Chemistry: Enthalpy and Formation Reactions

Chapter 6: Thermochemistry

• Thermochemistry is the study of heat changes in chemical reactions.
• If a reaction can be written as the sum of two or more reactions, the enthalpy change (ΔH°) for the overall reaction is the sum of the ΔH° values of the contributing reactions (Hess's Law).
• Enthalpy of formation (ΔH_f) is the heat change when one mole of a compound is formed from its constituent elements in their standard states.
• The standard enthalpy of formation of the most stable form of an element is zero (ΔH°_f = 0)
• Standard enthalpy of reaction, ΔH°_rxn = Σn ΔH°_f(products) - Σm ΔH°_f(reactants) (n & m are stoichiometric coefficients)

Practice: Heating Curves

• Calculate the amount of energy needed to convert 237 g of ice at 0.0°C to water at 80.0°C, using given values for ΔH_fusion, ice and specific heat capacity of water.

Practice: Hess's Law

• Calculate the reaction enthalpy for the formation of methane (CH₄) from coal and hydrogen gas (is this an exothermic or endothermic process).
• Use the given reactions, and known enthalpy changes for the reactions in part 1.

Practice: Hess's Law

• Determine ΔH° for the following reaction, 2 NH₃(g) + 5/2 O₂(g) → 2 NO(g) + 3 H₂O(g).

Practice: Hess's Law

• Calculate the enthalpy of formation of CCl₄ (g) using the thermochemical equations provided.

Enthalpy of Formation

• Enthalpy of formation, ΔH_f, is the enthalpy change when one mole of a substance is formed from its constituent elements, all in their standard states. A standard state is the most stable form of a substance at a given temperature (typically 25°C or 298.15K)

Standard Enthalpy of Formation

• The standard enthalpy of formation of the most stable form of an element is zero (ΔH°_f = 0).

Standard Enthalpy of Reaction

• The enthalpy change of a reaction when all reactants and products are in their standard states (ΔH°_rxn) can be calculated from the standard enthalpies of formation for the reactants and products.

Estimating Enthalpy Changes

• Breaking chemical bonds requires energy, while forming bonds releases energy.

Determining Bond Enthalpy

• Bond enthalpy depends on what elements are in a bond, the type of bond between the atoms, and the polarity of the rest of the molecule
• Bond enthalpy values are averages over a variety of molecules.

Chapter 7: Quantum Chemistry

• Electrons are the ultimate determinant of atomic properties and reactivity.
• Understanding how electrons behave and what influences them is crucial for understanding chemistry.
• A model of the atom explaining electronic structure, electronic energy, and electronic transitions.

The Electromagnetic Spectrum

• Radiant energy moves through space as perpendicular oscillating electric and magnetic waves.
• The electromagnetic spectrum includes radio waves, microwaves, infrared, visible light, ultraviolet, X-rays, and gamma rays. Visible light is part of the spectrum.

Wave Properties

• Wavelength (λ): the distance between two corresponding points on a wave (crest to crest or trough to trough). Units are often metres or nanometers
• Frequency (ν): the number of complete waves passing a point per unit time. Typical units are Hz (s⁻¹) or cycles per second
• Amplitude (A): the height of a wave, from the midpoint to the peak or trough; measure of energy in a wave.

Electromagnetic Radiation Values

• Different regions of the electromagnetic spectrum have different wavelength ranges.
• Wavelength determines color for the visible light range.

Wave Properties and Light

• Wavelength and frequency are related by: λν = c (c is the speed of light).
• Higher frequency correlates to shorter wavelengths.
• Lower frequency correlates to longer wavelengths.

Conversions for Energy, Wavelength, and Frequency

• CD players use lasers to emit red light with a wavelength of 685 nm; calculate frequency.

Photoelectric Effect

• Irradiating matter with light can cause electrons to be ejected from a material

Kinetic Energy of EM Radiation

• Light shining on metals ejects electrons; light is a carrier of energy.

Atomic View of Electron Removal

• Energy removal is required to overcome the attraction between the negative electron and positive nucleus.

The Photoelectric Effect

• For low-frequency light, no electrons are emitted regardless of the intensity of light
• Electrons are emitted by given metal only at or above a specific threshold frequency (ν₀).
• The number of emitted electrons increases with light intensity for light frequencies above ν₀; Kinetic Energy (KE) of electrons increases linearly with the frequency of the light.

Photon Energy

• Individual light packets are photons with energy determined by frequency. The equation for calculating the energy of a photon is: E = hv or E = hc/λ (h = Planck's constant).

