chem 121 exam 3

linear, 180

2 atoms and 4 lone pairs around the central atom:

Sentence 1: Partly incorrect; although single covalent bonds can be made by sharing two electrons, a single electron can still be shared between two nuclei to form a bonding and antibonding molecular orbital. Sentence 2: Incorrect; the H2+ ion exists and two molecular orbitals (bonding and antibonding) are still formed. Although only one electron is shared between two nuclei, the electron occupies the bonding molecular orbital, lowering its PE and KE, making it favorable to form a bond between the two H nuclei. The bond order = ½. Sentence 3: Incorrect; the H22- ion cannot exist as no bonds are formed between the two nuclei. Of the four shared electrons, two electrons occupy the bonding orbital and two electrons occupy the antibonding orbital. Therefore, the electron energies in sum are raised (i.e. the bond order = 0), which means that bond formation is not favored.

A CHEM 121 student was asked to consider the possible interactions between two hydrogen nuclei and explain whether the H2+ and H22- ions exist. The student's response was as follows: A covalent bond can only be made by sharing two electrons to form a bonding and antibonding molecular orbital. The H2+ ion cannot exist because there is only one shared electron and only one bonding orbital formed. Conversely, the H22- ion exists because sharing four electrons creates a double bond between the two hydrogen nuclei.

similar EN, high EN, share electrons, low melting point similar EN, low EN, sea of electrons, high melting point big difference in EN, middle EN, do not share electrons, solid at room temp

model for bonding: EN, electrons, melting point? covalent-- metallic-- ionic--

lower -- non-metal bonded to a non-metal; higher -- metal bonded to a non-metal as you go down the periodic table, EN decreases and the dipole moment decreases, as seen in the dipole data with F, Cl, Br, and I bonded with H as you go across, EN increase, as seen where H is bonded to F and O, HF is higher the H2O, NH3 and CH4 (CH4 lower) the more symmetrical, the closer to 0--CO and CO2

molecular dipole moments: trends?

note about electronegativity vs electron affinity

note about electronegativity vs electron affinity note contrast: electron affinity is attracting an electron to a free atom from "infinity"

a) Whereas KBr is composed of ±1 ions, MgO has ±2 ions, which increases the strength of Coulombic attractions in the lattice. This translates to a higher melting point for MgO, as more energy is required to overcome its larger lattice energy. b) A Cd lattice would have more electrons filling its antibonding molecular orbitals, which would weaken the metallic bonds between the Cd atoms, thereby lowering its melting point. c) KBr (an ionic solid) and SiO2 (a covalent solid) have localized valence electrons, which would not allow for electrical conductance. d) SiO2 is a covalent solid, which does not closely pack like Mo atoms in a metal lattice.

or each of the following questions, circle the correct substance(s) and briefly explain your reasoning: a) Which substance would have a higher melting point: MgO or KBr ? b) Which substance would have a lower melting point: Mo or Cd ? c) Which solids(s) would be an electrical insulator: Cd, KBr, or SiO2 ? d) Which substance would be less dense: Mo or SiO2 ?

clump above and below like a dumbbells when looking at the bonding axis

pi symmetry

seesaw, < 90 and < 120

4 atoms and 1 electron around the central atom:

square planar, 90

4 atoms and 2 lone pairs around the central atom:

tetrahedral, 109.5

4 atoms around the central atom:

square pyramid, <90

5 atoms and 1 lone pair around the central atom:

trigonal bipyramid, 90

5 atoms around the central atom:

octahedral, 90

6 atoms around the central atom:

electrostatic attraction between cation and anion lowers the PE

Transferring an electron from a Na atom to a Cl atom is not energetically favorable. Why does NaCl still form?

Valence Shell Electron Pair Repulsion Theory repulsion of electron pairs in the valence shell, want to minimize the repulsion of electron pairs; electron pairs will be as far away from each other as possible

VSEPR: determined by:

The electrons in a metal are delocalized, moving in the positive environment created by positively charged cores of the metal atoms. As these atoms are displaced, the energy of the electrons will change, but the structure will remain bonded. Given the flexibility, the metal atoms will return to the shape which provides the lowest energy for the delocalized electrons. Hence, a spring will return to its "rest" position. In a salt, the bonding is due to localized attractions of positive and negative ions. When these ions are displaced, the localized interactions are disrupted and the bonding is broken. Once broken, it takes high energy to restore these interactions.

