Inorganic Chemistry Supplementary Review 2

Define ionization energy.

The energy required to remove the least tightly bound electron from a neutral atom in the gas phase

Describe the periodic trend in ionization energy.

Lowest I.E. starts at the bottom left corner of the periodic table and increases going up and to the right

What is the real reason that a half-filled subshell is so stable?

It serves to maximize the stabilizing interactions, while minimizing the destabilizing repulsions among electrons.

What is the exchange interaction? Is it stabilizing or destabilizing?

The exchange interaction, πe, a stabilizing interaction, and it is a result of electrons pairing in degenerate orbitals with parallel spin

What is the pairing energy? Is it stabilizing or destabilizing?

Pairing energy is a destabilizing interaction, the coulomb interaction, πC, and it is the energy of electron-electron repulsion in a filled orbital.

What is harder to ionize, a high energy electron or a low energy electron?

A high energy electron is easier to ionize. It already contains more energy, so the input to remove it from the atom is less.

If a 3d series metal is ionized, what orbital loses the electron? What happens to the valence configuration?

If a d series metal has the configuration 4s²3d^n, the first electron to be ionized will come from the 4s orbital. The remaining s orbital will then enter a d orbital, giving the configuration of the +1 cation 4s03dn+1.

What is the Lanthanide Contraction?

It is the reduction in atomic radius following the lanthanide series, contrary to the overall trend observed for the periodic table. The lanthanides are elements 57 through 71, and are the first appearance of filled f orbitals. F orbitals are poor at shielding, so any electrons added into the subshell following the f orbital experience a higher Zeff, shrinking their radius over what is to be expected from overall trends.

Describe Slater's Rule.

Slater's Rules tell us what the effective nuclear charge will be.

Formally, it is written as Zeff= Z - σ.

Z is the atomic number, and σ=Σsi

Si is the number of electrons in a given subshell multiplied by a weighting coefficient.

The weighting coefficients are:

1) A. For s and p valence electrons, (n-2) groups or lower contribute 1.00 to si (weighting coefficient of 100%). 1.00xne

B. (n-1) groups contribute 0.85 to si. 0.85xne

C. N(s,p) groups (the valence groups) contribute 0.35 to si. 0.35xne

2) For d and f valence electrons:

A. electrons in the valence shell contribute 0.35 to si. 0.35xne

B. electrons in any lower group contribute 1.00 to si. 1.00xne

What is shielding?

Shielding is the reduction in charge attraction between the nucleus and electrons, due to electrons between the nucleus and the electron in question. An electron is 'between' the nucleus and the electron being considered if it has a lower energy.

What is penetration?

Penetration is when an electron of a higher atomic orbital is found within the shell of electrons of lower atomic number. That is to say that an electron of higher energy is found within an orbital of lower energy.

What is electron affinity?

Electron affinity is the difference in energy for a neutral gaseous atom, and the gaseous anion. It is often used interchangeable with electron gain enthalpy, which is the change in standard molar enthalpy when an atom or molecule gains an electron. A more positive electron affinity means that the atom is more stable with the additional electron, while a more negative electron gain enthalpy is indicative of greater stability.

What is the combination of electron affinity and ionization energy called?

The total effect of electron affinity and ionization energy is known as electronegativity. It is the overall measure of an atom's ability to attract electrons to itself when part of a compound. Fluorine has the highest electronegativity.

To explain electronegativity, theories have been proposed by Mulliken, Allred-Rochow, and Pauling. Pauling's theory is used most commonly.

What is polarizability?

It is an atom's ability to be distorted by an electric field. Basically, it means that regions of a molecule can take on partial positive, or partial negative charge.

Why do we have to use the hydrogenic system approximation?

Systems involving multiple electrons are much more complex, and they require the use of quantum mechanics.

What is the formula for the energy of a hydrogenic orbital?

13.6(eV) is the result of h*c*R, and converting to eV. The negative is a result of the sign convention of energy levels. R is the Rydberg constant, which = 1.09737x10^5 cm-1, c is the speed of light in cm, 2.998x10^10 ms-1 , and h is Plank's constant, 6.626x10^-34 Js.

What is another way of expressing the energy? Joules, and...

Wavenumbers, which are the inverse of wavelength, expressed as inverse centimeters.

Define and describe the quantum numbers n, l, ml, ms.

N: The principle quantum number. It defines the energy and size of the orbital.

L: Orbital angular momentum quantum number: it defines the magnitude of the orbital angular momentum, as well as the angular shape of the orbital. L can have values of 0 to n-1.

