-displacement reactions, double-displacement reactions, and combustion reactions. These reactions can be classified based on the type of chemical bonds broken and formed during the reaction.
  • Reaction Rates: Reaction rates refer to the speed at which a chemical reaction occurs. Several factors can affect reaction rates, including temperature, concentration of reactants, surface area of reactants, and the presence of catalysts or inhibitors. The rate law and rate constant are used to mathematically describe the relationship between the reaction rate and the concentration of reactants.
  • Equilibrium: Chemical equilibrium is a state in which the forward and reverse reactions occur at equal rates, resulting in no net change in the concentrations of reactants and products. The equilibrium constant is used to mathematically describe the relationship between the concentrations of reactants and products at equilibrium. Le Chatelier’s principle is used to predict the effects of changes in concentration, temperature, and pressure on the position of an equilibrium.
4.2 Net ionic equations
  • A net ionic equation is a chemical equation that shows the species that are involved in a chemical reaction and the net charge of each species, but omits spectator ions.
  • Spectator ions are ions that do not participate in the reaction, they exist in the same form on both the reactant and product sides of the equation.
  • Net ionic equations are useful in determining the species that are actually involved in a chemical reaction, and in understanding the chemical changes that occur during the reaction.
4.3 Representations of reactions
  • Chemical Equations: A representation of a chemical reaction that shows the reactants and products involved in the reaction. Chemical equations are typically written using chemical formulas and symbols to represent the different elements and compounds involved in the reaction.
  • Balancing Equations: Chemical equations must be balanced to satisfy the law of conservation of mass, which states that matter cannot be created or destroyed. 
  • Word Equations: An alternative way of representing chemical reactions that uses words to describe the reactants and products.
  • Skeleton Equations: A shorthand way of representing chemical equations that only shows the chemical formulas of the reactants and products, without indicating the coefficients or balancing the equation.
  • Types of Chemical Reactions: There are several types of chemical reactions including: synthesis, decomposition, single replacement, double replacement, combustion, and acid-base reactions. These types of reactions can be identified by the types of reactants and products involved.
  • Reaction Stoichiometry: The quantitative relationship between the amounts of reactants and products in a chemical reaction. Stoichiometric calculations involve using balanced chemical equations to determine the amounts of reactants needed or products produced in a chemical reaction.
4.4 Physical and chemical changes
  • Physical changes: Physical changes are changes in the state or form of matter that do not involve the creation of new substances. The original substance retains its identity even though it may have a different appearance or physical state. Examples of physical changes include changes in state (such as melting, boiling, freezing), changes in shape (such as cutting or tearing), and changes in size (such as grinding or crushing).
  • Chemical changes: Chemical changes, also known as chemical reactions, involve the creation of new substances through the rearrangement of atoms or molecules. Chemical changes can result in the formation of new substances that have different chemical properties and characteristics than the original substances. Examples of chemical changes include the rusting of iron, burning of wood, and cooking of food. Chemical changes involve the breaking of chemical bonds and the formation of new bonds.
4.5 Stoichiometry
  • Mole Concept: The mole is a unit used in chemistry to express quantities of chemical substances. This subtopic involves the calculation of the number of moles of a given substance in a sample and the relationship between moles, mass, and volume.
  • Chemical Equations: Chemical equations represent the reaction between reactants and products in terms of their molecular formulas. This subtopic involves balancing chemical equations and using stoichiometric coefficients to determine the relative amounts of reactants and products.
  • Limiting Reactants: The limiting reactant is the reactant that is completely consumed in a chemical reaction, thus limiting the amount of product that can be formed. This subtopic involves identifying the limiting reactant and using it to calculate the maximum amount of product that can be formed.
  • Percent Yield: The percent yield is a measure of the efficiency of a chemical reaction, defined as the actual yield of a product divided by the theoretical yield that could be obtained. This subtopic involves calculating percent yield and identifying sources of loss in a chemical reaction.
  • Stoichiometry Calculations: This subtopic involves applying stoichiometric relationships to solve problems related to chemical reactions, such as determining the mass or volume of a reactant or product, or predicting the amount of reactant required to produce a certain amount of product.