Wave-Particle Duality

• Light exhibits both wave-like and particle-like behavior. The photoelectric effect demonstrates the particle-like property of light.

Chapter 7: Quantum Chemistry

• Discuss the history of atomic models; describe the different model versions of the atom by scientists during the era of the late 1800s to early 1900s (e.g. Thomson's Plum Pudding Model) and compare those models.
• Explain the relationship between the energy change of a single electron and the emitted photon
• Explain how Rydberg developed a constant from the Planck's constant, speed of light, & Rydberg constant; equation for calculating the energy of a photon (E) from its wavelength.

Bohr and the Hydrogen Atom

• Electrons must exist only at specific energy levels based on their distance from the nucleus, and these are quantized/discrete values, as opposed to a continuous value.
• Energy can be absorbed or emitted to cause electrons to move between energy levels (an electron absorbing energy becomes electronically excited; emitted energy causes the electron to drop to an energy level).

Electronic States

• Energy level refers to the potential energy of an electron/shell within an atom.
• Ground state is an atom or electron in its lowest energy level.
• Excited state is an atom or electron at a higher energy level than its ground state.
• Principle quantum number (n) represents a specific atomic shell, is associated with a whole number.

The Bohr Model and Emission Spectra

• Explain how energy level changes in electrons correlate to different wavelengths of light emitted.

Poll Question: Predicting Photon Emission

• Which electron transition would emit a photon with the highest frequency.

Predicting Color of Light Emitted

• Equation for predicting the color of light emitted when an electron transitions between energy levels in an atom

Practice: The Rydberg Equation

• Calculate the wavelength of a line within the visible spectrum for a given electronic transition in a Hydrogen atom.

Emission Spectra of Other Elements

• Emission spectra provided by excited gaseous elements are unique to each element.

The Bohr Model Evaluation

• Strengths of the Bohr atomic model.
• Limitations of the Bohr atomic model.

Are Electrons Particles or Waves?

• De Broglie proposed electrons could also exhibit wave-like behavior. This is different than the predictable particle nature of the Bohr model.
• Combine equations for kinetic energy with equations for waves (λ = h/mv for the relation between wavelength (λ), Planck's constant(h), mass(m), and velocity(v)) to quantify the De Broglie wavelength of electrons.

Electrons as Waves?

• Discuss the result of wave-like properties of electrons on Bohr's model.

Waves and Nodes

• Explain standing waves as a good model for wave behavior associated with electrons.
• Define node as a point where amplitude = 0 to quantify wave properties.

Electrons as Waves

• Electrons behave like circular waves oscillating around the nucleus. The stable circular waves can only have a circumference defined to a specific value (nλ)

The Uncertainty Principle

• Discuss similarities and differences between wave theory and particle theory; neither describe electron behavior accurately within the same atom/situation.
• Heisenberg Uncertainty Principle - It is impossible to know both position and velocity accurately at the same time of a quantum particle such as an electron because it would require knowing the precise momentum.

Schrödinger Equation

• The Schrödinger equation accounts for more properties of the electron including energy due to motion, distance from nucleus, and wave-like behavior of the electron.

Plotting Probability Clouds

• Wave function is the result of solving Schrödinger's equation.
• Probability function is plotted as a probability cloud to reflect the probability of finding an electron within a given space (or 90% probability within a specified range).

Quantum Model of Electron Behavior

• Schrödinger's model is a refinement/expansion of Bohr's model; Schrödinger's model incorporates energy quantization for electrons in an atom in relationship to wave-behavior.
• Quantization of energy is a result of electron wave behavior.
• Electron energy is known but the exact position of an electron is not possible, as determined by the Heisenberg uncertainty principle.
• Electrons are most probably found in orbitals, which are determined by the Schrödinger equation.
• Wave functions' energies depend on the values for the three variables: n, l and m.

Quantum Numbers

• To solve Schrodinger's equations, we introduce three quantum numbers (n, l, and m). These describe probable electron location (orbital)
• Pauli Exclusion Principle - No two electrons in the same atom can have the same four quantum numbers.

Quantum Number Variables: Summary

• Principle quantum number (n): related to energy level. Possible values are positive integers.
• Angular momentum quantum number (l): related to the shape of the orbital and its spatial orientation. Possible values depend on n (range 0 to n-1).
• Magnetic quantum number (m_l): related to the orientation of the orbital in space. Possible values are integers from -l to +l.

