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AP CHEMISTRY Unit Guides

UNIT 1. Atomic Structure and Properties

UNIT 2 Compound Structure and Properties

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Guiding question

How does the periodic table help us to predict patterns and trends in the properties of the elements?

Learning outcomes
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • the periodic table consists of periods, groups and blocks

  • the period number shows the outer energy level that is occupied by electrons

  • elements in a group have a common number of valence electrons

  • periodicity refers to trends in properties of elements across a period and down a group

  • trends in properties of elements down a group include the increasing metallic character of group 1

  • elements and decreasing non-metallic character of group 17 elements

  • metallic and non-metallic properties show a continuum; this includes the trend from basic metal

  • oxides through amphoteric to acidic non-metal oxides

  • the oxidation state is a number assigned to an atom to show the number of electrons transferred in

  • forming a bond; it is the charge that atom would have if the compound were composed of ions

  • identify the positions of metals, metalloids and non-metals in the periodic table

  • deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa

  • explain the periodicity of atomic radius, ionic radius, ionization energy, electron affinity and electronegativity

  • describe and explain the reactions of group 1 metals with water, and of group 17 elements with halide ions

  • deduce equations for the reactions with water of the oxides of group 1 and group 2 metals, carbon and sulfur

  • deduce the oxidation state of an atom in an ion or compound

  • explain why the oxidation state of an element is zero

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  • discontinuities occur in the trend of increasing first ionization energy across a period

  • transition elements have incomplete d-sublevels that give them characteristic properties

  • the formation of variable oxidation states in transition elements can be explained by the fact that their successive ionization energies are close in value

  • transition element complexes are coloured due to the absorption of light when an electron is promoted between the orbitals in the split d-sublevels; the colour absorbed is complementary to the colour observed

  • explain how discontinuities in first ionization energy across a period provide evidence for the existence of sublevels

  • recognize properties of transition elements, including: variable oxidation state, high melting points, magnetic properties, catalytic properties, formation of coloured compounds and formation of complex ions with ligands

  • deduce the electron configurations of ions of the first-row transition elements

  • apply the colour wheel to deduce the wavelengths and frequencies of light absorbed and/or observed

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Guiding Question

How does the classification of organic molecules help you to predict their properties?

Learning Outcomes
After studying this topic you should be able to:

Understand

Apply your knowledge to

  • organic compounds can be represented by different types of formulas. These include empirical, molecular, structural (full and condensed), stereochemical and skeletal.

  • functional groups give characteristic physical and chemical properties to a compound.

  • Organic compounds are divided into classes according to the functional groups present in their molecules.

  • a homologous series is a family of compounds in which successive members differ by a common structural unit, typically CH2. Each homologous series can be described by a general formula.

  • successive members of a homologous series show a trend in physical properties

  • “IUPAC nomenclature” refers to a set of rules used by the International Union of Pure and Applied Chemistry to apply systematic names to organic and inorganic compounds.

  • structural isomers are molecules which have the same molecular formula but different connectivities.

  • identify different formulas, and interconvert between molecular, skeletal and structural formulas.

  • construct 3-D models (real or virtual) of organic molecules.

  • identify the following functional groups by name and structure: halogeno, hydroxyl, carbonyl, carboxyl, alkoxy, amino, amido, ester, phenyl.

  • recognize the following homologous series: alkanes, alkenes, alkynes, halogenoalkanes, alcohols, aldehydes, ketones, carboxylic acids, ethers, amines, amides and esters.

  • describe and explain the trend in melting and boiling points of members of a homologous series.

  • Recognize isomers, including branched, straight-chain, position and functional group isomers.

  • Apply IUPAC nomenclature to saturated or mono-unsaturated compounds having up to six carbon atoms in the parent chain and containing one type of the following functional groups: halogeno, hydroxyl, carbonyl, carboxyl (to include straight-chain and branched-chain isomers).

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  • stereoisomers have the same constitution (atom identities, connectivities and bond multiplicities) but different spatial arrangements of atoms.

  • mass spectrometry of organic compounds can cause fragmentation of molecules.

  • infrared (IR) spectra can be used to identify the type of bond present in a molecule.

  • proton nuclear magnetic resonance spectroscopy (1H NMR) gives information on the different chemical environments of hydrogen atoms in a molecule.

  • individual NMR signals can be split into clusters of peaks.

  • data from different techniques are often combined in structural analysis.

  • describe and explain the features which give rise to cis–trans isomerism; recognition in non-cyclic alkenes and C3 and C4 cycloalkanes.

  • drawing stereochemical formulas showing the tetrahedral arrangement around a chiral carbon.

  • describe and explain chiral carbon atom giving rise to stereoisomers with different optical properties.

  • recognize of a pair of enantiomers as non-superimposable mirror images from 3-D modelling (real or virtual).

  • deduce information about the structural features of a compound from specific MS fragmentation patterns.

  • interpret the functional group region of an IR spectrum, using a table of characteristic frequencies (wavenumber / cm−1).

  • Interpret 1H NMR spectra to deduce the structures of organic molecules from the number of signals, the chemical shifts, and the relative areas under signals (integration traces).

  • Interpret 1H NMR spectra from splitting patterns showing singlets, doublets, triplets and quartets to deduce greater structural detail.

  • interpret a variety of data including analytical spectra to determine the structure of a molecule.

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Required Course Content

LEARNING OBJECTIVE

2.3.A Represent an ionic solid with a particulate model that is consistent with Coulomb’s law and the properties of the constituent ions.

ESSENTIAL KNOWLEDGE

2.3.A.1 The cations and anions in an ionic crystal are arranged in a systematic, periodic 3-D array that maximizes the attractive forces among cations and anions while minimizing the repulsive forces.

Exclusion Statement: Knowledge of specific crystal structures is not essential to an understanding of the learning objective and will not be assessed on the AP Exam.

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LEARNING OBJECTIVE

2.4.A Represent a metallic solid and/or alloy using a model to show essential characteristics of the structure and interactions present in the substance.

ESSENTIAL KNOWLEDGE

2.4.A.1 Metallic bonding can be represented as an array of positive metal ions surrounded by delocalized valence electrons (i.e., a “sea of electrons”).

2.4.A.2 Interstitial alloys form between atoms of significantly different radii, where the smaller atoms fill the interstitial spaces between the larger atoms (e.g., with steel in which carbon occupies the interstices in iron).

2.4.A.3 Substitutional alloys form between atoms of comparable radius, where one atom substitutes for the other in the lattice (e.g., in certain brass alloys, other elements, usually zinc, substitute for copper).

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LEARNING OBJECTIVE

2.5.ARepresent a molecule with a Lewis diagram.

