Lecture Stoichiometry and Acids and Bases

Flashcards covering stoichiometry, water chemistry, aqueous solutions, equilibrium, solubility, and acid-base chemistry based on lecture notes.

Moles (from mass)

Calculated using the formula $n = \frac{m}{M}$, where moles equals mass divided by molar mass.

Moles (from concentration)

Calculated using the formula $n = c \times v$, where moles equals concentration times volume.

Limiting Reactant

The reactant that produces the smaller amount of moles; determined by converting both reactants to moles and comparing.

Water Polarity

Occurs because oxygen is more electronegative than hydrogen, resulting in a $\delta^-$ on oxygen and a $\delta^+$ on hydrogens.

Hydrogen Bonding

An interaction that must include $O-H$, $N-H$, or $F-H$ bonds; contributing to high boiling and melting points.

Strong electrolyte

A substance that $100\%$ ionises in solution, such as $NaCl$ and $HCl$.

Weak electrolyte

A substance that partially ionises in solution, such as $CH_3COOH$ and $NH_3$.

Non-electrolyte

A substance that does not ionise in solution, such as sucrose or ethanol.

Dynamic Equilibrium

A state where the reaction never stops, the forward rate equals the reverse rate, and concentrations stay constant.

Reaction Quotient ($Q$)

Represents the current position of a reaction.

Equilibrium Quotient ($K$)

Represents the final destination situation of a reaction.

The Golden Rule for Equilibrium direction

If $Q < K$, the reaction moves right; if $Q > K$, the reaction moves left; if $Q = K$, it is at equilibrium.

ICE Table

Used for reversible reactions to find $K_c$ when starting or equilibrium concentration amounts are given.

Equilibrium Constant Expression ($K$)

The ratio of products on top and reactants on bottom, where coefficients become powers; solids are never included.

$$Ksp$$

The solubility constant product; a large value means highly soluble and a small value means insoluble. It only changes with temperature.

Saturated Solution Rule

If a solid is present, the ion concentration stays fixed.

Bronsted Lowry Acid

A proton donor; the strongest acid is the best donor.

Bronsted Lowry Base

A proton acceptor.

Conjugate Pairs

Species that differ by exactly one proton, such as $CH_3COOH$ and $CH_3COO^-$.

Amphiprotic

A substance that can act as either an acid or a base.

Weak Acid pH Formula

$$[H_3O^+] = \sqrt{K_a \times \text{concentration}}$$

Weak Base pH Formula

$$[OH^-] = \sqrt{K_b \times \text{concentration}}$$

Buffer

A mixture of a weak acid and its conjugate base.

Henderson Hasselbach Equation

The equation used for buffers: $pH = pK_a + \log \left( \frac{[A^-]}{[HA]} \right)$. If $[A^-] = [HA]$, then $pH = pK_a$.

Equivalence Point

The point in a titration where all of the original acid has reacted; for normal acids, there is one, but amino acids have multiple corresponding to their ionisable groups.

Half Equivalence Point

The point where $[A^-] = [HA]$ and $pH = pK_a$ because $\log(1) = 0$.

Isoelectric Point ($pI$)

The pH where the molecule has a net charge of $0$; calculated for amino acids by averaging the $pK_a$ values surrounding the neutral form.

Zwitterion

A molecule that has a positive and negative net charge at the same time, though the overall charge is $0$.

Hydrophilic

Water loving molecules that are charged and polar.

Hydrophobic

Water hating molecules, such as non-polar hydrocarbons.

pH vs pKa condition (HA)

When $pH < pK_a$, the molecule exists in the $HA$ form.

pH vs pKa condition ($A^-$)

When $pH > pK_a$, the molecule exists in the $A^-$ form.