Chapter 1 Flashcards
What is the “Constitution” of a molecule?
*The order in which the atoms of a molecule are connected.
*This connectivity of the molecule is commonly referred to as simply the molecules “structure”.
What is the correct structure of Ozone, O3?
What is Condensed Structure Formulae?
Lewis structures in which many (or all) covalent bonds and electron pairs are omitted.
What is Bond-line formulas?
*Shows bonds between atoms (except those that go from C to H).
Shows atoms that are not carbon and hydrogen (heteroatoms)
How would this molecule look in bond-line Structure?
What do curved arrows mean?
Curved arrows are used to track the flow of electrons in chemical reactions.
Draw the curved arrow(s) representing the electron transfers.
Draw the curved arrow(s) representing the electron transfers.
Are most A-B reactions reversible?
Yes
What differentiates a Strong acid from a weak acid in strength?
Strong acid equilibrium favour products and weak acid equilibrium favour reactants.
The stronger the acid, the ______ the Ka. The stronger the acid, the ______ the pKa.
The stronger the acid, the larger the Ka. The stronger the acid, the smaller the pKa.
pH = ____
pH = -log [H+]
What are the most common type of Organic Acids?
Carboxylic Acids (with carboxyl functional group: C==O + C–OH (-COOH)
What is and atomic orbital?
*A three-dimensional region around the nucleus where an electron is most likely to be found.
*Because the Heisenberg uncertainty principle states that both the precise location and the exact momentum of an atomic particle cannot be simultaneously determined
*We can never say precisely where an electron is—we can only describe its probable location.
What’s the shape of an s orbital.
s atomic orbital is a sphere with the nucleus at its center
Is an 2s orbital closer or further to the nucleus than an 1s orbital?
further
What are the two types of properties of electrons? (“behavioural”)
Electrons have particle-like and wave-like properties.
What is a node? (of a 2s orbital)
*2s orbital has a region where the probability of finding an electron falls to zero.
*Nodes occur because electrons have both particle-like and wave-like properties. A node is a consequence of the wave-like properties of an electron.
What’s the shape of an p orbital.
*p orbitals have two lobes.
*Generally, the lobes are depicted as teardrop shaped, but computer-generated representations reveal that they are shaped more like doorknobs
*(Notice that in this context, + and - indicate the phase of the orbital; they do not indicate charge.)
Where is the node of a p orbital?
The node of the p orbital is a plane—called a nodal plane—that passes through the center of the nucleus, between its two lobes.
How many degenerated p-orbital is there?
*Three
*Each corresponding lobe is perpendicular to another
*(This picture shows the three degenerated 2p orbitals
What is different from 1p orbitals to 2p orbitals? What about the energies?
*The energy of a 2p orbital is slightly greater than that of a 2s orbital because the average location of an electron in a 2p orbital is farther away from the nucleus.
*Drawing would be bigger tear drops than 1p orbital
Draw a 3s orbital.
What is the Molecular Orbital Theory?
*Like an atomic orbital, which describes the volume of space around an atom’s nucleus where an electron is likely to be found, a molecular orbital describes the volume of space around a molecule where an electron is likely to be found.
*And like atomic orbitals, molecular orbitals, too, have specific sizes, shapes, and energies.
*An atomic orbital surrounds an atom.
A molecular orbital surrounds a molecule.
Explain the formation of a sigma bond.
Is energy released or absorbed?
*Imagine a meeting of two separate H atoms. As the 1s atomic orbital of one hydrogen atom approaches the 1s atomic orbital of the other hydrogen atom, the orbitals begin to overlap. The atoms continue to move closer, and the amount of overlap increases until the orbitals combine to form a molecular orbital. The covalent bond that is formed when the two s orbitals overlap is called a sigma (S) bond. A S bond is cylindrically symmetrical—the electrons in the bond are symmetrically distributed about an imaginary line connecting the nuclei of the two atoms joined by the bond.
*As the two orbitals begin to overlap, energy is released because the electron in each atom is attracted to its own nucleus and to the nucleus of the other atom
In the formation of bonds, the decreased and increased energy means…
*The attraction between the electrons of and atom to the nuclei of another causes a decrease of energy until the nuclei are so close that they start to repulse one another.
*The repulsion of the two approaching nuclei causes a large increase in energy. The minimum energy (maximum stability) is achieved when the nuclei are a particular distance apart. This distance is the bond length of the new covalent bond
*aka: Minimum energy corresponds to maximum stability.
*At the end energy is released when a covalent bond forms and breaking the bond requires precisely the same amount of energy its used to be formed
What is “bond dissociation energy”
*a measure of bond strength—is the
energy required to break a bond or the energy released when a bond is formed.
*Every covalent bond has a characteristic bond length and bond dissociation energy.
What is the conservation principle related to orbitals in molecular formations?
Orbitals are conserved, meaning the number of molecular orbitals formed must equal the number of atomic orbitals combined.
