SBI 4U1 Biochemistry: Acids, Bases, Buffers, Water Properties, and Intermolecular Forces

Comprehensive vocabulary flashcards covering acids, bases, biological buffer systems, properties of water, molecular geometry, and intermolecular forces from SBI 4U1 Biochemistry notes.

Weak Acid

An acid that only partially ionizes in water and can establish a reversible equilibrium reaction, such as acetic acid ($CH_3COOH_{(aq)}$).

Weak Base

A base that only partially ionizes in water and participates in a reversible equilibrium reaction, such as ammonia ($NH_3(aq)$).

Equilibrium Reaction

A reversible chemical reaction represented by doubled-up forward-and-backward arrows ($\rightleftharpoons$) that can shift in either direction depending on the availability of $H^+$ and $OH^-$ ions.

Acetic Acid ($CH_3COOH_{(aq)}$)

A weak acid found in vinegar that reacts reversibly with water to form hydronium and acetate ions.

Acetate Ion ($CH_3COO^-$)

The conjugate base formed when acetic acid loses a hydrogen ion ($H^+$).

Conjugate Acid-Base Pair

Two chemical species with the same overall chemical structure but differing ionic charges, related by the loss or gain of a single hydrogen ion ($H^+$).

Ammonia ($NH_3(aq)$)

A weak base that reacts reversibly with water to form ammonium and hydroxide ions.

Ammonium Ion ($NH_4^+$)

The conjugate acid formed when an ammonia molecule accepts a hydrogen ion ($H^+$).

Buffer

A chemical substance or system that equalizes acidity or basicity in aqueous solutions and resists severe changes in the pH of blood, body fluids, or tissues.

Carbonic Acid-Bicarbonate Buffer System

The primary buffer system in living tissues represented by the equilibrium $H_2O + CO_2(aq) \rightleftharpoons H_2CO_3 \rightleftharpoons HCO_3^- + H^+_{(aq)}$.

Carbonic Acid ($H_2CO_3$)

An intermediate molecule in the bicarbonate buffer system found in low levels in all tissues.

Bicarbonate Ion ($HCO_3^-$)

The conjugate base in the blood buffer system whose excess accumulation is regulated by the kidneys releasing more into urine.

Acidemia

A minor acidic condition in body tissues that causes the buffer equilibrium to shift towards the left, producing more water and carbon dioxide to increase basicity.

Alkalemia

A minor basic condition in body tissues that causes the buffer equilibrium to shift towards the right to generate more free $H^+$ ions and increase acidity.

Acidosis

A major systemic clinical issue resulting from significant and uncorrected acidity in the body.

Alkalosis

A major systemic clinical issue resulting from significant and uncorrected basicity in the body.

Carbon Dioxide Exhalation

The respiratory process of exhaling excess $CO_2$ to counteract an acidic localized environment in the body.

Urinary $H^+$ Ion Elimination

The process by which the kidneys counteract systemic acidity by urinating out accumulating hydrogen ions.

Acid

A compound that ionizes in water to form $H^+$ ions, acts as an $H^+$ donor, and has a pH of less than $7$.

Base

A compound that ionizes in water to form $OH^-$ ions, acts as an $H^+$ acceptor, and has a pH of more than $7$.

Neutralization Reaction

A reaction in which an acid and a base react to form a salt and water molecules (e.g., $HCl + NaOH \rightarrow NaCl + H_2O$).

Salt

An ionic compound, such as sodium chloride ($NaCl$), formed alongside water during a neutralization reaction.

Litmus Paper

An experimental paper indicator used to detect whether a solution is acidic or basic.

Phenolphthalein

A liquid chemical indicator used experimentally to determine the acidic or basic nature of a solution.

Bromothymol Blue

A liquid chemical indicator used to experimentally determine or estimate the pH of aqueous solutions.

pH Formula

The mathematical calculation of acidity defined as $pH = -\log[H^+]$ or more specifically $pH = -\log_{10}[H_3O^+_{(aq)}]$.

Human Blood pH

The steady, homeostatically regulated pH range of human blood, which is $7.35$ to $7.45$.

Autoionization of Water

The reversible reaction in which two water molecules react to form hydronium and hydroxide ions: $H_2O + H_2O \rightleftharpoons H_3O^+ + OH^-$.

Hydronium Ion ($H_3O^+$)

The ion that forms when hydrogen ions ($H^+$) are diluted in water, often used interchangeably with $H^+$ to simplify equations.

Hydroxide Ion ($OH^-$)

The negatively charged ion ($OH^-$) produced by bases in water and present alongside hydronium to keep pure water neutral.

Strong Acid

An acid, such as hydrochloric acid ($HCl$), that ionizes completely in water without undergoing equilibrium, yielding no $OH^-$ ions.