ESSENTIAL KNOWLEDGE

2.5.A.1 Lewis diagrams can be constructed according to an established set of principles.

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LEARNING OBJECTIVE

2.6.A Represent a molecule with a Lewis diagram that accounts for resonance between equivalent structures or that uses formal charge to select between nonequivalent structures.

ESSENTIAL KNOWLEDGE

2.6.A.1 In cases where more than one equivalent Lewis structure can be constructed, resonance must be included as a refinement to the Lewis structure. In many such cases, this refinement is needed to provide qualitatively accurate predictions of molecular structure and properties.

2.6.A.2 The octet rule and formal charge can be used as criteria for determining which of several possible valid Lewis diagrams provides the best model for predicting molecular structure and properties.

2.6.A.3 As with any model, there are limitations to the use of the Lewis structure model, particularly in cases with an odd number of valence electrons.

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LEARNING OBJECTIVE

2.7.A Based on the relationship between Lewis diagrams, VSEPR theory, bond orders, and bond polarities:VSEPR theory uses the Coulombic repulsion between electrons as a basis for predicting the arrangement of electron pairs around a central atom.

i. Explain structural properties of molecules.

ii. Explain electron properties of molecules.

ESSENTIAL KNOWLEDGE

2.7.A.1 VSEPR theory uses Coulombic repulsion between electron pairs around a central atom to predict:

i. Structural properties of molecules.

ii. Electron properties of molecules.

2.7.A.2 Both Lewis diagrams and VSEPR theory must be used for predicting electronic and structural properties of many covalently bonded molecules and polyatomic ions, including:

i. Molecular geometry (linear, trigonal planar, tetrahedral, trigonal pyramidal, bent, trigonal bipyramidal, seesaw, T-shaped, octahedral, square pyramidal, square planar)

ii. Bond angles

iii. Relative bond energies based on bond order

iv. Relative bond lengths (multiple bonds, effects of atomic radius)

v. Presence of a dipole moment

vi. Hybridization of valence orbitals for atoms within a molecule or polyatomic ion.

2.7.A.3 The terms “hybridization” and “hybrid atomic orbital” describe the arrangement of electrons around a central atom:- sp hybridized atoms have ideal bond angles of 180°- sp² hybridized atoms have ideal bond angles of 120°- sp³ hybridized atoms have ideal bond angles of 109.5°

Exclusion Statements:

-Understanding of the derivation and depiction of hybrid orbitals will not be assessed on the AP Exam. The course includes the distinction between sigma and pi bonding, the use of VSEPR to explain molecular shapes, and the sp, sp², and sp³ nomenclature.

– Hybridization involving d orbitals will not be assessed on the AP Exam.

– When an atom has more than four pairs of electrons surrounding the central atom, students are only responsible for the shape of the resulting molecule.

2.7.A.4 Bond formation is associated with overlap between atomic orbitals. In multiple bonds, such overlap leads to the formation of both sigma and pi bonds. The overlap is stronger in sigma than pi bonds, which is reflected in sigma bonds having greater bond energy than pi bonds. The presence of a pi bond also prevents rotation of the bond and leads to geometric isomers.

Exclusion Statement: Molecular orbital theory is recommended for deeper bonding insight but will not be explicitly assessed on the AP Exam, including molecular orbital diagrams, filling, or distinctions between bonding, nonbonding, and antibonding orbitals.

UNIT 3 Properties of Substances and Mixtures

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Required Course Content

LEARNING OBJECTIVE

3.1.A Explain the relationship between the chemical structures of molecules and the relative strength of their intermolecular forces when:i. The molecules are of the same chemical species.ii. The molecules are of two different chemical species.

ESSENTIAL KNOWLEDGE

3.1.A.1 London dispersion forces are a result of the Coulombic interactions between temporary, fluctuating dipoles. London dispersion forces are often the strongest net intermolecular force between large molecules.

i. Dispersion forces increase with increasing contact area between molecules and with increasing polarizability of the molecules.

ii. The polarizability of a molecule increases with an increasing number of electrons in the molecule and the size of the electron cloud. It is enhanced by the presence of pi bonding.

iii. The term “London Dispersion Forces” should not be used synonymously with the term “van der Waals forces.”

3.1.A.2 The dipole moment of a polar molecule leads to additional interactions with other chemical species.

i. Dipole-induced dipole interactions are present between a polar and nonpolar molecule. These forces are always attractive. The strength of these forces increases with the magnitude of the dipole of the polar molecule and with the polarizability of the nonpolar molecule.

ii. Dipole-dipole interactions are present between polar molecules. The interaction strength depends on the magnitudes of the dipoles and their relative orientation. Interactions between polar molecules are typically greater than those between nonpolar molecules of comparable size because these interactions act in addition to London dispersion forces.

iii. Ion-dipole forces of attraction are present between ions and polar molecules. These tend to be stronger than dipole-dipole forces.

3.1.A.3 The relative strength and orientation dependence of dipole-dipole and ion-dipole forces can be understood qualitatively by considering the sign of the partial charges responsible for the molecular dipole moment, and how these partial charges interact with an ion or with an adjacent dipole.

3.1.A.4 Hydrogen bonding is a strong type of intermolecular interaction that exists when hydrogen atoms covalently bonded to the highly electronegative atoms (N, O, and F) are attracted to the negative end of a dipole formed by the electronegative atom (N, O, and F) in a different molecule, or a different part of the same molecule.

3.1.A.5 In large biomolecules, noncovalent interactions may occur between different molecules or between different regions of the same large biomolecule.

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LEARNING OBJECTIVE
3.2.A Explain the relationship among the macroscopic properties of a substance, the particulate-level structure of the substance, and the interactions between these particles.
 
ESSENTIAL KNOWLEDGE

3.2.A.1 Many properties of liquids and solids are determined by the strengths and types of intermolecular forces present. Because intermolecular interactions are overcome completely when a substance vaporizes, the vapor pressure and boiling point are directly related to the strength of those interactions. Melting points also tend to correlate with interaction strength, but because the interactions are only rearranged in melting, the relations can be more subtle.

3.2.A.2 Particulate-level representations, showing multiple interacting chemical species, are a useful means to communicate or understand how intermolecular interactions help to establish macroscopic properties.

3.2.A.3 Due to strong interactions between ions, ionic solids tend to have low vapor pressures, high melting points, and high boiling points. They tend to be brittle due to the repulsion of like charges caused when one layer slides across another layer. They conduct electricity only when the ions are mobile, as when the ionic solid is melted (i.e., in a molten state) or dissolved in water or another solvent.