Ml: Magnetic Quantum number: Describes the orientation of the angular momentum. Ml can have values of 0 to +/- l.

Ms: Spin magnetic quantum number: Defines the intrinsic angular momentum of an electron. Ms can have values of either +1/2 or -1/2.

It can be seen from these definitions that every electron in an atom or molecule will have a different set of quantum numbers.

What is the radial wavefunction?

The radial wavefunction (R(r)), along with the angular wavefunction (spherical harmonic), gives us the orbitals. With a wavefunction, it is possible to completely characterize a particle. The wavefunction does not go to zero at the nucleus, they go to zero at infinity. When graphed, these produce the characteristic shapes of the orbitals we are used to seeing.

What is the radial distribution function?

It is a plot of R²(r)r² vs the radius. It tells us the probability of finding an electron at a certain distance from the nucleus. Every orbital has a different radial distribution function, and a node on the graph is a region of zero probability.

d. What is the Bohr radius?

The Bohr radius is the most probably distance to find the electron in a one proton, one electron system.

It is: a0= (ε0ђ²)/(π(me)e²)= 52.9 pico-meters.

What orbitals correspond to l=0 through l=4?

=0=s, l=1=p, l=2=d, l=3=f, l=4=g.

What is the building up principle/Hund's Rules?

When degenerate orbitals are available for occupation, electrons occupy separate orbitals with parallel spin.

What is the Pauli exclusion principle?

No more than two electrons can occupy a single orbital, and to do so, their spins must be paired.

Describe Valence Shell Electron Pair Repulsion Theory.

The purpose of VSEPR theory is to predict molecular geometries. The basic assumption to the theory is that regions of enhanced electron density take positions as far apart as possible, in order to minimize repulsive forces. The relative repulsive strengths of electron dense regions are: lone pairs > multiple bonds > single bonds.

Be able to predict molecular shapes based on VSEPR theory

IDK how to import pictures into quizlet so you're on your own.

What is the effect of a lone pair on geometry?

A lone pair pushes strongly against all other substituents. It is the strongest steric force governing the shape of a molecule.

Describe Molecular Orbital Theory and the assumptions.

It is an improvement over valence bond theory in that the bonding description extends to all atoms in a molecule, and handles polyatomic molecules easily. Atomic orbitals combine to form molecular orbitals, which are delocalized descriptions of electron distribution.

MO theory makes several key assumptions:

1. The orbital approximation: The wavefunction describing all of the electrons of a molecule can be written as a product of the one electron wavefunctions.

2. Linear Combinations of Atomic Orbitals: the superposition of multiple atomic orbitals of same type along with weighting coefficients.

What is bond order, and how is it calculated?

Bond order is a method to assessing overall bond strength between two atoms in a molecule. It directly assesses the net number of bonds between them. The higher the bond order, the stronger the bond.

B=1/2(n-n*)

n= # electrons in bonding orbital, n*= # electrons in antibonding orbital

Know how to assign bonding character (bonding, nonbonding, and antibonding).

A bonding orbital will lie lower in energy than its substituent atomic orbitals. An antibonding orbital will lie higher in energy than its substituent atomic orbitals. A nonbonding orbital will equal in energy to its corresponding atomic orbitals.

Know the significance of HOMO/LUMO interactions.

HOMO: highest occupied molecular orbital. It is where the highest energy electrons reside.

LUMO: lowest unoccupied molecular orbital. It is the energy level directly above the HOMO.

HOMO/LUMO interactions are critical to electronic structure. When an excited state is formed, an electron is promoted from the HOMO to the LUMO. The bonding and antibonding characteristics of the frontier orbitals are of great importance to the stability of the molecule. Basically, whereas in neutral atoms we treat valence shell electrons as the point of interaction, the frontier orbitals are the point of interaction and basis of understanding for molecular interactions.

How is paramagnetism identified through a MO diagram?

If there are two unpaired electrons, the species is paramagnetic.

Define Metallic bonding:

the bonding and ordering of metals into pure solids, or solid solutions, viewed as enormous molecules with continually overlapping atomic orbitals.

Define ionic bonding.

ions of different elements held together in rigid, symmetrical arrays as a result of attraction between their opposite charges

Define Lattice

A 3D infinite array of (lattice) points which define the repeating structure of a crystal.

