Sample Question #1 PDF

Preview

Summary

lecture...

Description

Sample Question #1 Select the most appropriate Lewis structure for chlorate anion: ClO3-

The correct answer is C. It is generally best to start from scratch and go through all of the steps to obtain the correct Lewis structure, but in brief: the three structures show the correct total number of electrons (26 in total), but different degrees of charge minimization, where ‘b’ is not charge minimized at all, ‘a’ is partly charge minimized, and ‘c’ is fully charge minimized. Structure ‘c’ has 12 electrons on chlorine, but that is OK, given that chlorine is in the 3rdperiod, so the fully charged minimized structure ‘c’ is correct. ●

●

The formal charge on chlorine in structure ‘c’ is 7-2-5= 0. As noted earlier, the formal charge is the charge that an atom would have if all bonds are entirely covalent.By contrast, the oxidation state of chlorine in structure ‘c’ is 5+, because oxygen is more electronegative than chlorine, so we can view the structure as Cl(5+) bonded to three O(2-) groups, giving an overall The oxidation state is the charge that would arise if all of the electrons in each bond are transferred to the most electronegative atom. Therefore, oxidation states are the charges that arise if we assume complete ionic character.

Sample Question #2 Select from the following Lewis structures the one that is the most appropriate for CO?

●

●

●

Lewis structure drawing, then you end up with the structure shown on the bottom left of the slide, with one lone pair and a formal positive charge on carbon, and 3 lone pairs and a formal negative charge on oxygen. We can then charge minimize to afford structure ‘c’. However, there is a problem with structure ‘c’, which is that carbon does not have an octet of electrons, and to fix this problem,we need to push one of the lone pairs on oxygen over to form a triple bond, thereby forming structure ‘b’. This is an odd-looking structure, given that we have adjacent negative and positive charges, and because we have 3 bonds to oxygen resulting in a positive charge on oxygen. However, it is the only structure that gives carbon an octet of electrons,which is more important than charge minimization. It is also worth noting that structure ‘a’ is not sensible, given that it features a quadruple bond, which is impossible outside of transition metal chemistry. Also, oxygen in this structure has 10 electrons, which is not possible for a period 2 element. Structure ‘d’ is also not sensible, given that both oxygen and carbon have just 6 electrons

Part 3 Resonance Structures for PO4 3● For PO4 3- there are 4 equivalent charge-minimized structures (resonances structures) ● The molecule exists as a hybrid of all formal resonance structures, known as the resonance hybrid ○ Resonance hybrid bond lengths, orders and charges are the average of all equivalent resonance states ○ Resonance structures contributing to a resonance hybrid must have the same atomic arrangement; they differ only in how the electrons are arranged, they differ only in how the electrons are arranged. but it is worth noting that most polyatomic anions, like nitrate, sulfate, chlorate, acetate etc. have resonance structures ○ We show contributing resonance structures joined by a double headed arrow, which is not the same as an equilibrium arrow. ○ This does not mean that the molecule has one structure part of the time and another structure the rest of the time. It exists as a single structure all of the time, which is an average of all the resonance contributors. ● Most polyatomic anions have resonance structures

●

●

We can generate the Lewis structure for phosphate, which has phosphorus in the centre, with one double bond to an oxygen atom with a formal charge of zero, and three single bonds to oxygen atoms with a formal negative charge. This might suggest that one of the oxygen atoms is different from the rest, but this is in fact not the case, as a result of resonance structures. Specifically, we can use a lone pair of electrons from one of the negatively charged oxygen atoms to form a new double bond with phosphorus, and at the same time, move a pair of electrons from the original double bond onto oxygen, so that there are now 3 lone pairs on that oxygen, accompanied by a formal negative charge. ○ In this way, we can generate 4 resonance structures, where the double bond between phosphorus and oxygen involves each of the four oxygen atoms

Resonance Structures for PO4 3To calculate the average formal charge on an atom, we take the total charge on that type of atom and divide by the total number of that atom

To calculate the average bond order, we take the total number of bonds involved in the resonance structure and divide by the total number of places that the bond is formed

Sample Question #3 Rank the following molecules in order of increasing average formal charge on terminal O’s (from negative to positive):

●

●

●

The correct answer is D. In HCO3–, the negative charge can be shared between either of the two TERMINAL oxygen atoms(but not the internal oxygen atom), so the average formal charge is –1/2. The internal oxygen atom is not included in this process because no charge minimized resonance structure can be drawn that involves delocalized the negative charge to the internal oxygen atom.In CO32–, we can draw 3 resonance structures, allowing the two negative charges to be spread equally between the 3 oxygen atoms. Therefore, the average formal charge is –2/3.In CO2, the average formal charge on oxygen is of course zero.