3.2.A.4 In covalent network solids, the atoms are covalently bonded together into a three-dimensional network (e.g., diamond) or layers of two-dimensional networks (e.g., graphite). These are only formed from nonmetals and metalloids: elemental (e.g., diamond, graphite) or binary compounds (e.g., silicon dioxide and silicon carbide). Due to the strong covalent interactions, covalent solids have high melting points. Three-dimensional network solids are also rigid and hard because the covalent bond angles are fixed. However, graphite is soft because adjacent layers can slide past each other relatively easily.

3.2.A.5 Molecular solids are composed of distinct, individual units of covalently bonded molecules attracted to each other through relatively weak intermolecular forces. Molecular solids generally have a low melting point because of the relatively weak intermolecular forces present between the molecules. They do not conduct electricity because their valence electrons are tightly held within the covalent bonds and the lone pairs of each constituent molecule. Molecular solids are sometimes composed of very large molecules or polymers.

3.2.A.6 Metallic solids are good conductors of electricity and heat, due to the presence of free valence electrons. They also tend to be malleable and ductile, due to the ease with which the metal cores can rearrange their structure. In an interstitial alloy, interstitial atoms tend to make the lattice more rigid, decreasing malleability and ductility. Alloys typically retain a sea of mobile electrons and so remain conducting.

3.2.A.7 In large biomolecules or polymers, noncovalent interactions may occur between different molecules or between different regions of the same large biomolecule. The functionality and properties of such molecules depend strongly on the shape of the molecule, which is largely dictated by noncovalent interactions. |

 
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Required Course Content
LEARNING OBJECTIVE
3.3.A
Represent the differences between solid, liquid, and gas phases using a particulate-level model.
ESSENTIAL KNOWLEDGE
3.3.A.1
Solids can be crystalline, where the particles are arranged in a regular three-dimensional structure, or they can be amorphous, where the particles do not have a regular, orderly arrangement. In both cases, the motion of the individual particles is limited, and the particles do not undergo overall translation with respect to each other. The structure of the solid is influenced by interparticle interactions and the ability of the particles to pack together.
3.3.A.2
The constituent particles in liquids are in close contact with each other, and they are continually moving and colliding. The arrangement and movement of particles are influenced by the nature and strength of the forces (e.g., polarity, hydrogen bonding, and temperature) between the particles.
3.3.A.3
The solid and liquid phases for a particular substance typically have similar molar volume because, in both phases, the constituent particles are in close contact at all times.
3.3.A.4
In the gas phase, the particles are in constant motion. Their frequencies of collision and the average spacing between them are dependent on temperature, pressure, and volume. Because of this constant motion, and minimal effects of forces between particles, a gas has neither a definite volume nor a definite shape.
Exclusion Statement: Understanding/interpreting phase diagrams will not be assessed on the AP Exam.
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LEARNING OBJECTIVE
3.4.A Explain the relationship between the macroscopic properties of a sample of gas or mixture of gases using the ideal gas law.
ESSENTIAL KNOWLEDGE

3.4.A.1 The macroscopic properties of ideal gases are related through the ideal gas law:

EQN: PV = nRT.

3.4.A.2 In a sample containing a mixture of ideal gases, the pressure exerted by each component (the partial pressure) is independent of the other components. Therefore, the partial pressure of a gas within the mixture is proportional to its mole fraction (X), and the total pressure of the sample is the sum of the partial pressures.

EQN: Pₐ = Pₜₒₜₐₗ × Xₐ,

where Xₐ = moles A / total moles;

EQN: Pₜₒₜₐₗ = Pₐ + Pᵦ + P𝒸 + …

3.4.A.3 Graphical representations of the relationships between P, V, T, and n are useful to describe gas behavior.
 
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LEARNING OBJECTIVE
3.5.A Explain the relationship between the motion of particles and the macroscopic properties of gases with:
i. The kinetic molecular theory (KMT).
ii. A particulate model.
iii. A graphical representation.
ESSENTIAL KNOWLEDGE
3.5.A.1 The kinetic molecular theory (KMT) relates the macroscopic properties of gases to motions of the particles in the gas. The Maxwell-Boltzmann distribution describes the distribution of the kinetic energies of particles at a given temperature.
3.5.A.2 All the particles in a sample of matter are in continuous, random motion. The average kinetic energy of a particle is related to its average velocity by the equation:
KE = ½ mv²
3.5.A.3 The Kelvin temperature of a sample of matter is proportional to the average kinetic energy of the particles in the sample.
3.5.A.4 The Maxwell-Boltzmann distribution provides a graphical representation of the energies/velocities of particles at a given temperature.
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LEARNING OBJECTIVE
3.6.A Explain the relationship among non-ideal behaviors of gases, interparticle forces, and/or volumes.
ESSENTIAL KNOWLEDGE
3.6.A.1 The ideal gas law does not explain the actual behavior of real gases. Deviations from the ideal gas law may result from interparticle attractions among gas molecules, particularly at conditions that are close to those resulting in condensation. Deviations may also arise from particle volumes, particularly at extremely high pressures.
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LEARNING OBJECTIVE
3.7.A Calculate the number of solute particles, volume, or molarity of solutions.
ESSENTIAL KNOWLEDGE
3.7.A.1 Solutions, also sometimes called homogeneous mixtures, can be solids, liquids, or gases. In a solution, the macroscopic properties do not vary throughout the sample. In a heterogeneous mixture, the macroscopic properties depend on location in the mixture.
3.7.A.2 Solution composition can be expressed in a variety of ways; molarity is the most common method used in the laboratory.
Equation: M = nₛₒₗᵤₜₑ / Lₛₒₗᵤₜᵢₒₙ
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Required Course Content
LEARNING OBJECTIVE

3.8.A Using particulate models for mixtures:

i. Represent interactions between components.

ii. Represent concentrations of components.