Define Unit Cell

an imaginary, parallel sided region from which the entire crystal can be built, such that it fits perfectly together, giving rise to a crystal system. There are 7 types of crystal systems: cubic, tetragonal, orthorhombic (3 different length dimensions, identical angles), monoclinic (3 different dimensions, defining angle between two sides), triclinic (3 different dimensions, 3 different angles), rhombohedral (all 3 sides the same length, all angles are equal and non-90), and hexagonal.

Define primitive cell

Contains only one lattice point.

Define Body Centered

two lattice points per unit cell

define face centered

Four lattice points per unit cell

Define centered cubic

3D version of body centered

Define Close packed structure

structure type with the least unfilled space

Define coordination number.

The number of nearest neighbors. It is the number of other spheres that a single sphere is touching in its unit cell

Define cubic close-close packed

ABCABC repeating layers (ABAB results in hexagonally close packed)

Define face-centered cubic

the same thing as cubic close-packed.

Define hole.

unoccupied space between spheres. Holes can be octahedral, where they lie between two triangles and form a hexagon. They can also be tetrahedral, where it is formed by a planar triangle of touching spheres.

Define polymorphism.

The ability of a metal to adopt different crystal structures based on the temperature and pressure.

Define Alloy

A blend of different metals. More formally called a solid solution.

Define Substitutional solid solution

When solute metal atoms take up the positions of solvent metal atoms in the crystal structure.

Define interstitial solid solution

When the solute metal atoms take up positions in the holes (interstices) of the solvent crystal structure.

Define lattice enthalpy.

The standard enthalpy change of accompanying the formation of a gas of ions from a solid. It is a measure of the strength of the solid.

Define and understand a Born-Haber Cycle.

A type of flow chart used to determine lattice enthalpies and other parameters governing how a solid behaves.

Define Van der Waals interaction.

Non-electrostatic contributions to the lattice enthalpy. The London Dispersion interaction is the most dominant van der Waals interaction. It is the transient fluctuations in electron density that result in temporary, induced dipoles.

Define non-stoichiometric compound

A substance that exhibits variable composition, but retains the same structure type.

Define conductor

A substance with an electric conductivity that decreases as temperature is increased.

Define semi-conductor

A substance with an electric conductivity that increases as temperature is increased.

Define insulator

Negligible conduction, but if it is possible to measure, it increases with temperature.

Define super-conductor

have zero electrical resistance below a critical temperature

Define band

A near continuous array of energy levels, due to the large number of symmetrically oriented atomic orbitals in a solid with very similar energy levels.

Define band gap

A separation of bands, due to there being no energy value for the molecular orbital.

define s/p band

The bands built from s and p orbitals.

define fermi level.

The highest occupied energy level in a solid at T=0.

What are the parameters for a substitutional solid solution?

1) The atomic radii of the elements are with about 15% of each other.

2) The crystal structures of the two pure metals are the same. This indicates that the directional forces between the two atom types are compatible with each other.

3) The electropositive characteristics of both compounds are similar. Otherwise, electron transfer would be likely.

What are the parameters for an interstitial solid solution?

Interstitial solid solutions usually form between metals and small atoms (group 2). In order for the compound to form, the small atom must not transfer electrons, or an ionic species would result. For the small atom to fit in the hole, its radius must be close to 0.414x(radius metal). Interstitial compounds can have a recognizable molecular formula, or they can form non-stoichiometric compounds.

Describe how to write a Born-Haber cycle, and give an example. (formatting looks like shit be warned)

Starting from the solid of interest, of form MX(s):

1. Standard molar enthalpy of decomposition, ΔdH0 (the negative of the standard molar enthalpy of formation, ΔfH0).

MX(s) -> M(s) + X(s, l, g) (elements are now in their reference states)

2. Standard molar enthalpy of atomization for a solid, ΔatomH0 is the same thing as sublimation

M(s) -> M(g), ΔatomH0

3. Standard molar enthalpy of atomization for an element is the same thing as the standard enthalpy of dissociation, ΔdisH0

X2(g) -> 2X(g), ΔdisH0

4. The standard molar enthalpy of formation of ions from neutral atoms is the enthalpy of ionization (cations) and election gain enthalpy (anions).

M(g) -> M+(g) + e-, ΔionH0

X(g) + e- -> X-(g), ΔegH0

For a complete cycle, the total change in enthalpy will be zero. ΔdH0, ΔatomH0, ΔdisH0, and ΔionH0 are all positive, and ΔegH0 is negative. The value required to make this equal zero is the value for the lattice enthalpy, ΔHL0, and it will be negative.

What is the most stabilizing influence on a solid? Set up a proportion to show how it changes along with the values, and relate this to different types of systems.