Sample Question #4 Rank the following molecules in order of increasing terminal C-O bond order:

The correct answer is D. ● In HCO3–, the average terminal C–O bond order is 3/2 = 1.5. ● In CO32–, the average terminal C–O bond order is 4/3 = 1.33. ● In CO2, the average terminal C–O bond order is 2 Bond Order, Length & Energy •Covalent bond length •Distance between 2 nuclei involved in a covalent bond •Bond dissociation energy (homolysis) •Approx. energy required to break 1 mol of bonds in gas phase •which is the energy required for bond homolysis, meaning that when we split the bond, bonding electrons are divided equally between the atoms involved •As an example, bond homolysis of H2 gives two hydrogen atoms, each with one electron. It does not give an H+and an H–. Bond dissociation energy is the energy required to break 1 mole of a particular bond in the gas phase •As bond order increases, bond length decreases •As bond length decreases, bond energy increases.

If we are talking about a specific bond in a specific molecule, such as the C–C single bond in ethane, or the N–N triple bond in N2, we can define a specific bond length and bond dissociation energy. However, often we use tables that give approximate values for particular types of bonds, that are an average of the values for many different types of molecules containing this type of bond

We can note several trends: 1. For a particular type of bond (e.g. a bond between two carbon atoms), as bond order increases, bond length decreases. This makes sense, since a stronger bond would be expected to pull the atoms involved in the bond closer together. 2. For a particular type of bond (e.g. a bond between two carbon atoms), as bond length decreases, bond energy increases. This also makes intuitive sense, since a shorter bond would be expected to occur when two atoms are more strongly bonded together, in which case more energy will be required for bond homolysis Molecular Shape ● VSEPR theory, which stands for Valence Shell Electron Pair Repulsion theory, provides a method to predict the 3-dimensional structures of molecules. ○ This theory is sometimes called Gillespie–Nyholm theory after its originators, one of whom, Ron Gillespie, was in the Department of Chemistry at McMaster. ● This theory recognizes that the electron groups of a central atom in a structure will repel each other, and tries to minimize this repulsion. ○ Electron groups refer to lone pairs or bonds, whether or not the bond is a single, double or triple bond. Note that in VSEPR, electron groups used to be called electron pairs. However,in recognition of the fact that a double or triple bond contains more than a pair of electrons, we now refer to electron groups, rather than electron pairs (unless we are referring only to lone pairs and single bonds, in which case the terms are interchangeable). ● Repulsion between electron groups decreases in the order: lone pair –lone pair > bonded pair –lone pair > bonded pair –bonded pair. This effectively means that a lone pair takes up more space than a single bond. ○ By contrast, a double bond occupies slightly more space than a lone pairs. VSEPR Classes AXnEm ● A = central atom ● X = atoms bonded to central atom ● E = number of lone electron pair For example, NH3 would be an AX3E molecule, because there are 3 atoms bonded to nitrogen, and there is one lone pair on nitrogen. ●

●

Electron group geometry dictates the observed molecular geometry. Note that these are often not the same thing. The electron group geometry includes any lone pairs, whereas molecular geometry does not. Therefore, the electron group geometry and the molecular geometry for an AX3E molecule like ammonia will be different

2 Electron Group

● ●

●

●

Electron Group Geometry –Linear The only molecules that we need to consider with 2 electron groups are those with an AX2 formula (i.e. two atoms bound to a central atom, with no lone pairs on the central atom). In this case,the molecular geometry will be linear,just like the electron group geometry, with an angle of 180° around the central atom. ○ Examples are BeCl2 and CO2. AX2 molecules have a high degree of symmetry, and we will refer to them here as“symmetrical”. ○ Without additional information, the term “symmetrical” is not very specific. ○ However, it is useful to classify certain molecules as “symmetrical” at this point, so that we can refer back to them later when we are discussing molecular polarity. ○ If you imagine that each bond from A to X is a rope and each X represents an equally matched team in a tug of war, the molecules designated as “symmetrical” are those where the efforts of all the teams would cancel each other out, and the central atom will stay in position.

3 Electron Group ● For a molecule with 3 electron groups, the electron group geometry is trigonal planar, so as to place all electron groups as far apart as possible. ○ Example is BF3. ●

AX3 molecules will have a trigonal planar molecular geometry (the same as the electron group geometry), and the angles around the central atom will be 120°. ○ We will classify these molecules as symmetrical

●

●

AX2E molecules, such as the BH2–anion, will have a bent molecular geometry, and are classified as “asymmetrical”. ○ The X–A–X angle in an AX2E molecule will be approximately 120°. However, the exact angle will be a bit less than 120° because lone pair–bonding pair repulsion is higher than bonding pair–bonding pair repulsion(i.e. a lone pair takes up more space than a single bond). Whenever deviations from ideal geometry are indicated, they are based on the fact that a lone pair takes up more space than a single bond. Therefore, they may only be correct for molecules that contain exclusively...