Chemistry Mcat

Front 1

Empirical Formula

Back 1

A formula that simply gives the ratio of the number of atoms of each element in a compound.
Example: The empirical formula of hydrogen peroxide, H2O2, is HO.

Front 2

Molecular Formula

Back 2

A formula that gives the actual numbers of atoms of each element in each molecule of a compound.
Example: The molecular formula of hydrogen peroxide is H2O2.

Front 3

Mole

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A mole is a measure of the number of atoms, molecules, or ions there are present. One mole of particles is 6.02 x 1023 particles (Avogadro's number).
Number of moles = mass in grams / molecular weight

Front 4

Density

Back 4

Density = mass / volume

Front 5

Oxidation Number

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The oxidation number of an atom is the charge that atom would carry if the compound it is a part of were composed of ions.

Front 6

Redox Reactions

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When at least one atom undergoes a change in oxidation state, the reaction is a redox reaction.
Oxidation is the loss of electrons, while reduction is the gain of electrons.

Front 7

Metathesis reactions

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reactions in which none of the atoms undergoes oxidation or reduction.
Examples of metathesis reactions are:
Acid-base reactions that involve a transfer of H+,
e.g. NaOH + HCl « NaCl + H2O
Acid-base reactions that involve the sharing of a pair of electrons by a Lewis base and a Lewis acid,
e.g. Cu2+ + 4NH3 « Cu(NH3)42+

Front 8

Balancing Redox Equations

Back 8

Consider the reaction
I3- + S2O32- « I- + S4O62-
This is not balanced. The steps below illustrate how to balance a redox equation.

Determine the oxidation number of each atom on both sides of the equation.

On the left side:
I: -1/3, S: +2, O: -2
On the right:
I: -1, S: +5, O: -2

Determine which atoms are oxidized and which are reduced.

S is oxidized and I is reduced.

Divide the reaction into oxidation and reduction half-reactions and balance these.

Oxidation:
S2O32- ® S4O62- is balanced to 2S2O32- ® S4O62- + 2e-

Reduction:
I3- ® I- is balanced to I3- + 2e- ® 3I-

Combine the half-reactions
I3- + 2S2O32- ® 3I- + S4O62-

Further balance the equation by inspection if necessary.
This is unnecessary in this case.

Front 9

Redox Titration

Back 9

Supposing we had a solution with an unknown concentration of I3- ions. We could determine this concentration by titrating a known volume of the solution against a known concentration of S2O32-. At the point where the color changes, twice as many S2O32- ions have been added as there were I3- ions, as the above example shows. From these data the concentration of I3- ions in the solution can be easily calculated.
This is an example of a redox titration. Redox titrations are similar in principle to acid-base titrations using pH indicators.

Front 10

Oxidizing Agents

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An oxidizing agent is a substance that oxidizes another substance. In doing so, it becomes reduced.
Relatively strong oxidizing agents tend to have a large affinity for electrons. Such agents are:
very electronegative (e.g. F2, Cl2, O2, O3)
or have high oxidation states (e.g. permanganate [MnO4-], chromate [CrO42-], dichromate [Cr2O72-], nitric acid [HNO3], sulfuric acid [H2SO4]).

Front 11

Reducing Agents

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A reducing agent is a substance that reduces another substance. In doing so, it becomes oxidized.
Relatively strong reducing agents are able to lose electrons easily: they have small ionization energies and low electronegativities (e.g. metals such as sodium, magnesium, aluminum, zinc; metal hydrides such as NaH, CaH2).

Front 12

Relative Strengths of Oxidizing and Reducing Agents

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A measure of how easily a substance is reduced is given by its standard state reduction potential, E0. The higher the E0, the more easily it is reduced, and the stronger an oxidizing agent it is. Conversely, the more negative the E0, the more easily it is oxidized, and the stronger a reducing agent it is.

Front 13

Laws of Thermodynamics
1st law:

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Energy is conserved. The total amount of energy in the universe is constant.
The universe can be thought of as simply a system we are interested in and its surroundings. For such a system,
DE (internal energy) = Q (heat) + W (work)
DE = Q + W
E is the internal energy of a system. It is the total of all the possible kinds of energy in the system. E cannot be measured, but DE can.
Q is positive if heat is added to a system from the surroundings and negative if the system gives up heat to its surroundings.
W is positive if work is done on the system by the surroundings (e.g. compressing gas in a container) and negative if a system does work on the surroundings (e.g. if a compressed gas is allowed to expand).