In the context of the formation of an H ¬ H bond, how many molecular orbitals are discussed when two atomic orbitals are combined?
When describing the formation of an H ¬ H bond, two atomic orbitals are combined, but only one molecular orbital is discussed.
Even though only one molecular orbital is discussed in the formation of an H ¬ H bond, where is the other molecular orbital, and why is it not discussed?
The other molecular orbital is present, but it doesn’t contain any electrons. It exists due to the wave-like properties of electrons.
What property of electrons leads to the formation of two molecular orbitals when two atomic orbitals combine?
The wave-like properties of electrons cause two atomic orbitals to form two molecular orbitals.
How can two atomic orbitals combine in the context of bonding molecular orbitals, and what is the term used to describe the constructive combination of two s atomic orbitals?
Two atomic orbitals can combine in an additive (constructive) manner, similar to how two light waves or sound waves can reinforce each other. The constructive combination of two s atomic orbitals is called a s (sigma) bonding molecular orbital.
How can two atomic orbitals combine in a way that cancels each other, and what is this cancellation similar to in the context of light waves and sound waves?
Two atomic orbitals can combine in a destructive way, canceling each other. This cancellation is similar to the darkness resulting from the cancellation of two light waves or the silence resulting from the cancellation of two sound waves.
What is the term used to describe the destructive combination of two s atomic orbitals, and how is an antibonding orbital indicated in notation?
The destructive combination of two s atomic orbitals is called a σ* antibonding molecular orbital. An antibonding orbital is indicated by an asterisk (), and chemists read it as “star.” Therefore, σ is read as “sigma star.”
How are the σ bonding molecular orbital and the σ antibonding molecular orbital represented in a molecular orbital (MO) diagram, and what do the horizontal lines in an MO diagram signify?*
The σ bonding molecular orbital and the σ* antibonding molecular orbital are shown in the molecular orbital (MO) diagram. In an MO diagram, the energies of both the atomic orbitals and the molecular orbitals are represented as horizontal lines. The bottom line represents the lowest energy level, and the top line represents the highest energy level.
When two atomic orbitals overlap, how many molecular orbitals are formed, and what is the relative energy relationship between these molecular orbitals and the atomic orbitals?
When two atomic orbitals overlap, two molecular orbitals are formed—one lower in energy and one higher in energy than the atomic orbitals. The energies of these molecular orbitals are represented in relation to the energies of the atomic orbitals in the MO diagram.
Where are electrons in the bonding molecular orbital most likely to be found, and what is the significance of this location in terms of atomic nuclei attraction?
Electrons in the bonding molecular orbital are most likely to be found between the nuclei, where they can more easily attract both nuclei simultaneously. This increased electron density between the nuclei is what binds the atoms together.
Why are electrons in the antibonding molecular orbital likely to be found anywhere except between the nuclei, and what impact does this have on the exposure of positively charged nuclei to each other?
Electrons in the antibonding molecular orbital are likely to be found anywhere except between the nuclei because a nodal plane lies between the nuclei. This leaves the positively charged nuclei more exposed to each other, detracting from the formation of a bond.
How does the energy level of the bonding molecular orbital compare to that of the antibonding molecular orbital, and what is the relationship between electron density and stability in molecular orbitals?
The bonding molecular orbital is lower in energy and more stable than the antibonding molecular orbital. Higher electron density between the nuclei contributes to greater stability.
How are electrons assigned to molecular orbitals, and what principles govern their occupation in these orbitals?
Electrons are assigned to molecular orbitals using the same rules as those used for assigning electrons to atomic orbitals:
*Electrons always occupy available orbitals with the lowest energy (the aufbau principle).
*No more than two electrons can occupy a molecular orbital (the Pauli exclusion principle).
According to the molecular orbital diagram, why is H2+ predicted to be less stable than H2, and why is He2 predicted not to exist?
H2+ is predicted to be less stable than H2 because it has only one electron in the bonding molecular orbital. He2 is predicted not to exist because the four electrons of He2 would fill both the lower energy bonding MO and the higher energy antibonding MO, canceling the advantage gained by the electrons in the bonding MO.
What type of bond is formed when two p atomic orbitals overlap, and how is this overlap characterized in terms of the orientation of the orbitals?
When two p atomic orbitals overlap, a pi (π) bond is formed. This overlap is characterized by the side of one orbital overlapping the side of the other.
What is constructive overlap, and how does it relate to the formation of a pi bonding molecular orbital in the context of p orbitals?
Constructive overlap occurs when two in-phase p atomic orbitals overlap (e.g., blue lobes overlap with blue lobes, and green lobes overlap with green lobes). This overlap forms a pi (π) bonding molecular orbital.
What is destructive overlap, and how does it contribute to the formation of a π* (“pi star”) antibonding molecular orbital in the context of p orbitals?*
Destructive overlap occurs when two out-of-phase p atomic orbitals overlap (e.g., blue lobes overlap with green lobes). This overlap forms a π* (pi star) antibonding molecular orbital.
