Chemical Bonds, Interactions, Acids, Bases, and pH

Vocabulary practice flashcards covering chemical bonds, partial charges, acid-base theories, pH calculations, dissociation constants, buffers, and biological pH systems.

Covalent Bond

A chemical bond formed by a shared pair of electrons between atoms, where bonds like C-O, N-C, and O-H are polar, and C-H is non-polar.

Ionic Bond

A chemical bond resulting from the transfer of electron(s) between atoms to form positive and negative ions that attract each other due to opposite charges.

Hydrogen Bond

A slight electrostatic attraction formed when a hydrogen atom covalently linked to an electronegative atom (such as N or O) is attracted to another electronegative atom.

Hydrophobic Effect

The exclusion of non-polar structures by a surrounding network of polar structures in an aqueous solution.

Van der Waals Interactions

Interactions involving attraction between the electron cloud of one atom and the positively charged nucleus of another, as well as repulsion between electron clouds of two atoms.

Partial Atomic Charge

The description of an uneven distribution of electrons in a molecule creating slightly positive ($\delta+$) and negative ($\delta-$) regions on covalently bonded atoms due to electronegativity differences.

Electronegativity

The inherent ability of an atom in a covalent bond to attract shared electrons toward itself.

Molecular Dipole Moment

A vector property created by the presence of partial positive and negative regions within a molecule, rendering the overall molecule polar.

Hydronium Ion

An ion ($\text{H}_3\text{O}^+$) formed when a free proton ($\text{H}^+$) released by an acid attaches to a water molecule ($\text{H}_2\text{O}$) in solution.

Hydroxide Anion

An anion ($\text{OH}^-$) consisting of an oxygen atom covalently bonded to a hydrogen atom, central to the concept of bases.

Arrhenius Acid

A substance that provides hydrogen cations ($\text{H}^+$) when dissolved in water.

Arrhenius Base

A substance that provides hydroxide anions ($\text{OH}^-$) in water or absorbs $\text{H}^+$ cations.

Br(8nsted Acid

A molecule or ion that acts as a proton donor in a chemical reaction ($\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$).

Conjugate Base

The species ($\text{A}^-$) formed after a Br(8nsted acid donates a proton, acting as a proton acceptor.

pH

A measure of acidity defined as the negative logarithm of the hydrogen ion concentration, $\text{pH} = -\log([\text{H}^+])$.

Ion Product Constant of Water ($K_w$)

The equilibrium constant for the self-dissociation of water, equal to $K_w = [\text{H}^+][\text{OH}^-] = 1 \times 10^{-14}\,\text{M}^2$ at $25\,^\circ\text{C}$.

Strong Acid

An acid that dissociates completely into its constituent ions when dissolved in water, such as $\text{HCl}$ and $\text{H}_2\text{SO}_4$.

Weak Acid

An acid that only partially dissociates when dissolved in water, such as acetic acid ($\text{CH}_3\text{COOH}$) and other carboxylic acids.

Dissociation Constant ($K_d$)

The ratio expressed as the product of the concentrations of dissociated products divided by the concentration of the associated starting material ($K_d = \frac{[\text{A}^-][\text{H}^+]}{[\text{HA}]}$).

Henderson-Hasselbalch Equation

An equation relating pH, pK, and component concentrations: $\text{pH} = \text{pK} + \log\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$.

pKa

The negative logarithm of the acid dissociation constant ($\text{pK} = -\log(K)$), where lower values signify stronger acids that release protons readily.

Histidine Side Chain pKa

The pKa value of approximately $6.5$ for the amino acid histidine, indicating that its positively charged state equals its neutral state at that pH.

Protonation State when $\text{pH} < \text{pK}$

The condition in a weak acid system where the concentration of the protonated acid form is greater than the conjugate base ($[\text{HA}] > [\text{A}^-]$).

Protonation State when $\text{pH} > \text{pK}$

The condition in a weak acid system where the concentration of the deprotonated conjugate base is greater than the protonated acid ($[\text{A}^-] > [\text{HA}]$).

Protonation State when $\text{pH} = \text{pK}$

The condition where the concentration of the deprotonated conjugate base equals the protonated acid ($[\text{A}^-] = [\text{HA}]$).

Buffer

A solution consisting of a weak acid and its conjugate base in equilibrium that resists changes in pH by absorbing added $\text{H}^+$ or $\text{OH}^-$.

Buffer Capacity

The concentration or amount of strong acid or base a buffer solution can neutralize before undergoing a significant change in pH.

Buffer Region

The pH range around the pK of a weak acid where buffering action is most effective, highest when the ratio of weak acid to conjugate base is close to $1:1$.

Bicarbonate Buffer System

The primary buffer system in human blood serum composed of carbonic acid ($\text{H}_2\text{CO}_3$) and bicarbonate ($\text{HCO}_3^-$) that maintains blood at $\text{pH} \approx 7.4$.

Open Buffer System

A physiological buffer mechanism where carbonic acid ($\text{H}_2\text{CO}_3$) is regenerated from metabolic waste $\text{CO}_2$ and $\text{H}_2\text{O}$, giving a net $\text{pK} = 6.2$.

Gastric Juice pH

The biological fluid in the human stomach with an acidic pH range between $1.0$ and $2.0$.

Blood Serum pH

The tightly regulated physiological pH level of human blood serum, maintained at approximately $7.4$.

Tissue pH

The physiological pH level of most tissues in the human body, maintained at approximately $7.0$.

Critical Enamel Demineralization pH

The critical threshold value of $\text{pH} = 5.5$ in the oral cavity below which tooth enamel begins to dissolve.

Salivary pH

The pH of human saliva, which normally ranges between $6.4$ and $7.2$ with an average value of $6.8$.

Nocturnal Salivary Acidification

The decrease in salivary pH during sleep caused by slower breathing, which increases $\text{CO}_2$ levels and yields higher carbonic acid concentration.

Acidogenic Bacteria

Oral microorganisms such as Lactobacilli and Streptococci whose growth is favored in low pH environments and which produce additional acids.

Acetate Ion

The conjugate base anion ($\text{CH}_3\text{COO}^-$) formed when acetic acid ($\text{CH}_3\text{COOH}$) donates a proton.

Bicarbonate Ion

The anion ($\text{HCO}_3^-$), also known as hydrogen carbonate, derived from carbonic acid ($\text{H}_2\text{CO}_3$).

Concentration of Water ($[\text{H}_2\text{O}]$)

The molar concentration of liquid water in aqueous solutions, which remains essentially constant at $55\,\text{M}$.