Angular Momentum Quantum Number

• The angular momentum quantum number (l) describes the shape/shape and spatial orientation of the electron's orbital in space. Possible quantum numbers are integers from 0 to n-1. The possible shapes are s, p, d and f orbitals.

s Orbitals

• s orbitals are spheres.
• As the principal quantum number (n) increases, the s orbital becomes larger, increasing its energy level and the furthest distance from the nucleus.

The Three p Orbitals

• p orbitals (l = 1) can be present in energy levels n=2 and higher than 2.
• Unlike s orbitals, p orbitals are not spheres. This means they need to have more orbitals to cover all space around the nucleus.

The Five d Orbitals

• d orbitals (l=2) exist in energy levels 3 or higher.
• More complex than p orbitals, but analogous in concept.
• Five orientations/shapes are possible (m_l = -2, -1, 0, +1, +2).

f Orbitals (l = 3)

• f orbitals (l=3): seven orientations, complex.

Magnetic Quantum Number

• The magnetic quantum number (m_l) describes the orientation of an orbital in space. Possible values range from -l to +l (including 0).

Poll Question: Quantum Numbers

• Which quantum number set is valid for describing a 5d orbital?

Spin Magnetic Quantum Number

• Not all spectral features of atoms and ions can be fully explained by just the wave equations of atoms and ions. Further, some unique atomic/ion properties were observed but could not be explained by wave equations (e.g. "doublets")
• Explains the "doublets" as a result of the spinning of the electron that creates a magnetic field oriented "up" or "down."
• The fourth quantum number (m_s) describes the electron spin. m_s can be either +½ or -½.

The Four Quantum Numbers

• The Four Quantum Numbers characterize a unique electron within an atom for its energy, shape, and spatial orientation within an atom.

Orbital Energies

• In a single electron/atom, orbitals at the same energy level that have the same principal quantum numbers (n) are degenerate.
• As more electrons are added to multi-electron atoms, orbitals in the same energy level are no longer degenerate due to electron-electron repulsion.

Rules for Electrons Filling Orbitals

• Each principal energy level has one or more orbitals with different energies based on shape.
• Atoms need to be sequentially filled from lowest to highest energy orbitals.

Electrons in Orbitals

• Each orbital can hold up to two electrons.
• Electrons have opposite spins within an orbital The Pauli exclusion principle states that no two electrons in an atom may have the same four quantum numbers, thus each orbital can only hold two electrons.

Pauli Exclusion Principle: Two Electrons Per Orbital

• No two electrons in the same atom can have exactly the same energy; thus, no two electrons within an atom can have the same four quantum numbers.

Aufbau Principle: Lowest to Highest Energy

• Electrons enter the lowest-energy orbitals first when filling orbitals sequentially with electrons.
• Filling of orbitals follow the sequence s → p → d → f. The numerical values of the principle quantum number (n) increase according to the sequence of orbital filling.

Energies of Orbital Types

• The sequence of electron filling from lowest to highest energy level.

Electron Filling From the Periodic Table

• Electron filling follows a specific order determined by the periodic table

Hund's Rule: Degenerate Orbitals

• Hund's rule states that when filling orbitals of equal energy, electrons first occupy each orbital singly with parallel spins before pairing up.
• Orbitals of equivalent energy are termed degenerate. This applies to p, d and f orbitals. Thus, half-filled/filled degenerate orbitals are more stable than other configurations.

Writing Electron Configurations

• Electron configuration notation is used to represent the arrangement of electrons within an atom or an ion, using spdf notation and noble gas abbreviation.
• Orbital box notation follows Hund's rule and utilizes diagrams to visualize the assignment of electrons in orbitals based on filling rules.

Orbital Box Diagrams

• Draw orbital box diagrams for given elements.

Poll Question: Electron Configuration

• What is the election configuration of potassium (Z=19). What is the electron configuration abbreviation that utilizes noble gas configuration?

Electron Configuration Abbreviations

• Use of noble gas abbreviations in writing electron configuration

Noble Gas Configuration and Valence Electrons

• Completely filled shells (valence levels) of electrons are especially stable.
• Core electrons are part of completely filled shells.
• Valence electrons are electrons beyond the filled shells, and participate in chemical reactions.

Practice: Transition Metal Configurations

• Determine the electron configurations of iron (Fe) and cobalt (Co) using noble gas abbreviation.

Electron Configuration Abbreviations

• Box diagrams and electron configurations: Unusual electron configurations (e.g. chromium and copper).