Front 14

Enthalpy

Back 14

Under the condition of constant pressure, Q = DH, the change in enthalpy. Most reactions occur under the essentially constant pressure of 1 atm.
If DH is negative, heat is generated and the reaction is said to be exothermic.
If DH is positive, heat is absorbed from the surroundings and the reaction is said to be endothermic.

Front 15

Laws of Thermodynamics

2nd law:

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Any spontaneous change that occurs must be accompanied by an increase in the entropy of the universe.

Front 16

Laws of Thermodynamics
2nd law:

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Any spontaneous change that occurs must be accompanied by an increase in the entropy of the universe.

Front 17

Laws of Thermodynamics
3rd law:

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The entropy of a pure crystal at 0 K is zero. (Thus the absolute amount of entropy of a system is measurable.)

Front 18

Gibb's Free Energy

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Whether a chemical reaction is spontaneous (i.e. the formation of the products is favored) or not depends on two driving forces: DH and DS. Spontaneous reactions tend to occur when potential energy is released as heat through the breaking and formation of bonds (DH <0), or when the system becomes more disordered (DS >0), or both.
DG incorporates both DH and DS, as show below.
DG = DH - TDS
where G is Gibb's free energy (so named because it is the energy that is free to do work),
H is enthalpy (heat produced or consumed under conditions of constant pressure),
S is entropy (degree of disorder), and
T is temperature in K.
Why is DS multiplied by T? The higher the temperature, the greater the kinetic energy of the molecules and, thus, the greater the tendency for the system to go to a state of increased disorder (an analogy is the increased disorder an earthquake creates).
If DG <0, the forward reaction is spontaneous.
However, even if a reaction is spontaneous, if the activation energy is too high it will not proceed. Such a reaction is said to be under kinetic control and may require the use of a catalyst or energy input such as heat (which would also alter the equilibrium position).
If a reaction occurs as DG predicts, it is said to be under thermodynamic control.

Front 19

Relationship between Gibb's Free Energy and Keq and E (cell potential for a reaction)

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An indication of the equilibrium position of a reaction is provided by each of DG, Keq, and E. DG looks at the thermodynamics of a reaction, while E looks at the relative tendencies of reactants to be oxidized or reduced. Thus they are related to each other.
Keq = e DG°/-RT
Thus if DG is large and negative, Keq will be large.
And, Keq = enFE°/RT
If E is large and positive, Keq will be large.
Also, DG° = -nFE°
Thus, if E is positive, DG will be negative.
(Note: the symbol ° denotes that standard states exist, i.e. the concentration of each reactant and product is 1 M if in solution, and 0.1 MPa if gaseous.)

Front 20

State Functions

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A state function is a quantity whose value depends only on the state of the system and not on its history. X is a state function only if DX does not depend on the path used to go from the initial state to the final state of the system. V, G, P, H, E, S, and T are state functions.
Hess's law is a consequence of enthalpy being a state function. It states that DH of a reaction is the same regardless of whether the reaction occurs in a single step or in several steps.

Front 21

LeChatelier's principle

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When a system at equilibrium is subjected to a stress, it will shift in a direction that minimizes the effect of this stress.

Front 22

First ionization energy

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The energy required to remove the outermost electron from a neutral atom in the gas phase.

Front 23

Electron affinity

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The energy given off when a neutral atom in the gas phase gains an extra electron to form a negatively charged ion.

Front 24

Electronegativity

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The tendency of an atom to draw the electrons in a bond towards it.

Front 25

Relationship between 1st ionization energy, electron affinity and electronegativity

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The first ionization energies, electron affinities, and electronegativities of elements generally increase as one moves closer to the top right-hand corner of the periodic table. They increase as one moves across a row from left to right because the force of attraction between the nucleus and an electron increases as the number of protons in the nucleus increases. They decrease as one moves down a group in the periodic table because the electrons in inner orbitals tend to shield the outer electrons from the attractive force of the nucleus.
Group I elements have large second ionization energies compared with their first ionization energies, while group II elements have second ionization energies not much higher than their first ionization energies. This can be explained by their respective electronic configurations.