When a metal can be shaped into a spring, the metal can be stretched but will return to its original shape when released. Explain this behavior in terms of the bonding of a metal.

polarity

geometric isomers can have different....

bent, less than 120

2 atoms and 1 electron pair around the central atom:

bent, << 109

2 atoms and 2 electron pairs around the central atom:

linear, 180

2 atoms and 3 lone pairs around the central atom:

linear, 180

2 atoms around the central atom:

trigonal pyramid, less than 109

3 atoms 1 electron pair around the central atom:

T-shape, <90

3 atoms and 2 lone pairs around the central atom:

T-shape, <90

3 atoms and 3 lone pairs around the central atom:

trigonal planar, 120

3 atoms around the central atom:

Sentence 1: Incorrect. Metallic bonding occurs between atoms that both have very low electronegativity. Sentence 2: The first part is correct, in that neither atom attracts the bonding electrons more. However, the second part is not correct, as the electron are free to move between many (not just 2) atoms. Thus the electron sea is lots of electrons from lots of atoms moving freely between all of them. Sentence 3: Incorrect. The melting point only increases halfway across the periodic table. This can be explained by band theory, in which the atomic orbitals of the metal atoms combine to form bands of bonding (lower energy) and antibonding (higher energy) molecular orbitals. As you go across the periodic table, the increasing electrons go into bonding MOs and increase the melting point. However, once the bonding band is full (after d5), the increasing electrons go into antibonding MOs and the melting point decreases.

A CHEM 121 student was asked to explain metallic bonding and how it relates to a metal's bulk properties, and her answer is shown in italics below: Metallic bonding occurs between two atoms that both have very high electronegativity. Since the electronegativities are similar, neither atom attracts bonding electrons more than the other, so the electrons are free to move between the two atoms in an electron sea. The more valence electrons are shared in the sea, the higher the melting point, so melting point increases steadily as you go across the periodic table.

False. Red absorbs green light and blue absorbs orange. Since green is higher energy light than orange, the red solution must have a greater d-orbital splitting, so it must have the higher field ligand. False. Red absorbs green light and blue absorbs orange. Since green is higher energy light than orange, the red solution must have a greater d-orbital splitting. Octahedral splitting is higher than tetrahedral, so the red is octahedral. True. It is both bad practice and unsafe to leave unlabeled solutions in lab.

A CHEM 123 student comes across two unlabeled flasks of colorful solutions, both of which contain an unknown Co3+ complex dissolved in water. One is red and the other is blue. True or False: If both are octahedral complexes, the blue one has a higher field ligand. True or False: If one is octahedral and one is tetrahedral, the blue one is octahedral. True or False: The original student who made the solutions should have labelled the flasks.

1) The first sentence is correct 2) The second sentence is incorrect: the electrons in the antibonding orbitals are at higher energy, but they are still below the zero of energy where they would be ionized; it still requires an ionization energy to remove them. Thus they will not do so spontaneously (i.e. without input energy). In actuality, He2 would simply fall apart into two He atoms. 3) The third and fourth sentences are also incorrect: He2 would decompose into neutral He atoms, and the electrons from the antibonding orbital would just relax back to their energy in separate atoms. In fact, with only two electrons left, He22+ would have a single bond (bond order of 1) and exist as a stable molecule. Although (in general) 2 positive ions repel each other, they don't exist here to do so, instead the diatomic ion exists, meaning that the bond formed is stronger than any repulsion from the two positive nuclei.

A CHEM121 student was asked to use molecular orbital arguments to discuss the bonding in He2. The student replied: In He2, there would be two electrons in a bonding orbital, but there would also be two electrons in an anti-bonding orbital, therefore He2 does not exist. If He2 did exist, the high energy electrons would spontaneously ionize. This would form two He+ ions. These ions repel each other, which is why the He22+ ion is unstable.

True False False

A solid crystal of table salt forms because: True or False: forming a tightly packed lattice lowers the potential energy between positive and negative ions. True or False: a "sea of electrons" is created. True or False: it is energetically favorable to transfer an electron between elemental sodium and elemental chloride.

Tridentate means that a ligand can donate 3 separate lone pairs to a metal center. Dien can do this because it has 3 nitrogen atoms, each of which has a lone pair. These 3 nitrogens are spaced appropriately to facilitate the geometry needed for each one to bind the same metal.

Briefly explain what tridentate means, and how dien can act as a tridentate ligand.

The potential energy of a σ3p electron is lower than a 3p electron because an electron in a bonding molecular orbital feels Coulombic attraction to two nuclei (rather than only one in an atomic orbital). The kinetic energy of a σ3p electron is also lower than a 3p electron because the size of the σ3s molecular orbital is larger than the 3p atomic orbital, which lowers the confinement energy of the σ3p electron.