Electron Configuration Exceptions

• In certain instances, some elements will place electrons into orbitals in a manner that breaks the Aufbau Principle, in order to achieve more stable half-filled/filled sets of orbitals (d orbitals).

F-Block Elements

• Explain the elements of the F-block including lanthanides and actinides.

Poll Question: Electron Configuration for F-Block

• What is the noble gas abbreviation configuration for a specific f-block element?

Paramagnetism and Diamagnetism

• Paramagnetism - A property of a substance in which unpaired electrons cause the substance to align with a magnetic field.
• Diamagnetism - A property of a substance in which paired electrons cause the substance to repel or weakly oppose a magnetic field.

Valence Electrons & Chemistry

• Valence electron configurations have a correlation with chemical properties

Chapter 7: Periodic Trends

• Explain the periodic trends in atomic radius, ionization energy, electron affinity, and electronegativity.
• Ionic radii: The radius of an ion. For cations the size of the ion is smaller than the atom from which is originates, and for anions the size of the ion is larger than the originating atom.
• The periodic trends in chemical properties: Electron distributions in metals and nonmetals & periodic trends in reactivity of groups

Atomic Radius

• Atomic radius: The size of an atom.
• Atomic radius increases down a group and decreases across a period. This is due to an increase in electrons in each shell down a group, and increasing nuclear charge across the period (apparent nuclear charge).

Atomic Radius and Apparent Nuclear Charge

• Explain how apparent nuclear charge affects atomic size.

General Trends of Atoms

• Describe the general trends in atomic properties: Number of shells and shielding.

Poll Question: Atomic Radius

• Determine the order of atoms with increasing atomic radius.

Ionic Radius: Cations

• Cation - Atoms that lose electrons, and have a smaller ionic radius than the originating atom
• Ion size decreases going across a period, and increases going down the group. (More valence electrons = larger radius)

Ionic Size: Cations

• Explain the factors affecting cation size.

Ionic Radius: Anions

• Anion - Atoms that gain electrons
• Ion size decreases going across a period, and increases going down the group. (More valence electrons = larger radius)

Ionic Radius: Anions

• Discuss the factors affecting anion size.

Trends in Ionic Size

• Illustrate the trend in ionic sizes of atoms and ions within the periodic table.

Poll Question: Ionic Radius

• Which ion has the largest ionic radius in a list of various ions?

Ionization Energy

• Ionization energy: The minimum energy required to remove an electron from an atom.
• When an electron is removed from an atom, energy is required, thus ionization energy is always positive

Ionization Energy Example

• Explain why ionization energy increases as you move across a period, starting from the low side of the period and moving toward the nonmetal side of the period; describe the result.
• Explain why ionization energy decrease as you move down a group, starting from top to bottom; describe the result.
• Explain differences between the first, second, and third ionization energies among similar atoms (e.g. Mg)

Periodic Trends in Ionization Energy

• Describe the trends in ionization energy across a period and down a group

Periodic Trends in Ionization Energy

• Explain the trend and factors that cause the increase/decrease in ionization energy when moving up/down or across a periodic table row respectively.

Poll Question: Ionization Energy

• Which element would have the lowest ionization energy on the periodic table (in a given group or periodic table row).

Electron Affinity

• Electron Affinity: The energy change of an atom when gaining an electron.
• Electron affinity is negative when energy is released upon the addition of an electron on the atom

Periodic Trends: Electron Affinity

• Describe the trends in electron affinity across a period and down a group.

Electronegativity

• Electronegativity: A measure of the ability of an atom in a molecule to attract shared electrons to itself.
• Electronegativity scales atoms relative to one another.

Electronegativity

• Discuss the periodic trends of electronegativity.

Bond Types

• Describe the relationship between electronegativity difference and bond types.
• Pure covalent (or nonpolar covalent) bond - the difference in electronegativity of the bonding atoms is zero
• Polar covalent bond -The difference in electronegativities of the bonding atoms is between zero and 2 (a medium difference)
• Ionic bond –The difference in electronegativities of the bonding atoms is greater than 2, and electrons are transferred between atoms.

Polar Covalent Bond

• Discuss the relationship between electronegativity difference and bond polarity

Poll Question: Bond Polarity

• Identify the most polar bond in a list of various bonds.

Dipole Moments

• Dipole moment quantifies the polarity of a molecule.

Polarity and Molecular Shapes

• Polarity results from electronegativity differences on bonded atoms and their arrangement in space, following the VSEPR theory rules; ALL valence atoms in the molecule must be considered.

Practice: Determining Polarity

• Determine if a molecule is polar or nonpolar.