Hydrochloric Acid ($HCl$)

A strong acid that undergoes complete ionization in water: $HCl + H_2O \rightarrow H_3O^+_{(aq)} + Cl^-_{(aq)}$.

Strong Base

A base, such as sodium hydroxide ($NaOH$), that ionizes completely in water with a one-sided arrow, producing no $H^+$ or $H_3O^+$ ions.

Sodium Hydroxide ($NaOH$)

A strong base that ionizes completely in water to produce sodium ions ($Na^+_{(aq)}$) and hydroxide ions ($OH^-_{(aq)}$).

Cohesion

The property of water molecules forming hydrogen bonds with each other, resulting in high surface tension.

Adhesion

The property of water molecules forming hydrogen bonds with other polar molecules, contributing to capillary action and solubility.

Surface Tension

The cohesive force at the surface of a liquid that allows organisms like a water strider to walk on the surface of a pond.

Capillary Action

The movement of water upward through narrow spaces, such as xylem tubes in plants, driven by adhesion and cohesion.

High Specific Heat Capacity

The property caused by hydrogen bonding where water absorbs or loses large amounts of thermal energy with relatively small changes in temperature.

Temperature Moderation

The biological effect of water's high specific heat capacity that helps organisms maintain a constant body temperature.

High Specific Heat of Vaporization

The property where liquid water requires large amounts of thermal energy to break hydrogen bonds and become a gas.

Evaporative Cooling

The process by which organisms, including humans, dissipate body heat through the evaporation of water from body surfaces (such as sweating).

Ice Water Lattice

The open crystalline framework formed below $0^\circ\text{C}$ where hydrogen bonds keep water molecules spread apart, reducing ice's density below that of liquid water.

Density of Water at $4^\circ\text{C}$

The temperature at which liquid water achieves its highest density before expanding as it cools toward freezing.

Top-Down Freezing

The survival mechanism for aquatic life where ice floats and freezes on top of a body of water, insulating the liquid water beneath during winter.

Snow Insulation

The property of solid crystallized water acting as a very good thermal insulator to protect many organisms from extreme winter cold.

Lewis Dot Diagram

A representation of an atom or molecule that shows only the valence electrons surrounding the atomic symbol.

Water Molecule Structure

A molecule consisting of one highly electronegative oxygen atom bonded to two least electronegative hydrogen atoms, bearing a relatively negative side and a relatively positive side.

Angular (Bent) Shape

The geometric shape of a water molecule ($H_2O$) containing two polar bonds and non-bonding lone pairs on the central oxygen.

Methane ($CH_4$)

A non-polar molecule with a tetrahedral shape whose individual polar $C-H$ bond vectors cancel out due to symmetry.

Tetrahedral Shape

The symmetrical four-cornered molecular geometry exhibited by methane ($CH_4$).

Hydrogen Chloride ($HCl$ Molecule)

A polar molecule with a linear shape containing a single polar bond between hydrogen and chlorine.

Linear Shape

The straight-line geometric shape exhibited by diatomic molecules like hydrogen chloride ($HCl$).

Ammonia ($NH_3$ Molecule)

A polar molecule with a trigonal pyramid shape containing three polar $N-H$ bonds and a lone pair on nitrogen.

Trigonal Pyramid Shape

The molecular geometry of ammonia ($NH_3$) resulting from three bonding pairs and one non-bonding lone pair on the central nitrogen atom.

Nitrogen Valence Electrons

The $5$ valence electrons present in a nitrogen atom that govern its bonding and molecular shape in ammonia.

Intermolecular Forces

Forces of attraction that occur between separate molecules and are responsible for determining the physical state of substances.

Van der Waals Forces

A collective term used to describe intermolecular forces of attraction, including London forces and dipole-dipole forces.

London Forces

The weakest intermolecular forces that occur between all molecules as a result of temporary induced dipoles.

Temporary Dipole

A transient separation of charge created when electrons in two adjacent atoms momentary occupy positions that induce attractive poles.

Induced Dipole-Induced Dipole Attraction

An alternate name for London forces, referring to temporary dipoles created in adjacent molecules.

Condensation of Nonpolar Substances

The process driven by London forces that causes nonpolar substances to condense into liquids and freeze into solids when temperatures drop low enough.

Dipole-Dipole Forces

Medium-strength intermolecular forces occurring exclusively between polar molecules that possess permanent net dipoles.

Permanent Net Dipole

A lasting separation of electric charge across a polar molecule that allows electrostatic attraction between neighboring molecules.

Electrostatic Attraction in Dipoles

The attraction where the partial positive charge ($\delta^+$) on one molecule attracts the partial negative charge ($\delta^-$) on an adjacent molecule.

Hydrogen Bond

The strongest type of intermolecular force, occurring between molecules where hydrogen is covalently bonded to fluorine, oxygen, or nitrogen.