Briefly explain why the total energy of an electron in a σ3p molecular orbital is different from an electron in a 3p atomic orbital.

MO AO MO MO MO

Compare the following between that of an atomic orbital (AO) of H and a molecular orbital (MO) of H2: orbital size of an ____ is larger an electron in an ____ is more confined KE of electron in an ______ is lower IE of an electron in an _______ is larger PE of electron in an ______ is lower

The bonds between the ion and the ligand are coordinate covalent bonds, meaning that an electron pair is shared but both electrons are donated by the ligand. If the bonds were ionic, the ions were separate when dissolved in solution, i.e. there would be CN- ions free solvated in water. This does not happen, so the bonds are covalent.

Despite the charges on the ions, the bonds between the CN- ligands and the Fe ion are not ionic. Briefly describe the correct description of the bond that is formed, providing the experimental evidence which shows that these bonds are not ionic.

multiple bonds have the same VSEPR as single bonds

Electron Domain Theory:

The H-O-H angle is smaller than the HCH angle by a few degrees. In CH4, all four electron pairs are bonded to an H atom. In H2O, two of the electron pairs are lone pairs, not bonded to any other atom. Since the angle between the bonded electron pairs is less in H2O than in CH4, then the lone pairs in H2O must be repelling the bonded pairs more than the bonded pairs repel one another. This is reasonable physically, because the second nucleus in the bond will attract the electron pair and thus localize it to a greater extent than a single nucleus can. Since the bonded pairs of electrons are thus more localized, they generate a smaller repulsive effect

Explain how a comparison of the geometries of H2O and CH4 leads to a conclusion that lone pair electrons produce a greater repulsive effect than do bonded pairs of electrons. Give a physical reason why this might be expected.

Two of the four electron pairs on each carbon are shared by both carbons. These two pairs of electrons thus cannot be separated to as great an extent as if they were lone pairs or pairs shared by different atoms, for if they were, it would be physically impossible for them to be shared. Consequently, these two electron pairs, which must stay together, are separated as much as possible from the other two electron pairs. These three "domains" thus form an equilateral triangle about the central carbon atom.

Explain why the octet of electrons about each carbon atom in ethene, C2H4, are not arranged even approximately in a tetrahedron.

Ru2+ has 6 d electrons, so it needs 12 more to satisfy the 18 electron rule. Since each dien donates 3x2=6 electrons, 2 dien ligands are needed.

How many dien ligands would Ru2+ need to satisfy the 18 electron rule? Show all work for full credit

there is attraction (2 nuclei, 1 electron) and repulsion (nucleus, nucleus) bond energy is for H2 almost 2 times that of H2+-- 2 electrons and 2 nuclei = 4 attractions, electron-electron repulsion = 2x repulsion. not exactly 2x as much because electron-electron repulsion

How should the energy of the electron compare when it is shared between two nuclei vs not shared?

A substance is magnetic if it has unpaired electrons in its molecular orbitals, so that the individual electrons do not cancel out each other's magnetism. In O2, the highest energy molecular orbitals with electrons are the π* anti-bonding orbitals. Since there are two of these orbitals, the two π* electrons are not paired, making O2 magnetic. In Zn atoms, the 4s and 3d atomic orbitals are all filled. Therefore, the molecular orbitals in solid Zn are also completely filled, causing all of the electrons to be paired. Therefore, Zn is not magnetic.

In Dr. Kincaid's demo last week, we discovered that liquid O2 was magnetic but solid Zn was not magnetic. Based upon your understanding of molecular orbitals energies, briefly explain why this is the case (note that you do not need to sketch the molecular orbital energy diagrams).

VBT: 1/2 filled atomic orbitals overlap to form a bond MOT: atomic orbitals combine to form molecular orbitals, "fill up" orbitals from lower energy to higher energy with the total number of electrons; 2 atomic orbitals yields 2 molecular orbitals

Molecular Orbital Theory vs Valence Bond Theory

O, F, Ne s 2p lowest, then p 2p, p* 2p, s* 2p p 2p, s 2p, p* 2p, s* 2p BASICALLY SWITCH s 2p and p 2p

O looking shape for orbitals diagrams for... tree shape:

A polar bond is a covalent bond between 2 atoms of differing electronegativity, causing the more electronegative atom to have a stronger pull on the shared electrons. Polarity strength is higher for bonds between atoms of greater differing electronegativities; because H is less electronegative than C, the O-H bond has a greater polarity than O-C.

PRACTICE EXAM Q: Both the C-O and O-H single bonds in all of these compounds are described as polar. Briefly explain what a polar bond is and why the O-H bond is considered more polar than the C-O bond.