Lattice enthalpy is the most stabilizing influence on a solid.

ΔHL⁰ ≈ q¹q²/d, if the charge is high and distance small, lattice enthalpy is very strongly stabilizing. If the charge is small and the distance is great, the solid is not especially stable. ΔHL⁰ values are at their highest when the solid is tightly packed, and the ions are charge dense.

f. Discuss the importance of the Born-Mayer equation. (I did not format the equation because no)

The Born-Mayer equation allows us to estimate the lattice enthalpy from charge and separation, at T=0.

ΔHL0= (((6.022×〖10〗^23 )×√((q"1" q"2" )^2 )×e^2)/4πϵd)×(1-(d^*/d))×A

Where d is the distance between neighboring cations and anions, d* is a constant, A is the Madelung constant, and ε is the permittivity.

Why does solubility depend on lattice enthalpy?

Solubility can be described as a product of lattice enthalpy and hydration enthalpy. Since lattice enthalpy is the energy required to break apart an ionic solid, it is logical to conclude that solubility will depend on that value. It can be put concisely that if there is a large difference between ionic sizes, the compound is likely to be soluble in water. This is because of the ways in which lattice enthalpy and hydration enthalpy depend on the radius.

ΔHL0 ∝1/(r+r)

ΔhydH ∝ (1/r+) + (1/r-)

This means that as the difference in ionic size increases, the lattice enthalpy decreases, while the hydration enthalpy increases.

Separation of the ions in solution is also entropically favored, given that the hydration enthalpy for one is greater than the other.

What is solvent levelling?

Solvent levelling is the ordering of water molecules around a dissolved ion. The more charge dense an ion is, the more it orders the water. This is entropically unfavorable, so it acts against processes that seek to break apart ions in solution.

Considering the discussion in the previous section, this demonstrates that when considering enthalpy, entropy can't be ignored. While solubility is proportional to the lattice and hydration enthalpies, it is also proportional to the overall enthalpy AND entropy. The entropic effects of solvent levelling push back against the favorable hydration enthalpy.

Discuss band theory, including the consequences for conductors and insulators.

The best way to describe band theory is to immediately draw analogies to molecular orbital theory. In MO theory, the molecular orbitals formed by the constituent atomic orbitals reside at different energy levels; have specific orientations, and a degree of overlap. The combination of all of these factors leads to the characterization of MO's as bonding, nonbonding, or antibonding, separated by the previously mentioned energy levels. In band theory, instead of the orbitals of just a few atoms overlapping, there is a near continuum of overlapping orbitals. The lowest energy orbitals have no nodes between neighboring atoms, whereas the highest energy orbitals have a node between every pair of neighbors. Because of the near continuum of MO's, the energy separation between neighboring orbitals approaches zero, creating a near continuous energy spectrum.

The s-band forms between s orbitals, the p-band between p orbitals, etc...

There may be a gap (band gap) between orbital bands, or they may overlap.

Conductors have partially filled bands at low energy levels, so electrons are promoted easily. Vigorous vibrations, such as those due to thermal energy, disrupt the band structure, and decrease conductivity.

Insulators have filled bands with a large energy differential separating them. The filled band is called the valence band, and the next higher band is called the conduction band.(Refer to picture in study guide

Define Instrinsic semiconductor

like insulators, they have a filled s-band and empty p-band, but the energy gap is very small. This makes it so that electrons can be promoted by thermal energy. The promotion of electrons results in positive holes in the valence band, amounting to a charge separation, which allows current to flow. The population of the conduction band shows an Arrhenius-like dependence on temperature.

Define Extrinsic Semiconductors. Define N- and P-type.

A material gains the ability to be a semiconductor because of intentionally added impurities (dopants), which increase the number of electron carriers.

a. N-type: the N meaning negative, the dopant supplies additional electrons to the solid structure. These electrons lie above the valence band in energy, and close to the conduction band, so the gap is much easier to jump, thanks to the donor band.

b. P-type: The P means positive, the dopant material has less valence electrons than the parent material. The holes it creates lie just above the valence band in energy, so electrons can jump to the acceptor band to allow current to flow.

What conditions are needed for electrons to flow through a solid?

A small energy difference and unfilled orbitals

Define a Lewis acid/base.

A Lewis acid is an electron pair acceptor, and a Lewis base is an electron pair donor.

Define a Bronsted-Lowry acid/base.

A Bronsted acid is a proton donor, and a Bronsted base is a proton acceptor. Obviously, this theory focuses on proton transfer as the source of acidity.

