ATOMIC TRENDS

(atomic radius)

(attract electrons to itself)

Acids & Bases

Strong Acids

Hydroiodic acid	HI
Hydrobromic acid	HBr
Hydrochloric acid	HCl
Perchloric acid	HClO4
Sulfuric acid	H2SO4
Nitric acid	HNO3

Strong Bases

Group 1 hydroxides	NaOH, LiOH, KOH
Group 1 oxides	Li2O, etc.
Some group 2 hydroxides	Ba(OH)2, Sr(OH)2, Ca(OH)2
Metal amides	NaNH2, LiNH2, Mg(NH2)2, Ca(NH2)2, etc.

- The conjugate base of a strong acid has no basic properties in water.

- The conjugate base of a weak acid is a weak base.

- The conjugate acid of a strong base has no acidic properties in water. For example, the conjugate acid of LiOH is Li+, which does not act as an acid in water.

- The conjugate acid of a weak base is a weak acid (and the weaker the base, the stronger the conjugate acid). For ex, the conjugate acid of NH3 is NH4+, a weak acid.

Amphoteric: can act as either an ACID or BASE. The conjugate base of a weak polyprotic (more than one H) acid is always amphoteric, because it can either donate or accept another proton.

- ex: H₂CO₃ à HCO₃^- à CO₃^2-

Reaction Rate

1. the lower the activation energy, the faster the reaction rate

2. the greater the concentrations of the reactants, the faster the rate (more favorable collisions)

3. the higher the temperature of the reaction mixture, the faster the reaction rate

ΔG has no bearing on rate of the reaction!

Thermodynamic & Kinetic factors DO NOT AFFECT EACH OTHER!!!

Catalyst - lowers the activation energy of the rate-determining step (therefore the energy of the higher-energy transition state), remains UNCHANGED at the end of the rxn

***even though the catalyst MAY physically change in the middle...will always revert back for no net change

Rate Law - reactants of rate-determining step, NOT products...(previous post)...can only be determined EXPERIMENTALLY

Solutions

Concentration Measurements

- solubility of gases DECREASES with INCREASING temperature (think: soda left out, flat tasting boiled water)

- unlike molaRity, molaLity does NOT change with temperature & pressure (also nearly the same b/c 1L = 1kg)

- mole fraction useful way to express concentration with more than 1 solute is present

Electrolytes

- ionizability factor (i) = tells us how many ions 1 unit of substance will produce in a solution (i=1, covalent and won't dissociate... i = > 1, a strong electrolyte)

7-3) concentration of bromide ions in solution --> concentration (M) = # molecules x molaLity (m)

**Memorize the Solubility Rules!

Phase Solubility Rules

1. The solubility of solids in liquids tends to INCREASE with INCREASING temperature.

2. The solubility of gases in liquids tends to DECREASE with INCREASING temperature.

- ex: soda outside, flat boiled water

3. The solubility of gases in liquids tends to INCREASE with INCREASING pressure.

7-6) Why choice C is the truth, and the above is applied:

Scuba diving safety: The most important issue here is that, in accordance with Henry's law, the increasing pressure correlates to increased absorption of nitrogen. Once the pressure decreases as you ascend, it takes that nitrogen time to get safely released from the body. If we ascend too quickly, the nitrogen bubbles can get too large to safely be eliminated through breathing, and that can have dire consequences (gas bubbles trapped in body). We'll discuss all this in some detail.

Mendelian Genetics

http://mcat-review.org/genetics.php

http://www.merckmanuals.com/professional/sec22/ch327/ch327c.html

pKa (opposite for Ka!!), H-H Titration

Water is not included in the acid-dissociation equilibrium expression because the [H₂O] has no effect on the equilibrium.

As the K_a value of an acid increases, so does the strength of the acid. By definition:

• strong acid: K_a > 1
• weak acid: K_a < 1

The larger the value of pK_a, the smaller the extent of dissociation. A weak acid has a pK_a value in the approximate range −2 to 12 in water. Acids with a pK_a value of less than about −2 are said to be strong acids; a strong acid is almost completely dissociated in aqueous solution, to the extent that the concentration of the undissociated acid becomes undetectable.

Larger the Ka, stronger the acid!

Smaller the pKa, stronger the acid! (THE REVERSE)

pKa + pKb = 14

pH + pOH = 14

pH = -log[H⁺] à [H⁺] = 10^-pH

pOH = -log[OH^-] à [OH^-] = 10^-pOH

^{Henderson-Hasselbalch:}

^{Equivalence point: amount of ACID EQUAL to amount of BASE present, only ions exist in solution}

^{1/2 Equivalence point: (pH = pKa) volume added is half of what it will be at equivalence point & the protonated and deprotonated states are equal}

Type of Titration	When Equivalence Point will occur
Weak Acid w/ STRONG BASE	pH > 7
Weak Base w/ STRONG ACID	pH < 7
STRONG ACID w/ STRONG BASE	pH = 7

The graph above is a titration curve of three different solutions. Solution I is titrated with a base of a lower pH. Rank the solutions in terms of the strength of the acid in the solution.

(a) I < II < III
(b) III < I < II
(c) II < III < I
(d) There is no difference among the strength of the acids.

Explanation: Consider what a titration involves and what this curve tells us about the solutions involved. In this titration we are starting with an acid solution (low pH before titrant is added) and are titrating with a base (high pH after titrant is added). After adding a certain amount, the proportion of acid/base begins to approach 1/1 and the pH will start to increase. As we continue adding base, the pH will eventually level off as the ratio of base far exceeds the ratio of acid.

Using a titrant of lower pH would result in graph I, since the starting point is the same but the end point is lower. Using a higher concentration of acid would result in graph II (compared to graph III), because it takes more titrant to reach the same end point of the titration (the start and end pH values are the same, but the amount of base required is greater; therefore the concentration of starting acid is greater). Thus we can conclude that there is no difference in the strength of the acids; we are looking at three different experimental conditions.

http://www.mcatquestion.com/findquestion.php?arg1=1062

***Compare pH and pKa's quickly:

- Recall that when the pH is below the pK_a of a titratable group, the group will be predominantly protonated

- More detail: Basic groups, like amino groups, have high pK_as, whereas acidic groups like carboxylic acids have low pK_as (recall that K_a = [A^-][H⁺]/[HA]; therefore an acidic group which favors dissociation ([A^-][H⁺]) will have a higher K_a. The pK_a is the -logK_a, so a higher K_a will result in a lower pKa (do the math to compare the pK_as of high and low K_a values).

http://www.mcatquestion.com/findquestion.php?arg1=1051

Back in action, gracias a dios

It astonishes me that it's been nearly a YEAR now since the last time I posted... things really need to start getting back in gear. The beauty of having escaped school, escaped the country, escaped my general essence of being and this whole exam mindset, is that it has helped me re-focus upon what I know I want to do with my life...become a doctor and do all the right things to take me to that place.

In this past year alone I've had perhaps the most distractions I think I've ever had in my life... topics of which I won't go into detail but aside from everything, I'm taking a newfound comfort in this space to reel me back into the direction I should be pointing--and that's a positive attitude and the spirit of learning...ya dig?