4.6 Introduction to titration
  • Titration is a laboratory technique used to determine the concentration of a solution by reacting it with a solution of known concentration. The process involves slowly adding a standard solution to the sample solution until a reaction endpoint is reached, at which point the amount of standard solution added can be used to calculate the concentration of the sample solution.
  • Titration curve: A titration curve is a graph that shows the change in pH or other property of a solution as a function of the amount of titrant added during a titration.
  • Equivalence point: The equivalence point is the point during a titration when the number of moles of the titrant added is stoichiometrically equivalent to the number of moles of the analyte in the sample.
  • Indicator: An indicator is a substance that changes color in response to a change in pH. In acid-base titrations, indicators are often used to signal the endpoint of the titration.
  • Strong acids and bases: Strong acids and bases are substances that completely dissociate in water to form H+ and OH- ions, respectively. Strong acids and bases can be titrated using a standard solution of a strong base or acid, respectively.
  • Weak acids and bases: Weak acids and bases are substances that only partially dissociate in water. Titration curves for weak acid-base systems have a different shape than those for strong acid-base systems, and the equivalence point is often more difficult to determine.
  • Back titration: Back titration is a technique used to determine the concentration of a substance that is difficult to titrate directly. The process involves reacting the substance with an excess of a known standard solution, then titrating the excess solution with a second standard solution to determine the amount of the excess solution remaining. The amount of the original substance can then be calculated from the difference in the amount of excess solution before and after the second titration.
4.7 Types of chemical reactions
  • Combination Reactions: In a combination reaction, two or more reactants combine to form a single product. The general equation for a combination reaction is A + B → AB.
  • Decomposition Reactions: In a decomposition reaction, a single reactant breaks down into two or more products. The general equation for a decomposition reaction is AB → A + B.
  • Combustion Reactions: In a combustion reaction, a fuel (typically a hydrocarbon) reacts with oxygen gas to produce carbon dioxide and water vapor. The general equation for a combustion reaction is CxHy + O2 → CO2 + H2O.
  • Single Displacement Reactions: In a single displacement reaction, an element or ion in a compound is replaced by another element or ion. The general equation for a single displacement reaction is A + BC → AC + B.
  • Double Displacement Reactions: In a double displacement reaction, two compounds exchange ions to form two new compounds. The general equation for a double displacement reaction is AB + CD → AD + CB.
  • Acid-Base Reactions: In an acid-base reaction, an acid and a base react to form salt and water. The general equation for an acid-base reaction is HA + BOH → H2O + BA, where HA is an acid, BOH is a base, and BA is a salt.
  • Redox Reactions: In a redox (reduction-oxidation) reaction, there is a transfer of electrons from one reactant to another. This can occur through the transfer of an electron, as in single displacement reactions, or through the sharing of electrons, as in covalent bonding.
4.8 Introduction to acid-base reactions
  • Acid and base reactions are some of the most common types of chemical reactions in chemistry. They involve the transfer of a proton (H+) from an acid to a base, resulting in the formation of a conjugate acid and a conjugate base. Acids are substances that donate protons while bases are substances that accept protons.
  • Acid-base reactions can be classified as strong or weak depending on the extent of ionization of the acid or base involved. Strong acids and bases completely dissociate in water, while weak acids and bases only partially dissociate.
  • The pH scale is commonly used to measure the acidity or basicity of a solution. It ranges from 0 to 14, with 0 being the most acidic and 14 being the most basic. A pH of 7 is considered neutral.
  • Some common acid-base reactions include neutralization reactions, in which an acid and a base react to form a salt and water, and acid-base titrations, in which a solution of known concentration is added to a solution of unknown concentration to determine its concentration.
4.9 Oxidation-reduction (redox) reactions
  • Oxidation-Reduction (Redox) Reactions: A type of chemical reaction in which one or more electrons are transferred between reactants. Oxidation involves the loss of electrons, while reduction involves the gain of electrons. Redox reactions can be used to generate electricity in batteries and are important in many industrial and biological processes.