ESSENTIAL KNOWLEDGE
3.8.A.1 Particulate representations of solutions communicate the structure and properties of solutions, by illustration of the relative concentrations of the components in the solution and/or drawings that show interactions among the components. 
Exclusion Statement: Colligative properties will not be assessed on the AP Exam.
Exclusion Statement: Calculations of molality, percent by mass, and percent by volume for solutions will not be assessed on the AP Exam.
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Required Course Content
LEARNING OBJECTIVE
3.10.A Explain the relationship between the solubility of ionic and molecular compounds in aqueous and nonaqueous solvents, and the intermolecular interactions between particles.
ESSENTIAL KNOWLEDGE
3.10.A.1 Substances with similar intermolecular interactions tend to be miscible or soluble in one another.
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LEARNING OBJECTIVE
3.11.A Explain the relationship between a region of the electromagnetic spectrum and the types of molecular or electronic transitions associated with that region.
ESSENTIAL KNOWLEDGE
3.11.A.1 Differences in absorption or emission of photons in different spectral regions are related to the different types of molecular motion or electronic transition:
i. Microwave radiation is associated with transitions in molecular rotational levels.
ii. Infrared radiation is associated with transitions in molecular vibrational levels.
iii. Ultraviolet/visible radiation is associated with transitions in electronic energy levels.
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LEARNING OBJECTIVE
3.12.A Explain the properties of an absorbed or emitted photon in relationship to an electronic transition in an atom or molecule.
ESSENTIAL KNOWLEDGE
3.12.A.1 When a photon is absorbed (or emitted) by an atom or molecule, the energy of the species is increased (or decreased) by an amount equal to the energy of the photon.
3.12.A.2 The wavelength of the electromagnetic wave is related to its frequency and the speed of light by the equation:
EQN: c = λν.
The energy of a photon is related to the frequency of the electromagnetic wave through Planck’s equation:
EQN: E = hν.
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LEARNING OBJECTIVE

3.13.A Explain the amount of light absorbed by a solution of molecules or ions in relationship to the concentration, path length, and molar absorptivity.

ESSENTIAL KNOWLEDGE

3.13.A.1 The Beer-Lambert law relates the absorption of light by a solution to three variables according to the equation:
EQN: A = εbc.

The molar absorptivity, ε, describes how intensely a chemical species absorbs light of a specific wavelength. The path length, b, and concentration, c, are proportional to the number of light-absorbing particles in the light path.

3.13.A.2 In most experiments the path length and wavelength of light are held constant. In such cases, the absorbance is proportional only to the concentration of absorbing molecules or ions. The spectrophotometer is typically set to the wavelength of maximum absorbance (optimum wavelength) for the species being analyzed to ensure the maximum sensitivity of measurement.

UNIT 4 Chemical Reactions

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Guiding Question

What can be deduced from the temperature change that accompanies chemical or physical change?​

Learning outcomes.
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • Chemical reactions involve a transfer of energy between the system and the surroundings, while total energy is conserved.

  • Reactions are described as endothermic or exothermic, depending on the direction of energy transfer between the system and the surroundings.

  • Temperature changes accompany endothermic and exothermic reactions

  • The relative stability of reactants and products determines whether reactions are endothermic or exothermic.

  • The standard enthalpy change for a chemical reaction, ΔH⦵, refers to the heat transferred at constant pressure under standard conditions and states. It can be determined from the change in temperature of a pure substance.

  • Distinguish between heat and temperature.

  • Understand the temperature change (decrease or increase) that accompanies endothermic and exothermic reactions, respectively.

  • Sketch and interpret potential energy profiles for endothermic and exothermic reactions.

  • apply the equations Q = mcΔT and ΔH⦵ = − Q/n in the calculation of the enthalpy change of a reaction.​

IB Clarification Notes

  • Axes for energy profiles should be labelled as reaction coordinate (x), potential energy (y).

  • The units of ΔH⦵ are kJ mol-1.

  • The equation Q = mcΔT and the value of c, the specific heat capacity of water, are given in the data booklet

  • There is no higher-level only material in R1.1.

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Guiding Question

How does application of the law of conservation of energy help us to predict energy changes during reactions?

Learning outcomes.
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • Bond-breaking absorbs and bond-forming releases energy.

  • Hess’s law states that the enthalpy change for a reaction is independent of the pathway between the initial and final states.

  • Calculate the enthalpy change of a reaction from given average bond enthalpy data.

  • Apply hess’s law to calculate enthalpy changes in multi-step reactions.

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  • Standard enthalpy changes of combustion, ΔHc⦵ and formation, ΔHf⦵, data are used in thermodynamic calculations.

  • An application of Hess’s Law uses enthalpy of formation and/or enthalpy of combustion to calculate the enthalpy change of a reaction.

  • A Born-Haber cycle is an application of Hess’s law, used to show energy changes in the formation of an ionic compound.

  • Deduce equations and solutions to problems involving these terms.

  • Calculate enthalpy changes of a reaction using ΔHf⦵ data or ΔHc⦵ data:

ΔH⦵ = Σ ΔHf⦵products − Σ ΔHf⦵reactants
ΔH⦵ = Σ ΔHc⦵reactants − Σ ΔHc⦵products

  • Interpret and determine values from a Born–Haber cycle for compounds composed of univalent and divalent ions.

IB Clarification Notes

  • Enthalpy of combustion and formation data are given in the data booklet.

  • The equations to determine the enthalpy change of a reaction using data are given in the data booklet

ΔH⦵ = Σ ΔHf⦵products − Σ ΔHf ⦵reactants
ΔH⦵ = Σ ΔHc⦵reactants − Σ ΔHc⦵products

  • The Born-Haber cycle includes ionization energies, enthalpy of atomization (using sublimation and/or bond enthalpies), electron affinities, lattice enthalpy, enthalpy of formation.

  • The construction of a complete Born-Haber cycle will not be assessed.

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Guiding question

What are the challenges of using chemical energy to address our energy needs?

Learning outcomes
​After studying this topic students should be able to:

Understand

Apply their knowledge to

  • Reactive metals, non-metals and organic compounds undergo combustion reactions when heated in oxygen.

  • Incomplete combustion of organic compounds, especially hydrocarbons, leads to the production of carbon monoxide and carbon.

  • Fossil fuels, which include coal, crude oil and natural gas, have different advantages and disadvantages.

  • Understand the link between carbon dioxide levels and the greenhouse effect.

  • Biofuels are produced from the biological fixation of carbon over a short period of time through photosynthesis.

  • Understand the difference between renewable and non-renewable energy sources.

  • A fuel cell can be used to convert chemical energy from a fuel directly to electrical energy.

  • Deduce equations for reactions of combustion, including hydrocarbons and alcohols.

  • Deduce equations for the incomplete combustion of hydrocarbons and alcohols.

  • Evaluate the amount of carbon dioxide added to the atmosphere when different fuels burn.

  • Consider the advantages and disadvantages of biofuels.

  • Deduce the half-equations for the electrode reactions in a fuel cell.

IB Clarification Notes

  • The tendency for incomplete combustion and energy released per unit mass should be covered.

  • The reactants and products of photosynthesis should be known.

  • Hydrogen and methanol should be considered as fuels for fuel cells.

  • The use of proton exchange membranes will not be assessed.

  • There is no higher-level only material in R1.3.

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Guiding Question

What determines the direction of chemical change?

Learning outcomes
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • Entropy, S, is a measure of the dispersal or distribution of matter and/or energy in a system.

  • The more ways the energy can be distributed, the higher the entropy.