The pi bonds between the first two carbon atoms and the last two carbon atoms in allene are on perpendicular planes, at the p atomic orbitals used to form these two bonds (lobes 5/7/10/13 and 9/12/15/18) are perpendicular to one another. Therefore, the terminal CH2 must also reside on perpendicular planes.

PRACTICE EXAM Q: Briefly explain why the CH2 groups at the opposite ends of allene do not lie in the same plane.

s* is higher p* is higher p and p are the same depends on specific molecule!

Predict the relativeenergies of the orbitals: s vs s* p vs p* p(from 2py) vs p(from 2pz) s vs p

the two nuclei will repel one another at close distance the Pauli exclusion principle causes filled AO's to repel each other.

Provide one reason, with a brief explanation, for why there is a limit to how close two atoms can get to one another.

1) the potential energy is lower because the electrons feel the attraction of 2 nuclei instead of just 1 2) the kinetic energy is lower because the molecular orbital is larger

Provide two reasons, each with a brief explanation, for why shared electrons in a covalent bond have lower energy than when the bond is broken.

attractions between opposite charges, repulsions

if Na+ and Cl- are put into one structure, what is maximized? minimized?

Rb and F have ionic charges of +1 and -1 respectively while Mg and O have ionic charges of +2 and -2 respectively in the compounds. Therefore, Mg and O will be more strongly attracted to one another leading to a higher lattice energy. Rb is very large which creates a larger interatomic distance in the lattice of RbF as compared to MgO (which contains to relatively small ions). This leads to a lower lattice energy for RbF.

Rubidium fluoride (RbF) has a lattice energy of 785 kJ/mol, while Magnesium oxide (MgO) has a lattice energy of 3791 kJ/mol. Provide two reasons, with brief explanations, why these lattice energies are so different.

False; Metal atoms almost never pack in the simple cubic packing because it is the least efficient, and results in less overlap of orbitals for bonding.

T/F? Metal atoms typically pack in the "simple cubic packing" pattern, since it is the simplest.

True; The middle of the d-block is where the bonding "band" would be full and give the largest bond order. Going further right in the d-block, the anti-bonding "band" becomes populated and so bond order decreases.

T/F? Metallic bond order is generally highest in the middle of the d-block on the periodic table.

False; Only valence electrons (not core electrons) move freely through atoms.

T/F? The electron sea model describes a bonding mode where all of the electrons in the substance are moving freely between atoms.

The highest energy, most weakly attached electrons in these metals are 4s electrons, not the 3d electrons. As such, the ionization energy of these electrons is determined by the shielding of the nucleus including the 3d electrons. With each added proton, there is an added 3d electron to provide the shield. Hence, the ionization energy does not vary much despite the increased nuclear charge.

The four metals Sc, Ti, V, and Cr have increasing nuclear charge in the order listed, but have only small differences in ionization energy. On the basis of the electron configurations of these elements, explain this unexpected lack of variation.

Each electron in H2 has a slightly higher energy than the single electron in H2+, due to electron-electron repulsion. However, breaking the H2 bond requires elevating the energy of two electrons instead of one, and this requires in total much more energy.

The ionization energy of H2 is slightly less than the ionization energy of H2+. But the bond energy of H2 is much larger than the bond energy of H2+. Explain how these two facts are consistent with each other.

there are more electrons contributed by a multiple bond than by a single bond for the central atom

Why do multiple bonds affect other bond angles around it?

1.simple < bcc < hcp=ccp 2. higher packing efficiency means greatest opportunity for overlap-- hcp and ccp 3. simple--6 bcc--8 hcp and ccp--12

atomic lattice packing: For the 4 packing arrangements shown: 1) Rank in order of increasing packing efficiency. 2) Which has the greatest opportunity for overlap of atomic orbitals? 3) Rank in order of increasing coordination number.

many bonds super close together, tons of molecular orbitals super close in energy, half bonding and half anti bonding

ban theory of metals

(1/2)* (bonding electrons - anti bonding electrons)

bond order:

An electron shared by two nuclei has a lower potential energy due to the sum of the attractions to both positive centers. Each attraction contributes a negative potential, so the sum is a greater negative number than for a single atom. An electron in a molecular orbital has a larger space to move and therefore has less confinement than in a single atom. By the uncertainty principle, this lowers the kinetic energy. Since the energy of the electron is lower in the molecule, the energy of the electron would have to be raised in order to break the chemical bond by separating the atoms. This is measured as the bond energy, which is the energy required to separate the atoms.