Define hard/soft acid/base theory (HSAB).

Hard acids bond in order of:

I- < Br- < Cl- < F-

Soft acids bond in order of:

F- < Cl- < Br- < I-

Soft acids/bases are larger and more polarizable. Hard acids/bases are small and charge dense.

Hard will tend to bond to hard, soft tends to bond to soft.

How does covalence apply to HSAB?

Use these as a rule of thumb, not a law.

1. For a given cation, covalence increases as the size of the anion increases.

2. For a given anion, covalence increases as the size of the cation decreases.

3. Covalence increases as charge increases.

4. Covalence is larger for non-noble configurations.

What contributes to the Gibb's energy of complex formation, other than the type of bond formed?

1. Competition with solvent

2. The rearrangement of the substituents of the acid and base that may be necessary to permit formation of the complex.

3. Steric repulsion between substituents on the acid and base.

f. What are the criteria for categorization of HSAB?

Other than what a different flashcard states there are no definite criteria. It is based on experimentally determined evidence, usually to do with complex formation.

What is an amphiprotic substance?

It is a molecule that can act as both a Bronsted acid and base.

What is the relationship between a strong acid and its conjugate base?

A strong acid will have a weak conjugate base. A strong base will have a weak conjugate acid. A strong acid seeks to remain deprotonated, while a strong base seeks to remain protonated.

Write the expression for the acidity constant (Ka) for a Bronsted acid. Discuss the range of experimentally observed Ka values and what they mean.

Ka=([A-][H3O+])/[HA]

The strongest (aqueous) acid out there is hydroiodic acid, with a Ka on the order of 10¹¹. This means that once introduced to solution, HI immediately and completely converts to I- and H3O+ (well, except for 1 molecule in every 10¹¹). Some very weak acids have Ka on the order of 10⁻¹⁵. Just about no molecules lose a proton with a Ka that small. Most acids fall somewhere between 10⁻⁵ and 10¹, such that they establish an actual equilibrium.

What is autoprotolysis? Why does this make some calculations much, much easier?

It is the proton transfer from one water molecule to another. This establishes a constant equilibrium, even in neutral water, between hydronium and hydroxyl ions.

This comes in handy in calculations, because we can set up the expression:

Kw=[H3O+][OH-].

Kw= 1.0x10⁻¹⁴ (at 25°C), and we can define Ka and Kb in terms of Kw, allowing us to easily convert between the two. Ka∙Kb=Kw

Discuss proton affinity and proton gain enthalpy.

Proton affinity (AP) and proton gain enthalpy (ΔpgH0) relate acidity to enthalpy changes during a proton transfer. Proton gain enthalpy is reported as a positive number, and the proton affinity is the negative of it. A high AP indicates strong basicity, if only slightly negative, it is only weakly basic or possibly acidic.

Proton gain can be described by these 3 steps:

1. Electron loss from A-:

A-(g) → A(g) + e-

- ΔegH⁰(A) = Ae(A)

The reverse of electron gain by A

2. Electron gain by H+:

H+(g) + e- → H(g)

- ΔiH0(H) = -I(H)

The reverse of ionization of H

3Combination of H and A:

H(g) + A(g) → HA(g)

-B(H-A)

The reverse of H-A bond dissociation

The proton gain enthalpy is the sum of the preceding enthalpy changes:

ΔpgH0(A-) = Ae(A) - I(H) - B(H-A),

AP(A-) = B(H-A) + I(H) - Ae(A)

What is the dominant factor in determining acid character?

Electron affinity is the dominant factor. Electron affinity and proton affinity are inversely related, so observing the trend of Ae increasing from left to right, the proton affinity increases right to left. As proton affinity decreases, acidity increases, so Ae and stronger acid character have a positive, direct correlation. A higher Ae corresponds to a more stable conjugate base.

Discuss solvent levelling.

Solvent levelling in water is the fact that any strong acid or strong base in aqueous solution is only as strong of an acid as H3O+ or as strong of a base as OH-. The relative strengths of strong acids can be measured by using an anhydrous weak acid as the solvent. The same idea can be applied to bases.

How is the Bronsted definition of acids and bases extended by the solvent system definition?

The Bronsted definition is only concerned with proton transfer. The solvent system definition defines an acid as any solute that increases the concentration of the cation generated by auto ionization of the solvent. It defines a base as anything that increases the concentration of the anion generated by the auto ionization of the solvent.

Define and discuss aqua acids.

The acidic proton is on a water molecule coordinated to the central metal ion.