Polar Molecules

• Table to illustrate the trends and values for various polar molecules

Poll Question: Polar or Nonpolar?

• Select the polar molecule based on the chemical elements for the molecule and its geometry.

Practice: VSEPR and Bond Angle Sheet

• Practice various VSEPR (Valence Shell Electron Pair Repulsion) problems. These problems typically involve determining molecular geometry/shapes and bond angles.

Molecules with Multiple Central Atoms

• Determine the molecular shape and bond angles among various molecules with multiple central atoms.

Practice: Large Molecule (Saccharine)

• Determine the molecular geometry of various large molecules (e.g. Saccharine)

Five Electron Groups

• Molecules displaying five electron pairs to the central atom.

Trigonal Bipyramidal Geometries

• Describe the molecular geometry of molecules with five electron pairs.

Axial and Equatorial Electrons

• Explain how to predict molecule shape for molecules having five electron pairs using the VSEPR theory rules in relation to axial and equatorial electron pairs.

Octahedral Geometries

• Describe the molecular geometry of molecules with six electron pairs.

Six Electron Groups

• Explain the various arrangements in molecular geometry when a molecule contains six electron pairs.

Poll Question: Molecular Geometry for Expanded Octets

• Predict the molecular shape/geometry and bond angles.

Chapter 8: Lewis Structures - Covalent Compounds

• Discuss the relationship between Lewis structures and covalent compounds.

Writing Lewis Structures: Covalent Molecules

• Discuss methods for determining Lewis structures for various covalent molecules and polyatomic ions.
  • Determine the central atom
  • Determine total number of valence electrons
  • Place single bonds between the central atom and the terminal atoms; subtract 2 electrons from the total for every bond formed.
  • Place remaining electron pairs on terminal atoms and the central atom as needed to complete octets
  • Consider multiple bonds from the central atom to terminal atoms as needed.

Example: Lewis Structures

• Draw the Lewis structure for NO₃⁻

Practice: Lewis Structures

• Draw Lewis structures for given compounds, including various different atoms and bond types.
• Example types of given compounds are Cl₂CO, SCI₂, and CINO₂

Exceptions to the Octet Rule

• List out reasons why some elements may have more than eight valence electrons, thus breaking the octet rule.

Exceptions to the Octet Rule

• Incomplete octets are possible for boron given proper Lewis structure drawing.
• Free radicals have an odd number of electrons, such as nitrogen containing molecules.

Poll Question: Lewis Structures

• Determine the number of double bonds in a molecule

Isoelectronic Species

• Define isoelectronic species: species having the same number of valence electrons and the same total number of electrons or both.

Resonance Structures

• Explain resonance structures: different connectivity of atoms but different arrangement of valence electrons within a molecule.
• Discuss the relationship among resonance structures within a given molecule.

Resonance in Benzene

• Describe the resonance structures of Benzene (a cyclic molecule).

Resonance in Nitrate Ion

• Draw resonance structures for the Nitrate ion (NO₃⁻).

Resonance in Sulfate Ion

• Draw resonance structures for the Sulfate ion (SO₄²⁻).

Resonance Structures

• Describe how to determine the most predominant resonance structure in a molecule. Use formal charges and electronegativity rules.

Formal Charge

• Formal charge is used to estimate best resonance structures (structure with lowest formal charges is the best).

Polyatomic Ion Formal Charges

• When applying formal charge calculations to polyatomic ions, the overall formal charge must be equal to the charge of the overall ion (for instance total formal charge will not be zero for an ion).

Electronegativity and Formal Charge

• Describe the relationship between electronegativity and formal charge.

Poll Question: Comparing Lewis Structures

• Determine which Lewis structure is most appropriate according to various criteria, including formal charge.

Practice: Resonance Structures

• Draw resonance structures given a molecule or ion structure.

Bond Properties

• Bond properties depend on the bond order, lengths, & strenghts

Bond Order

• Bond order is the number of bonding electron pairs shared by two covalently bonded atoms.

Bond Length

• Bond length is the distance between nuclei of two bonded atoms; bond length varies in relation to the number of bonds (e.g. single vs double vs triple bonds)

Bond Length

• Discuss the relationship between bond length and atomic size and the relationship among bond length, bond order, and bond strength.

Bond Energy

• Bond energy is the amount of energy required to break a bond in a gaseous molecule.
• Describe the relationship among bond energy & bond order & bond length

Poll Question: Bond Strength

• Identify the strongest covalent bond in a given list.