Why does an electron shared by two nuclei have a lower potential energy than an electron on a single atom? Why does an electron shared by two nuclei have lower kinetic energy than an electron on a single atom? How does this sharing result in a stable molecule and how can this be measured experimentally?

If the electronegativity difference between bonded atoms is large, it is easy to predict that the bond will be primarily ionic. However, if the electronegativity difference is small, the bonded electrons could both have high electronegativity or could both have low electronegativity. In the former case, the bonding would be covalent; in the latter case, the bonding would be metallic. Hence, the average electronegativity must be known to distinguish metallic bonding from covalent bonding.

Why is it necessary to consider both electronegativity differences between bonded atoms and the average electronegativity of bonded atoms when analyzing the type of bond which is formed?

raise anti-bonding orbitals raise the energy slightly more than bonding orbitals lower the energy

anti-bonding orbitals _______ the energy

polarity

asymmetric distribution of electron density (charge) around a molecule

diamagnetic--no unpaired electrons paramagnetic--unpaired electrons

diamagnetic vs paramagnetic:

v (atoms in unit cell) / v (cube) l = 2r for primitive cubice unit cell l = (4/ sqrt(3))r for body-centered cubic l = 2*sqrt(2) for face-centered cubic p-52% bcc-68% fcc-74%

efficiency % =

-valence electrons are not tightly held, each nucleus attracts electrons briefly the electrons move to another atom -delocalized, move freely from nucleus to nucleus

electron sea model of metals:

somewhere between s and p, energies of these orbitals are the same

energy of hybrid orbitals are located...

2--linear--sp 3--trigonal planar--sp2 4--tetrahedral--sp3

formation of hybrid orbitals--# of electron domains and hybridization 2 3 4 5--trigonal bipyramid, sp3d 6--octahedral--sp3d2

high core charge, smaller radius

high EN means...

stronger bond

higher bond order means...

1. go halfway, increase--add electron to bonding orbitals; more electrons in bonding orbitals means higher melting point more electrons in anti bonding orbitals decreases the melting point going across, adding electrons to d-orbitals, increases bonding and melting point; BUT halfway starts to decrease because half of the orbitals are bonding and other half are anti bonding, anti bonding are higher in energy and decrease the bond strength 2. E=hv where E is delta E; bands of orbitals means many delta Es; any wavelength of light can be absorbed and re-emitted; huge number of E levels, huge # of delta E, and freq. can be absorbed and re-emitted

how does band theory explain: 1. trends in melting point 2. shininess

1. ions are in a fixed position--electrons cannot move, localized to each atom 2. disrupt attractive forces (diagram with moving charge down) 3. lots of +/- attraction, takes a lot of energy to separate 4. once separated there are charges that can move

how does lattice energy explain the following properties of ionic salts: 1. electrical insulator-- 2. brittleness-- 3. high melting point-- 4. conductivity in liquid/ aqueous form-

1. electricity is the flow of electrons, electrons are free to flow around, apply electric potential--easy for electrons to flow around 2. delocalization of electrons--bonds are easy to break and reform, electrons shift with positive charge, electrons are not bound to 1 atom and can easily rearrange to a new shape 3. melt--pull atoms apart, sea of electrons holding the atoms together, sharing lots of electrons which holds the atoms together

how does the electron sea model explain: 1. conductivity 2. malleability 3. high melting point

electronegativity

how much an atoms attracts electrons to itself in context of a bond

higher charge, greater energy (PE more negative), smaller radius, greater energy--Coulomb's law and electrostatic attraction!

lattice energy trends

more

lone electron pairs take up _______ space

simple cubic, 1 atom, 52%, coord #=6 bcc, 2 atoms, 68%, coord #=8 fcc, 4 atoms, 74%, coord #=12

primitive-- bcc-- fcc--

circle/sphere when you look at the bonding axis

sigma symmetry:

sp--2 sp and 2 p sp2--3 sp2 and 1 p sp3--4 sp3 and no p

sp-- sp2-- sp3--

decrease down, increase then decrease across CANNOT explain with electron sea model

trends in melting point of transition metals:

-low IE, all similar -high PE--easy to remove electrons -removing electrons from the same shell -no huge increase in core charge in the d-block, electrons added to d-block are shielding electrons and 4s electrons are valence electrons -2+ ion is common because 2 valence electrons in 4s

what does the ionization energy data indicate about energies of valence electrons in d-block metals?

energy is absorbed--energy must be put into the system to break a bond

what happens when a chemical bond is broken?

electron/electron repulsion, there are only 2 electrons in a molecular orbital, the 3rd electron goes to a different energy level--next orbital (anti-bonding for H)

why 2 electrons and not 3?