How many nodal planes does the pi bonding molecular orbital have, and what is the significance of these nodal planes in its structure?
The pi bonding molecular orbital has one nodal plane—a nodal plane that passes through both nuclei.
What does the Valence-Shell Electron-Pair Repulsion (VSEPR) model aim to predict, and what is the underlying principle of this model in terms of electron density around an atom?
*The VSEPR model predicts molecular geometry based on minimizing electron repulsion between regions of electron density around an atom.
*The model is founded on the principle that atoms share electrons by overlapping their atomic orbitals, and, due to electron pair repulsion, bonding and lone-pair electrons around an atom are positioned as far apart as possible.
How do Lewis structures and the VSEPR model complement each other in understanding the structure of a molecule?
*Lewis structures provide a first approximation of the structure of a simple molecule by showing how atoms share electrons.
*The VSEPR model complements this by offering insight into the shape of the molecule, considering the minimization of electron repulsion and the positioning of electron pairs as far apart as possible.
Why do organic chemists often find the VSEPR model helpful when visualizing chemical reactions, and what aspect of chemical change does it primarily focus on?
*Because it allows them to visualize chemical reactions in terms of changes occurring in the bonds of reacting molecules.
*The model primarily focuses on molecular geometry and the arrangement of electron pairs around atoms.
Despite its usefulness, why is the VSEPR model considered inadequate for some molecules, and what limitation does it have in terms of molecular orbitals?
The VSEPR model is considered inadequate for some molecules because it does not allow for the consideration of antibonding molecular orbitals.
In this book, how does the use of both the Molecular Orbital (MO) model and the VSEPR model differ, and what factors influence the choice between these models when describing a molecule?
Both the MO and VSEPR models are used in this book. The choice between these models depends on which one provides the best description of the molecule under discussion.
(VSEPR for almost whole chap 1 [1.7-1.13])
How is the bonding in methane (CH4) characterized, and what are the key features of the four covalent C¬H bonds in methane?
*The bonding in methane is characterized by four covalent C¬H bonds.
*All four identical bonds have the same length (1.10 Å), and the bond angles are identical at 109.5°.
What compound is a methane but with two carbons? What is the type of bond in between the two carbons?
*ethane
*there is a carbon–carbon single bond between the two carbons.
How are bonds represented in a perspective formula, and what is the significance of solid lines, solid wedges, and hatched wedges in such a representation?
In a perspective formula, bonds in the plane of the paper are drawn as solid lines, while a bond protruding toward the viewer is represented as a solid wedge, and a bond projecting away from the viewer is drawn as a hatched wedge. Solid lines must be adjacent to each other.
What information does the potential map of methane provide regarding the charge distribution on carbon and hydrogen atoms, and what conclusion can be drawn from the absence of red and blue areas in the map?
The potential map of methane shows that neither carbon nor hydrogen carries much of a charge. There are no red areas representing partially negatively charged atoms, nor blue areas representing partially positively charged atoms. The absence of partially charged atoms indicates that methane is a nonpolar molecule. This absence is explained by the similar electronegativities of carbon and hydrogen, leading to the relatively equal sharing of bonding electrons.
Despite having only two unpaired valence electrons, how does carbon form four covalent bonds, and what explanation is provided for this phenomenon?
Carbon forms four covalent bonds despite having only two unpaired valence electrons by promoting one of the electrons in its 2s orbital into its empty 2p orbital. This promotion results in carbon having four unpaired valence electrons, allowing it to form four covalent bonds and complete its octet.
Despite using an s orbital and three p orbitals to form the four C ¬ H bonds in methane, why are they considered identical, and what concept explains this phenomenon?
The four C ¬ H bonds in methane are considered identical because carbon utilizes hybrid atomic orbitals. The concept that explains this phenomenon is hybridization.
What are hybrid orbitals, and how are they formed through the process of hybridization?
Hybrid orbitals are mixed orbitals resulting from the combination of atomic orbitals. They are formed through a process called hybridization, which involves combining atomic orbitals.
What is the result of combining the one s and three p orbitals of the second shell, and how are these resulting orbitals apportioned?
Combining the one s and three p orbitals of the second shell results in four equal orbitals, each being one part s and three parts p. These mixed orbitals are called sp3 orbitals.
What does the term “sp3” imply, and what is the composition of each sp3 orbital in terms of s and p character?
The term “sp3” implies that one s orbital and three p orbitals were mixed. Each sp3 orbital has 25% s character and 75% p character.
Why are the four sp3 orbitals considered degenerate, and what does degeneracy mean in this context?
The four sp3 orbitals are degenerate, meaning they all have the same energy level.
What does an sp3 orbital looks like?
*because of overlap, and 25% s character, 75% p character, lobe sizes look a bit different than a usual p orbital
*the s orbital adds to one lobe of the p orbital and subtracts from the other lobe