  • Half-Reactions: In a redox reaction, each reactant undergoes either oxidation or reduction. These two processes can be separated into half-reactions, which show the transfer of electrons from one species to another.
  • Oxidation Numbers: A way to keep track of the number of electrons gained or lost by an atom in a molecule or ion. Oxidation numbers can be used to identify which species are being oxidized or reduced in a redox reaction.
  • Balancing Redox Reactions: Redox reactions must be balanced to ensure that the number of electrons lost in oxidation is equal to the number of electrons gained in reduction. There are different methods to balance redox reactions, including the half-reaction method and the oxidation-number change method.
  • Applications of Redox Reactions: Redox reactions have many practical applications, such as in the production of metals, the generation of electricity in batteries, and in the processes used to treat wastewater and remove pollutants from the environment.

Unit 5: Kinetics

Subtopic Number Subtopic Name Outline
5.1 Reaction rates
  • Reaction rates: The rate of a chemical reaction is the measure of how quickly reactants are converted to products in a chemical reaction. The subtopic covers:
  • Factors affecting reaction rates: This includes the concentration of reactants, temperature, catalysts, and surface area of reactants.
  • Rate laws: Rate laws describe the mathematical relationship between the concentration of reactants and the rate of the reaction.
  • Reaction mechanisms: Reaction mechanisms describe the series of steps by which a reaction occurs, including the intermediates formed during the reaction.
  • Activation energy: Activation energy is the minimum energy required for a reaction to occur.
  • Collision theory: Collision theory explains how the rate of a chemical reaction depends on the frequency, orientation, and energy of collisions between reacting molecules.
  • Reaction kinetics: Reaction kinetics is the study of the rates of chemical reactions and the factors that affect them, including the study of reaction orders, rate constants, and reaction mechanisms.
5.2 Introduction to Rate law
  • Rate Law: The rate law is a mathematical expression that describes how the rate of a chemical reaction depends on the concentration of reactants. The rate law is determined experimentally, and it allows us to predict the rate of the reaction under different conditions.
  • Reaction Order: The reaction order is a measure of the dependence of the reaction rate on the concentration of reactants. It is determined from the rate law equation and can be fractional or whole numbers.
  • Rate Constant: The rate constant is a proportionality constant that relates the reaction rate to the concentrations of reactants in the rate law equation. It is determined experimentally and is specific to a particular reaction at a particular temperature.
5.3 Concentration changes over time
  • Reaction Rate: The reaction rate is the rate at which the reactants are consumed or the products are formed in a chemical reaction. It is usually expressed as a change in concentration per unit time.
  • Half-Life: The half-life of a reaction is the time it takes for the concentration of a reactant to decrease to half of its initial value. It is a characteristic property of a reaction and depends on the reaction order and rate constant.
  • Integrated Rate Laws: Integrated rate laws are mathematical equations that describe the concentration of reactants or products as a function of time. They can be derived from the rate law equation and can be used to determine the reaction order and rate constant.
5.4 Elementary reactions
  • Elementary reaction: An elementary reaction is a reaction that occurs in a single step and involves only one molecular collision.
  • Reaction mechanism: The reaction mechanism is the sequence of elementary reactions that leads to the overall reaction.
  • Intermediate: An intermediate is a species that is formed in one elementary reaction and then consumed in a subsequent elementary reaction.
  • Catalyst: A catalyst is a substance that increases the rate of a reaction by lowering the activation energy without being consumed in the reaction.
5.5 Collision model
  • The collision model is a concept used in chemical kinetics to explain how chemical reactions occur. It assumes that for a reaction to take place, the reacting molecules must collide with one another with sufficient energy and proper orientation.
  • Activation Energy: Activation energy is the minimum amount of energy required to initiate a chemical reaction. In the collision model, it is the energy required for a collision to result in a chemical reaction.
  • Reaction Rate: The reaction rate is the measure of how quickly a reaction takes place. The collision model helps explain why reaction rates vary depending on the con