H-Bonding Partner Elements

The three highly electronegative elements—fluorine, oxygen, and nitrogen—capable of forming hydrogen bonds when covalently attached to hydrogen.

Cancelled-Out $\Delta E$ Vectors

The physical explanation for why a molecule like methane ($CH_4$) is non-polar overall despite containing individual slightly polar bonds.

Polar Covalent Bond

A covalent bond formed by the unequal sharing of an electron pair, giving one atom a partial positive charge ($\delta^+$) and the other a partial negative charge ($\delta^-$).

Delta ($\delta$)

The lower-case Greek letter used to denote a partial or relative electric charge at one end of a polar covalent bond.

Partially Positive ($\delta^+$)

The relative positive electrical charge assigned to the less electronegative atom in a polar covalent bond.

Partially Negative ($\delta^-$)

The relative negative electrical charge assigned to the more electronegative atom in a polar covalent bond.

Electronegativity Difference ($\Delta E$)

The numerical difference in electronegativity values between two bonded atoms, used to assess bond type along a polarity continuum.

Nonpolar Covalent Bond Range

The range of electronegativity differences ($\Delta E$) from $0$ to $0.4$.

Polar Covalent Bond Range

The range of electronegativity differences ($\Delta E$) greater than $0.4$ and less than $1.7$.

Ionic Bond Range

The range of electronegativity differences ($\Delta E$) from $1.7$ up to $4.0$ where complete electron transfer occurs.

Carbon-Hydrogen ($C-H$) Bond Polarity

A bond with an electronegativity difference of $2.5 - 2.1 = 0.4$, classifying it as a nonpolar covalent bond.

Carbon-Oxygen ($C-O$) Bond Polarity

A bond with an electronegativity difference of $3.5 - 2.5 = 1.0$, classifying it as a polar covalent bond.

Sodium Chloride ($NaCl$) Bond Polarity

A chemical bond with an electronegativity difference of $3.0 - 0.9 = 2.1$, classifying it as an ionic bond.

Polar Molecule

A molecule with an unequal distribution of electrical charge resulting in a relatively positive end and a relatively negative end.

Dipole Moment

A measure of the overall polarity and separation of opposite charges across a molecule.

Like Dissolves Like

The solubility principle explaining that polar molecules dissolve in polar solvents (like water) because of matching polarities.

Molecular Asymmetry

The asymmetrical spatial arrangement of polar bonds or dipoles that prevents them from cancelling out, creating an overall polar molecule.

Orbital

A region in space occupied by electrons around an atomic nucleus.

Hybridization

The process where valence electron orbitals overlap and change their orientation when two or more atoms form covalent bonds.

VSEPR Theory

Valence Shell Electron Pair Repulsion theory; a model used to describe hybrid orbitals and molecular shapes by accounting for lone pairs and bond polarity.

Non-bonding (Lone) Pair

Electrons that orbit an atom but do not take part in covalent chemical bonding.

Central Atom Lone Pair Repulsion

The phenomenon where non-bonding pairs on a central atom take up more spatial volume than bonding pairs, repelling and compressing bond angles.

Bond Angle Compression

The reduction in angle between bonding pairs caused by the repulsion and extra volume occupied by lone pairs.

Intramolecular Bond

A chemical bond within a molecule, such as the covalent bond holding hydrogen and chlorine together in $HCl$.

Polarity Continuum

The continuous spectrum of electronegativity differences ($\Delta E$) spanning from nonpolar covalent to polar covalent to ionic bonding.

Concentration of Hydronium ($[H_3O^+]$)

The molar concentration of hydronium ions in an aqueous solution, denoted by square brackets in pH calculations.

Concentration of Hydroxide ($[OH^-]$)

The molar concentration of hydroxide ions in an aqueous solution, denoted by square brackets $[OH^-]$.

Xylem Tubes

Vascular tubes in plants through which water moves upward via capillary action.

Water Strider Pond Walking

An ecological example of water's high surface tension, allowing an insect to rest and walk on water surfaces.

Sugar Solubility in Water

An example of water acting as a versatile solvent where polar compounds dissolve rapidly due to adhesion and hydrogen bonding.

Sweating

A physiological mechanism in humans and other organisms that dissipates body heat through the high specific heat of vaporization of water.

Biological Preference for Weak Acids/Bases

The biological principle that living systems rely predominantly on weak acids and bases and buffer equilibriums rather than strong acids or bases.

Electronegativity of Hydrogen

The electronegativity value of hydrogen ($2.1$), which makes it relatively electropositive when bonded to carbon, nitrogen, oxygen, or chlorine.

Electronegativity of Carbon

The electronegativity value of carbon ($2.5$), used to calculate bond polarities in organic molecules.