  • Under the same conditions, entropy of gas > liquid > solid.

  • Change in Gibbs energy, ΔG, relates the energy that can be obtained from a chemical reaction to the change in enthalpy, ΔH, change in entropy, ΔS, and absolute temperature, T.

  • At constant pressure, a change is spontaneous if the change in Gibbs energy, ΔG, is negative.

  • As a reaction approaches equilibrium, ΔG becomes less negative and finally reaches zero.

  • Predict whether a physical or chemical change will result in an increase or decrease in entropy of a system.

  • Calculate standard entropy changes, ΔS⦵ , from standard entropy values, S⦵.

  • Apply the equation ΔG⦵ = ΔH⦵ – TΔS⦵ to calculate unknown values of these terms.

  • Interpret the sign of ΔG calculated from thermodynamic data.

  • Determine the temperature at which a reaction becomes spontaneous.

  • Perform calculations using the equation ΔG = ΔG⦵ + RT ln Q and its application to a system at equilibrium ΔG⦵ = − RT ln Kc.

IB Clarification Notes

  • Standard entropy values are given in the data booklet.

  • Thermodynamic data values are given in the data booklet.

  • Note the units: ΔH kJ mol−1; ΔS J K-1 mol−1; ΔG kJ mol−1

  • ΔG considers the direct entropy change resulting from the transformation of the chemicals and the indirect entropy change of the surroundings as a result of the transfer of heat energy.

  • The equations ΔG = ΔG⦵ + RT ln Q and ΔG⦵ = − RT ln Kc are given in the data booklet.

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LEARNING OBJECTIVE

4.5.A Explain changes in the amounts of reactants and products based on the balanced reaction equation for a chemical process.

ESSENTIAL KNOWLEDGE

4.5.A.1 Because atoms must be conserved during a chemical process, it is possible to calculate product amounts by using known reactant amounts, or to calculate reactant amounts given known product amounts.

4.5.A.2 Coefficients of balanced chemical equations contain information regarding the proportionality of the amounts of substances involved in the reaction. These values can be used in chemical calculations involving the mole concept.

4.5.A.3 Stoichiometric calculations can be combined with the ideal gas law and calculations involving molarity to quantitatively study gases and solutions.

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LEARNING OBJECTIVE
4.6.A
Identify the equivalence point in a titration based on the amounts of the titrant and analyte, assuming the titration reaction goes to completion.
ESSENTIAL KNOWLEDGE
4.6.A.1
Titrations may be used to determine the amount of an analyte in solution. The titrant has a known concentration of a species that reacts specifically and quantitatively with the analyte. The equivalence point of the titration occurs when the analyte is totally consumed by the reacting species in the titrant. The equivalence point is often indicated by a change in a property (such as color) that occurs when the equivalence point is reached. This observable event is called the endpoint of the titration.
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LEARNING OBJECTIVE
4.7.A
Identify a reaction as acid-base, oxidation-reduction, or precipitation.
ESSENTIAL KNOWLEDGE
4.7.A.1
Acid-base reactions involve transfer of one or more protons (H⁺ ions) between chemical species.
4.7.A.2
Oxidation-reduction (redox) reactions involve transfer of one or more electrons between chemical species, as indicated by changes in oxidation numbers of the involved species. Combustion is an important subclass of oxidation-reduction reactions, in which a species reacts with oxygen gas. In the case of hydrocarbons, carbon dioxide and water are products of complete combustion.
4.7.A.3
In a redox reaction, electrons are transferred from the species that is oxidized to the species that is reduced.
Exclusion Statement: The meaning of the terms “reducing agent” and “oxidizing agent” will not be assessed on the AP Exam.
4.7.A.4
Oxidation numbers may be assigned to each of the atoms in the reactants and products; this is often an effective way to identify the oxidized and reduced species in a redox reaction.
4.7.A.5
Precipitation reactions frequently involve mixing ions in aqueous solution to produce an insoluble or sparingly soluble ionic compound. All sodium, potassium, ammonium, and nitrate salts are soluble in water.
Exclusion Statement: Rote memorization of “solubility rules” other than those implied in 4.7.A.5 will not be assessed on the AP Exam.
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Required Course Content
LEARNING OBJECTIVE
4.8.A
Identify species as Brønsted-Lowry acids, bases, and/or conjugate acid-base pairs, based on proton-transfer involving those species.
ESSENTIAL KNOWLEDGE
4.8.A.1
Identify species as Brønsted-Lowry acids, bases, and/or conjugate acid-base pairs, based on proton-transfer involving those species. By definition, a Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor.
4.8.A.2
Only in aqueous solutions, water plays an important role in many acid-base reactions, as its molecular structure allows it to accept protons from and donate protons to dissolved species.
4.8.A.3
When an acid or base ionizes in water, the conjugate acid-base pairs can be identified and their relative strengths compared.
Exclusion Statement: Lewis acid-base concepts will not be assessed on the AP Exam. The emphasis in AP Chemistry is on reactions in aqueous solution.
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Required Course Content
LEARNING OBJECTIVE
4.9.A
Represent a balanced redox reaction equation using half-reactions.
ESSENTIAL KNOWLEDGE
4.9.A.1
Balanced chemical equations for redox reactions can be constructed from half-reactions.

UNIT 5 Kinetics

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Guiding Question

How are chemical equations used to calculate reacting ratios?

Learning outcomes
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • Chemical equations show the ratio of reactants and products in a reaction.

  • The mole ratio of an equation can be used to determine:

→ The masses and/or volumes of reactants and products.

→ The concentrations of reactants and products for reactions occurring in solution.

  • The limiting reactant determines the theoretical yield.

  • The percentage yield is calculated from the ratio of experimental yield to theoretical yield.

  • The atom economy is a measure of efficiency in Green Chemistry.

  • Deduce chemical equations when reactants and products are specified.

  • Calculate reacting masses and/or volumes and concentrations of reactants and products.

  • Identify the limiting and excess reactants from given data.

  • Distinguish between the theoretical yield and the experimental yield

  • Solve problems involving reacting quantities, limiting and excess reactants, theoretical, experimental and percentage yields.

  • Calculate the atom economy from the stoichiometry of a reaction.

IB Clarification Notes

  • Include the use of state symbols in chemical equations.

  • Avogadro’s law and definitions of molar concentration are covered in Structure 1.4. The values for Ar given in the data booklet to two decimal places should be used in calculations.

  • Include discussion of the inverse relationship between atom economy and wastage in industrial processes.

  • The equation for calculation of the atom economy is given in the data booklet.

  • There is no higher-level only material in R2.1

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Guiding question

How can the rate of a reaction be controlled?