Using Bond Energy to Calculate ΔH

• Discuss a method to estimate the enthalpy change (ΔH) associated with a reaction using bond energies.

Practice: Reaction Enthalpy From Bond Energy

• Determine ΔH_rxn for a given chemical reaction.

Example: Reaction Enthalpy From Bond Energy

• Determine ΔH_rxn for the combustion of H₂C.

Practice: Reaction Enthalpy From Bond Energy

• Determine the enthalpy change (ΔH) of a given reaction.

Chapter 9: Molecular Geometry

• The distribution of atoms in space.
• Valence Shell Electron Pair Repulsion (VSEPR) theory
• Number of bonding and nonbonding pairs of electrons on central atom; geometric arrangement that effectively separates these pairs in space.

Thinking About Bonding and Molecular Structure

• Pose/answer various questions related to bonding and molecular structure.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• Explain VSEPR theory

VSEPR Theory

• Valence electrons, whether present in bond pairs or lone pairs, always try to maximize their distance apart from one another in space.

Linear and Trigonal Planar Geometries

• Using VSEPR theory, explain the linear and trigonal planar geometry for molecules

Linear and Trigonal Planar Geometry

• Geometry diagrams related to the linear and trigonal planar geometries.

Tetrahedral Geometries

• Describe the tetrahedral molecular geometry; provide example of molecules with tetrahedral geometry.

Four Electron Groups

• Illustrate four electron pair geometries: Tetrahedral, Trigonal Pyramidal, and Bent.

Effect of Lone Pairs

• Non-bonding electron pairs in space exert greater repulsive forces on adjacent electron pairs, thus distorting the molecular geometry and changing bonding angles from their ideal values

Poll Question: Molecular Geometry with Lone Pairs

• Predict the molecular geometry of various molecules with lone electron pairs (e.g. sulfur dichloride)

Practice: Effect of Multiple Bonds

• Use the VSEPR theory rules to determine molecular geometry given multiple bonds

Molecular Geometries

• Determine the molecular geometry of various molecules (e.g. SO₂, PBr₃).

Molecules with Multiple Central Atoms

• Draw molecular geometry diagrams for large molecules with more than one central atom.

Practice: Large Molecule (Saccharine)

• Draw the molecular geometry of the saccharine molecule.

Five Electron Groups

• Describe the molecular shapes of molecules having five bonding pairs and lone pairs.

Trigonal Bipyramidal Geometries

• Describe the various molecular geometries in relation to the VSEPR theory for molecules with five electron groups.

Axial and Equatorial Electrons

• Distinguish between axial and equatorial electrons in determining the structure and shape of molecules with five electron groups using VSEPR theory.

Octahedral Geometries

• Describe various octahedral molecular geometries of molecules with 6 electron pairs (e.g. SF6, BrF₅, XeF₄).

Six Electron Groups

• Classify the various cases in octahedral molecular geometries according to the VSEPR theory rules with different numbers of bonding groups and lone pairs of electrons

Polarity

Electronegativity and Bond Polarity

• Describe how electronegativity differences result in bond polarity.

Measuring Polarity

• Discuss the quantitative measurements of molecular polarity called dipole moments.

Electronegativity

Bond Types

• Describe the different types of bonds from pure (nonpolar) covalent to polar covalent to ionic bonds; relate the different types of bonds to the differences in electronegativity of the atoms within the molecule/bond.

Polar Covalent Bond

• Explain the properties of a polar covalent bond and how it differs from other types of chemical bonds (e.g. purely covalent, ionic).

Poll Question: Bond Polarity

• Select the most polar covalent bond in a list

Dipole Moments

• Explain the concept of dipole moments and their relationship to molecular polarity

Polarity and Molecular Shapes

• Explain the relationship between molecular polarity & molecular geometry & all valence electrons in determining if a molecule is polar or nonpolar.

Practice: Determining Polarity

• Determine if a molecule is polar or nonpolar based on its geometry and electronegativities.

Polar Molecules

• Table to indicate various factors that indicate whether a molecule is polar or not.

Poll Question: Polar or Nonpolar?

• Classify a molecule as polar or nonpolar based on given information about the elements and their arrangement

Practice: VSEPR and Bond Angle Sheet

• Various VSEPR practice problems.

Test your understanding of the concepts of standard enthalpy of formation and enthalpy changes in chemical reactions. This quiz covers key topics such as reaction equations, energy changes, and specific reaction enthalpies. Perfect for chemistry students who want to reinforce their knowledge in thermodynamics.