Acid strength increases as the covalent character of the M-O(H2) bond increases. Across the periodic table, covalence increases left to right, and top to bottom. Generally, the heavier the metal, the stronger the acid will be.

Define and discuss hydroxoacids.

The acidic proton is from a hydroxyl group coordinated to a central metal ion, but without a neighboring oxo- group.

Define and discuss oxo-acids.

The acidic proton is from a hydroxyl group coordinated to a central metal ion, with an oxo- group adjacent to it. The stability of the conjugate base is the primary factor in determining acid strength. The conjugate base's stability is determined by resonance, polarizability, and charge density. A rough formula to estimate the pKa is 8 - 5(# of oxo groups).

A few more generalized points about oxo-acids

-Non-metals form acidic oxides, due to electronegativity

-Metals form basic oxides

-Going down a group, atoms are bigger, polarizes water better, and higher hydration enthalpy and entropy

-Amphoteric compounds form weak acids and weak bases. One member of group 2 and several toward the bottom of groups 13, 14, and 15 are amphoteric (Beryllium, aluminum, gallium, indium, germanium, tin, lead, arsenic, antimony, and bismuth).

-Bases donate electron density

-Hard/soft, soft/hard makes the conjugate less stable by making it more soluble

What are some types of non-aqueous solvents?

Liquid ammonia, hydrogen fluoride, anhydrous sulfuric acid, dinitrogen tetroxide

Discuss the Lewis acidity of the different columns of the periodic table.

Alkali metals act as Lewis acids with water, forming hydrated ions.

-Group 13: Boron tri-halides act as Lewis acids. Acid strength increases with atomic mass of the halogen. Aluminum halides form dimers in the gas phase, and are used as Lewis acid catalysts.

-Group 14: Except carbon, group 14 elements exhibit hypervalence and act as Lewis acids, tin(II) chloride acts as both a Lewis acid and base.

-Group 15: Oxides and halides of the heavier elements act as Lewis acids.

-Group 16: Sulfur dioxide can act as a Lewis base when it forms a complex with a trialkylamine that acts as a Lewis base. Sulfur trioxide is a strong Lewis acid.

-Halogens: Bromine and Iodine act as mild Lewis acids.

What's the difference between hard-hard and soft-soft interactions?

Hard preferentially bonds to hard, soft preferentially bonds to soft. H-H and S-S are more thermodynamically stable than H-S or S-H.

Soft-soft is more covalent in character, whereas hard-hard is more ionic in character.

Factors influencing the Gibb's energy of complex formation:

1. Competition with solvent

2. Rearrangement of the substituents of the acid and base that may be necessary to permit formation of the complex

3. Steric repulsion between substituents on the acid and the base

Soft acid/soft base interactions tend to be exothermic

Define reduction, oxidation, reducing agent, and oxidizing agent.

Reduction: The gain of electrons

Oxidation: The loss of electrons

Reducing agent: the species which supplies the electrons (it is oxidized)

Oxidizing agent: the species which receives the electrons (it is reduced)

Together, they make the redox couple.

What is the significance of oxidation numbers?

Oxidation are not tangible, or observable, they are merely a bookkeeping method. They are used to describe the ionic state of an atom.

Why do we write half reactions to express reduction potentials?

Because of the fact that in a redox reaction we are following the movement of electrons, breaking a reaction up into half reactions to separate reduction and oxidation gives us a clear view on the electron transfer, and simplifies balancing a reaction. When written as a complete reaction, the electrons cancel out, and are not written. Remember that all half reactions are written as reductions (in data tables, etc...), so the sign on the reduction potential may need to be reversed.

Why is the electrochemical series so hard to predict?

It is difficult to predict because it involves a number of thermochemical cycles, including heat of sublimation, ionization energy, standard enthalpy of hydration (for the metal and H), bond dissociation energy, and the overall reaction enthalpy.

What is the effect of multiplying a half reaction?

Multiplying the reaction does not change the potential, but it does change the free energy.

What values of K and Eᶲ indicate a spontaneous reaction?

If E° > 0, then K>1, and it is thermodynamically favorable (spontaneous).

Define anode, cathode, and galvanic cell. (if you get this wrong Becky will magically appear and bitch slap you)

Anode: the site of oxidation

Cathode: the site of reduction

Galvanic cell: An electrochemical cell in which a reaction drives an electric current, which travels through an external circuit for measurements.

Write the equation for Gibb's free energy change involving the standard cell potential.

ΔG°=-nFE°

F is the Faraday constant, F=9.648x104C/mol

N is the number of electrons