Learning outcomes
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • the rate of reaction is expressed as the change in concentration of a particular reactant/product per unit time.

  • species react as a result of collisions of sufficient energy and proper orientation.

  • factors that influence the rate of a reaction include pressure/concentration, surface area, temperature and the presence of a catalyst.

  • activation energy, Ea, is the minimum energy that colliding particles need for a successful collision leading to a reaction.

  • catalysts increase the rate of reaction by providing an alternative reaction pathway with lower Ea.

  • determine rates of reaction

  • explain the relationship between the kinetic energy of the particles and the temperature in kelvin, and the role of collision geometry.

  • predict and explain the effects of changing conditions on the rate of a reaction.

  • construct Maxwell-Boltzmann energy distribution curves to explain the effect of temperature on the probability of successful collisions.

  • sketch and explain energy profiles with and without catalysts, for endothermic and exothermic reactions.

  • construct Maxwell-Boltzmann energy distribution curves to explain the effect of different values for Ea on the probability of successful collisions affecting these, including the effect of a catalyst.

AHL

AHL

  • many reactions occur in a series of elementary steps, and the slowest step determines the rate of the reaction.

  • energy profiles can be used to show the activation energy and transition state of the rate determining step in a multi-step reaction.

  • the molecularity of an elementary step is the number of reacting particles taking part in that step.

  • rate equations depend on the mechanism of the reaction and can only be determined experimentally.

  • the order of a reaction with respect to a reactant is the exponent to which the concentration of the reactant is raised in the rate equation.

  • the order with respect to a reactant can describe the number of particles taking part in the rate-determining step.

  • the overall reaction order is the sum of the orders with respect to each reactant.

  • the rate constant, k, is temperature dependent and its units are determined from the overall order of the reaction.

  • the Arrhenius equation uses the temperature dependence of the rate constant to determine the activation energy.

  • the Arrhenius factor, A, takes into account the frequency of collisions with proper orientations.

  • evaluate proposed reaction mechanisms and recognise reaction intermediates.

  • distinguish between intermediates and transition states, and recognise both in energy profiles of reactions.

  • construct and interpret energy profiles from kinetic data.

  • interpret the terms unimolecular, bimolecular and termolecular.

  • deduce the rate equation for a reaction from experimental data.

  • sketch, identify and analyse graphical representations of zero, first and second order reactions.

  • solve problems involving the rate equation, including the units of k.

  • analyse graphical representations of the Arrhenius equation, including its linear form.

  • determine the activation energy and the Arrhenius factor from experimental data.

IB Clarification notes

  • Calculation of reaction rates from tangents of graphs of concentration, volume or mass against time should be covered.

  • Biological catalysts are called enzymes.

  • The different mechanisms of homogeneous and heterogeneous catalysts will not be assessed.

AHL

  • include examples where the rate-determining step is not the first step.

  • Proposed reaction mechanisms must be consistent with kinetic and stoichiometric data.

  • Concentration-time and rate-concentration graphs should be included.

  • Only integer values for order of reaction will be assessed.

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Guiding question

How can the extent of a reversible reaction be influenced?

Learning outcomes
After studying this topic students should be able to:

Understand

Apply their knowledge to

  • a state of dynamic equilibrium is reached in a closed system when the rates of forward and backward reactions are equal

  • the equilibrium law describes how the equilibrium constant, K, can be determined from the stoichiometry of a reaction

  • the magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent

  • Le Châtelier’s principle enables the prediction of the qualitative effects of changes in concentration, temperature and pressure on a system at equilibrium

  • describe the characteristics of a physical and chemical system at equilibrium.

  • deduce the equilibrium constant expression from an equation for a homogeneous reaction.

  • determine the relationships between K values for reactions that are the reverse of each other at the same temperature.

  • apply Le Châtelier’s principle to predict and explain responses to changes of systems at equilibrium.

AHL

AHL

  • the reaction quotient, Q, is calculated using the equilibrium expression with non-equilibrium concentrations of reactants and products.

  • the equilibrium law is the basis for quantifying the composition of an equilibrium mixture.

  • the equilibrium constant and Gibbs energy change, ΔG, can both be used to measure the position of an equilibrium reaction

  • calculate the reaction quotient, Q, from the concentrations of reactants and products at a particular time, and determine the direction in which the reaction will proceed to reach equilibrium.

  • solve problems involving values of K and initial and equilibrium concentrations of the components of an equilibrium mixture.

IB Clarification Notes

  • Include the extent of reaction for: K<<1, K<1, K=1, K>1, K>>1

  • Include the effects of K on the equilibrium position.

  • LeChatlier’s principle can be applied to heterogenous equilibria such as: X(g)⇌X(aq)