Front 26

Formal Charge

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Formal charge is the charge on an atom in its Lewis structure, or the overall charge on an ion or molecule.
Steps in determining formal charge:
1. Divide the electrons in each covalent bond between the atoms in the bond.
2. Determine whether the atom of interest has more electrons than protons (negative formal charge), less (positive formal charge), or the same number (zero formal charge).

Front 27

Colligative Properties

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A colligative property is any property that depends on the concentration of solute particles in a solution but not on their identity. Colligative properties only apply to solutions.

Front 28

Raoult's Law

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P = mole fraction of solvent x Po
where P is the partial pressure of solvent vapor from a solution, Po is the partial pressure of vapor from pure solvent, and the mole fraction of the solvent is concentration of solvent particles / concentration of all particles in a solution.
(Note: Vapor pressure is the pressure of a gas that collects above a liquid in a closed container. The gas and the liquid are in dynamic equilibrium.)
DP = mole fraction of solute x Po

Front 29

Boiling Point Elevation

Back 29

In dilute solutions, DTBP = kb x molality of solute particles
where kb is the molal B.P. elevation constant

Front 30

Freezing Point Depression

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DTFP = -kf x molality of solute particles
where kf is the molal freezing point depression constant

Front 31

Osmotic Pressure

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The osmotic pressure of a solution is the minimum pressure needed to prevent osmosis of pure water across a semipermeable membrane into this solution.
Osmotic pressure = RT [solute particles]

Front 32

Definitions of Acid/Base

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Bronsted-Lowry
Acid Donates proton
Base Accepts proton

Arrhenius
Acid Donates proton
Base Donates OH- (hydroxyl) group

Lewis
Acid Accepts electron pair Base Donates electron pair

Front 33

pH and pOH

Back 33

pH = -log10[H+]
pOH = -log10[OH-]

Front 34

Kw

Back 34

Water dissociates as follows:
H2O « H+ + OH-
Kw = [H+] [OH-] = 10-14 (at 25oC)
\ pKw = pH + pOH = 14
(e.g. A solution with pH=9 has pOH = 14 - 9 = 5)
\ In neutral water, [H+] = [OH-] = Ö(10-14) = 10-7
pH = -log10[H+] = 7.0
pOH = -log10[OH-] = 7.0

Front 35

Ka

Back 35

An acid dissociates in water as follows:
HA « H+ + A-
Ka = [H+] [A-] / [HA]
pKa = -log10Ka

Front 36

Kb

Back 36

A base reacts with water as follows:
A- + H2O « HA + OH-
Kb = [HA] [OH-] / [A-]
(Note: the concentration of H2O is usually ignored)
pKb = -log10Kb
Note, Ka Kb = ([H+] [A-] / [HA]) x ([HA] [OH-] / [A-])
= [H+] [OH-]
= Kw
= 10-14
And since Ka Kb = 10-14
pKa + pKb = 14 = pKw

Front 37

Henderson-Hasselbalch Equation

Back 37

Ka = [H+] [A-] / [HA]
[H+] = Ka [HA]/[A-]
pH = pKa - log10[HA]/[A-]
pH = pKa + log10[A-]/[HA]

Front 38

Buffers

Back 38

Buffers are solutions that resist changes in pH when small amounts of acid or base are added to it.

A buffer usually consists of:
a weak acid and its conjugate base or a weak base and its conjugate acid.

Such systems have the greatest buffering capability when the two components are in equal concentration.
(e.g. an acetic acid/acetate buffer has greatest buffering capability when [acetic acid] = [acetate], and therefore when
pH = pKa + log10[acetate]/[acetic acid]
= pKa + log10(1)
= pKa + 0
= pKa = 4.74)

Front 39

Types of Bonds

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Ionic
Strongest
Electrostatic attraction between positive and negative ions

Covalent
Strong
Sharing of electrons between atoms: the nuclei of both atoms are pulled toward the shared electrons


Dipole-dipole
Weak to moderate
Electrostatic attraction between partially positive and partially negative parts of neutral molecules
van der Waals (or London forces or dispersion forces)Weakest Transient dipoles random accumulation of electrons on one side of a molecule creates a dipole moment, which may in turn induce dipoles in other molecules

Metallic
Moderate to strong
Positive nuclei in a "sea" of mobile electrons