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Required Course Content
LEARNING OBJECTIVE
5.4.A Represent an elementary reaction as a rate law expression using stoichiometry.
ESSENTIAL KNOWLEDGE
5.4.A.1 The rate law of an elementary reaction can be inferred from the stoichiometry of the particles participating in a collision.
5.4.A.2 Elementary reactions involving the simultaneous collision of three or more particles are rare.
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Required Course Content
LEARNING OBJECTIVE
5.5.A Explain the relationship between the rate of an elementary reaction and the frequency, energy, and orientation of particle collisions.
ESSENTIAL KNOWLEDGE
5.5.A.1 For an elementary reaction to successfully produce products, reactants must successfully collide to initiate bond-breaking and bond-making events.
5.5.A.2 In most reactions, only a small fraction of the collisions leads to a reaction. Successful collisions have both sufficient energy to overcome the activation energy requirements and orientations that allow the bonds to rearrange in the required manner.
5.5.A.3 The Maxwell-Boltzmann distribution curve describes the distribution of particle energies; this distribution can be used to gain a qualitative estimate of the fraction of collisions with sufficient energy to lead to a reaction, and also how that fraction depends on temperature.
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Required Course Content
LEARNING OBJECTIVE
5.6.A Represent the activation energy and overall energy change in an elementary reaction using a reaction energy profile.
ESSENTIAL KNOWLEDGE
5.6.A.1 Elementary reactions typically involve the breaking of some bonds and the forming of new ones.
5.6.A.2 The reaction coordinate is the axis along which the complex set of motions involved in rearranging reactants to form products can be plotted.
5.6.A.3 The energy profile gives the energy along the reaction coordinate, which typically proceeds from reactants, through a transition state, to products. The energy difference between the reactants and the transition state is the activation energy for the forward reaction.
5.6.A.4 The rate of an elementary reaction is temperature dependent because the proportion of particle collisions that are energetic enough to reach the transition state varies with temperature. The Arrhenius equation relates the temperature dependence of the rate of an elementary reaction to the activation energy needed by molecular collisions to reach the transition state.
Exclusion Statement: Calculations involving the Arrhenius equation will not be assessed on the AP Exam.
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Required Course Content
LEARNING OBJECTIVE
5.7.A Identify the components of a reaction mechanism.
ESSENTIAL KNOWLEDGE
5.7.A.1 A reaction mechanism consists of a series of elementary reactions, or steps, that occur in sequence. The components may include reactants, intermediates, products, and catalysts.
5.7.A.2 The elementary steps when combined should align with the overall balanced equation of a chemical reaction.
5.7.A.3 A reaction intermediate is produced by some elementary steps and consumed by others, such that it is present only while a reaction is occurring.
5.7.A.4 Experimental detection of a reaction intermediate is a common way to build evidence in support of one reaction mechanism over an alternative mechanism.
Exclusion Statement: Collection of data pertaining to detection of a reaction intermediate will not be assessed on the AP Exam.
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Required Course Content
LEARNING OBJECTIVE
5.8.A Identify the rate law for a reaction from a mechanism in which the first step is rate limiting.
ESSENTIAL KNOWLEDGE
5.8.A.1 For reaction mechanisms in which each elementary step is irreversible, or in which the first step is rate limiting, the rate law of the reaction is set by the molecularity of the slowest elementary step (i.e., the rate-limiting step).
Exclusion Statement: Collection of data pertaining to detection of a reaction intermediate will not be assessed on the AP Exam.
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Required Course Content
LEARNING OBJECTIVE
5.9.A Identify the rate law for a reaction from a mechanism in which the first step is not rate limiting.
ESSENTIAL KNOWLEDGE
5.9.A.1 If the first elementary reaction is not rate limiting, approximations (such as pre-equilibrium) must be made to determine a rate law expression.
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Required Course Content
LEARNING OBJECTIVE
5.10.A Represent the activation energy and overall energy change in a multistep reaction with a reaction energy profile.
ESSENTIAL KNOWLEDGE
5.10.A.1 Knowledge of the energetics of each elementary reaction in a mechanism allows for the construction of an energy profile for a multistep reaction.
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Required Course Content
LEARNING OBJECTIVE
5.11.A Explain the relationship between the effect of a catalyst on a reaction and changes in the reaction mechanism.
ESSENTIAL KNOWLEDGE
5.11.A.1 In order for a catalyst to increase the rate of a reaction, the addition of the catalyst must increase the number of effective collisions and/or provide a reaction path with a lower activation energy relative to the original reaction coordinate.
5.11.A.2 In a reaction mechanism containing a catalyst, the net concentration of the catalyst is constant. However, the catalyst will frequently be consumed in the rate-determining step of the reaction, only to be regenerated in a subsequent step in the mechanism.
5.11.A.3 Some catalysts accelerate a reaction by binding to the reactant(s). The reactants are either oriented more favorably or react with lower activation energy. There is often a new reaction intermediate in which the catalyst is bound to the reactant(s). Many enzymes function in this manner.
5.11.A.4 Some catalysts involve covalent bonding between the catalyst and the reactant(s). An example is acid-base catalysis, in which a reactant or intermediate either gains or loses a proton. This introduces a new reaction intermediate and new elementary reactions involving that intermediate.
5.11.A.5 In surface catalysis, a reactant or intermediate binds to, or forms a covalent bond with, the surface. This introduces elementary reactions involving these new bound reaction intermediate(s).

UNIT 6 Thermochemistry

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Guiding question

What happens when protons are transferred?

Learning outcomes
After studying this topic you should be able to:

Understand

Apply your knowledge to

  • a Brønsted-Lowry acid is a proton donor and a Brønsted-Lowry base is a proton acceptor.

  • a pair of species differing by a single proton is called a conjugate acid-base pair.

  • some species can act as both Brønsted-Lowry acids and bases.

  • the pH scale is a convenient means to describe the [H+] of a solution;

pH = − log10[H+]; [H+] = 10−pH.

  • the ionic product constant of water, Kw, shows an inverse relationship between [H+] and [OH−]; Kw = [H+][OH−].

  • strong and weak acids and bases differ in the extent of ionization.

  • acids react with bases in neutralization reactions.

  • pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features.

  • deduce the Brønsted-Lowry acid and base in a reaction.

  • deduce the formula of the conjugate acid or base of any Brønsted-Lowry base or acid.

  • interpret and formulate equations to show acid-base reactions of these species.

  • perform calculations involving the logarithmic relationship between pH and [H+].

  • recognise solutions as acidic, neutral and basic by the relative values of [H+] and [OH−]

  • recognise that acid-base equilibria lie in the direction of the weaker conjugate

  • formulate equations for the reaction between acids and metal oxides, metal hydroxides, hydrogencarbonates and carbonates.

  • sketch and interpret the general shape of the pH curve for the neutralisation reaction between a strong acid and a strong base

AHL

AHL

  • the pOH scale describes the [OH−] of a solution. pOH = − log10[OH−]; [OH−] = 10−pOH

  • the strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values.

  • for a conjugate acid-base pair, the relationship Ka x Kb = Kw can be derived from the expressions for Ka and Kb.

  • the pH of a salt solution depends on the relative strengths of the parent acid and base.

  • pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features.

  • acid-base indicators are weak acids, where the components of the conjugate acid-base pair have different colours. The pH of the end point of an indicator, where it changes colour, approximately corresponds to its pKa value.

  • an appropriate indicator for an acid-base titration has an end point range that coincides with the pH at the equivalence point.

  • a buffer solution is one which resists change in pH on the addition of small amounts of acid or alkali.

  • the pH of a buffer solution depends on both:
    → the p    Ka or pKb of its acid or base.
    → the ratio of the concentration of acid or base to the concentration of the conjugate base or acid.

  • interconvert between [H+], [OH−], pH and pOH values.

  • interpret the relative strengths of acids and bases from Ka, Kb, pKa and pKb data.

  • solve problems involving Ka, Kb, Kw, pKa and pKb values

  • construct equations for the hydrolysis of ions in a salt, and the predict the effect of each ion on the pH of the salt solution.

  • interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases.

  • construct equilibria expressions to show why the colour of an indicator changes with pH.

  • identify an appropriate indicator for a titration from the identity of the salt and the pH range of the indicator.

  • describe the composition of acidic and basic buffers and explain their action.

  • solve problems involving the composition and pH of a buffer solution, using the equilibrium constant.

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Required Course Content
LEARNING OBJECTIVE
6.2.A Represent a chemical or physical transformation with an energy diagram.
ESSENTIAL KNOWLEDGE
6.2.A.1 A physical or chemical process can be described with an energy diagram that shows the endothermic or exothermic nature of that process.
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Required Course Content
LEARNING OBJECTIVE
6.3.A Explain the relationship between the transfer of thermal energy and molecular collisions.
ESSENTIAL KNOWLEDGE
6.3.A.1 The particles in a warmer body have a greater average kinetic energy than those in a cooler body.
6.3.A.2 Collisions between particles in thermal contact can result in the transfer of energy. This process is called “heat transfer,” “heat exchange,” or “transfer of energy as heat.”
6.3.A.3 Eventually, thermal equilibrium is reached as the particles continue to collide. At thermal equilibrium, the average kinetic energy of both bodies is the same, and hence, their temperatures are the same.
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Required Course Content
LEARNING OBJECTIVE
6.4.A Calculate the heat q absorbed or released by a system undergoing heating/cooling based on the amount of the substance, the heat capacity, and the change in temperature.
ESSENTIAL KNOWLEDGE
6.4.A.1 The heating of a cool body by a warmer body is an important form of energy transfer between two systems. The amount of heat transferred between two bodies may be quantified by the heat transfer equation:
EQN: q = mcΔT.
Calorimetry experiments are used to measure the transfer of heat.
6.4.A.2 The first law of thermodynamics states that energy is conserved in chemical and physical processes.
6.4.A.3 The transfer of a given amount of thermal energy will not produce the same temperature change in equal masses of matter with differing specific heat capacities.
6.4.A.4 Heating a system increases the energy of the system, while cooling a system decreases the energy of the system.
6.4.A.5 The specific heat capacity of a substance and the molar heat capacity are both used in energy calculations.
6.4.A.6 Chemical systems change their energy through three main processes: heating/cooling, phase transitions, and chemical reactions.
6.4.A.7 In calorimetry experiments involving dissolution, temperature changes of the mixture within the calorimeter can be used to determine the direction of energy flow. If the temperature of the mixture increases, thermal energy is released by the dissolution process (exothermic). If the temperature of the mixture decreases, thermal energy is absorbed by the dissolution process (endothermic).
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Required Course Content
LEARNING OBJECTIVE
6.5.A Explain changes in the heat q absorbed or released by a system undergoing a phase transition based on the amount of the substance in moles and the molar enthalpy of the phase transition.

ESSENTIAL KNOWLEDGE

6.5.A.1 Energy must be transferred to a system to cause a substance to melt (or boil). The energy of the system therefore increases as the system undergoes a solid-to-liquid (or liquid-to-gas) phase transition. Likewise, a system releases energy when it freezes (or condenses). The energy of the system decreases as the system undergoes a liquid-to-solid (or gas-to-liquid) phase transition. The temperature of a pure substance remains constant during a phase change.

6.5.A.2 The energy absorbed during a phase change is equal to the energy released during a complementary phase change in the opposite direction. For example, the molar enthalpy of condensation of a substance is equal to the negative of its molar enthalpy of vaporization. Similarly, the molar enthalpy of fusion can be used to calculate the energy absorbed when melting a substance and the energy released when freezing a substance.

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Required Course Content
LEARNING OBJECTIVE
6.6.A Calculate the heat q absorbed or released by a system undergoing a chemical reaction in relationship to the amount of the reacting substance in moles and the molar enthalpy of reaction.
ESSENTIAL KNOWLEDGE

6.6.A.1 The enthalpy change of a reaction gives the amount of heat energy released (for negative values) or absorbed (for positive values) by a chemical reaction at constant pressure.

6.6.A.2 When the products of a reaction are at a different temperature than their surroundings, they exchange energy with the surroundings to reach thermal equilibrium. Thermal energy is transferred to the surroundings as the reactants convert to products in an exothermic reaction. Thermal energy is transferred from the surroundings as the reactants convert to products in an endothermic reaction.

6.6.A.3 The chemical potential energy of the products of a reaction is different from that of the reactants because of the breaking and forming of bonds. The energy difference results in a change in the kinetic energy of the particles, which manifests as a temperature change.

Exclusion Statement: The technical distinctions between enthalpy and internal energy will not be assessed on the AP Exam. Most reactions studied at the AP level are carried out at constant pressure, where the enthalpy change of the process is equal to the heat (and by extension, the energy) of reaction.
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Required Course Content
LEARNING OBJECTIVE
6.7.A Calculate the enthalpy change of a reaction based on the average bond energies of bonds broken and formed in the reaction.
ESSENTIAL KNOWLEDGE
6.7.A.1 During a chemical reaction, bonds are broken and/or formed, and these events change the potential energy of the system.
6.7.A.2 The average energy required to break all of the bonds in the reactant molecules can be estimated by adding up the average bond energies of all the bonds in the reactant molecules. Likewise, the average energy released in forming the bonds in the product molecules can be estimated. If the energy released is greater than the energy required, the reaction is exothermic. If the energy required is greater than the energy released, the reaction is endothermic.
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Required Course Content
LEARNING OBJECTIVE
6.8.A Calculate the enthalpy change for a chemical or physical process based on the standard enthalpies of formation.
 
ESSENTIAL KNOWLEDGE
6.8.A.1 Tables of standard enthalpies of formation can be used to calculate the standard enthalpies of reactions.
EQN: ΔH°reaction = ΣΔHf°products − ΣΔHf°reactants
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Required Course Content
 LEARNING OBJECTIVE
6.9.A Represent a chemical or physical process as a sequence of steps.
6.9.B Explain the relationship between the enthalpy of a chemical or physical process and the sum of the enthalpies of the individual steps.
ESSENTIAL KNOWLEDGE

6.9.A.1 Many processes can be broken down into a series of steps. Each step in the series has its own energy change.

6.9.B.1  Because total energy is conserved (first law of thermodynamics), and each individual reaction in a sequence transfers thermal energy to or from the surroundings, the net thermal energy transferred in the sequence will be equal to the sum of the thermal energy transfers in each of the steps. These thermal energy transfers are the result of potential energy changes among the species in the reaction sequence; thus, at constant pressure, the enthalpy change of the overall process is equal to the sum of the enthalpy changes of the individual steps.

6.9.B.2 The following are essential principles of Hess’s law:

i. When a reaction is reversed, the enthalpy change stays constant in magnitude but becomes reversed in mathematical sign.

ii. When a reaction is multiplied by a factor c, the enthalpy change is multiplied by the same factor c.
iii. When two (or more) reactions are added to obtain an overall reaction, the individual enthalpy changes of each reaction are added to obtain the net enthalpy change of the overall reaction.
Exclusion Statement: The concept of state functions will not be assessed on the AP Exam.

UNIT 7 Equilibrium

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