Fundamentals of Chemistry: Quantities, Reactions, Gases, Solutions, and Acids-Bases

Comprehensive vocabulary flashcards covering chemical reactions and quantities, properties of gases and gas laws, solution behavior and concentrations, and acids, bases, and buffer systems.

Counting term

A word or expression that states a specific number of items grouped together, utilized when items are too small or numerous to count one-by-one.

Mole

A counting unit used to quantify atoms, ions, and molecules, defined by the equality $1\text{ mole} = 6.02 \times 10^{23}\text{ particles}$.

Avogadro’s number

The exact quantity of particles found in one mole of any chemical substance, which is $6.02 \times 10^{23}\text{ particles}$.

Molar mass

The mass in grams of $1\text{ mole}$ of an element or compound, serving as a conversion factor between grams and moles.

Chemical change

A process where a substance is converted into one or more new substances with different formulas and properties, commonly evidenced by bubbles, color change, solid formation, or thermal energy changes.

Reactants

The original starting substances consumed in a chemical reaction, positioned on the left side of a chemical equation.

Products

The new chemical substances produced during a chemical reaction, positioned on the right side of a chemical equation.

Balanced chemical equation

A representation of a chemical change where the total number of atoms for each element is identical on both the reactant side and the product side.

Energy of activation

The minimum kinetic energy that colliding reactant molecules must possess in order to successfully break and form bonds to react.

Exothermic reaction

A chemical reaction that releases heat energy, where the energy of the products is lower than the energy of the reactants and heat appears on the product side.

Endothermic reaction

A chemical reaction that absorbs heat energy, where the energy of the products is greater than the energy of the reactants and heat is supplied as a reactant.

Kinetic molecular theory (KMT)

A model stating that gas particles move constantly, randomly, and rapidly in straight lines, have negligible volume compared to container volume, experience no attractive or repulsive forces, and exhibit kinetic energies that rise with temperature.

Gas pressure ($P$)

A measure of the collisions that gas particles make against the walls of their container; the force exerted per unit area by the gas on the container.

Systolic pressure

The peak blood pressure exerted against arterial walls when the heart contracts to push blood out; normal values range from $100\text{ to }120\,\text{mmHg}$.

Diastolic pressure

The lowest blood pressure exerted on arterial walls while the heart muscle relaxes and refills with blood; normal values range from $60\text{ to }80\,\text{mmHg}$.

Sphygmomanometer

A medical instrument featuring an inflatable cuff and pressure gauge used to assess arterial blood pressure by listening for pulse sounds as cuff pressure falls.

Boyle’s law

The gas principle demonstrating that at constant temperature ($T$) and molar quantity ($n$), the pressure ($P$) of a gas varies inversely with its volume ($V$), expressed as $P_1 \times V_1 = P_2 \times V_2$.

Inspiration

The process of inhalation during which the diaphragm contracts and moves downward, expanding lung volume and dropping internal lung pressure below atmospheric pressure so air enters.

Expiration

The process of exhalation during which the diaphragm relaxes and curves upward, reducing lung volume and elevating internal pressure so air flows outward.

Gay-Lussac’s law

The gas principle stating that when container volume ($V$) and mole count ($n$) remain constant, gas pressure ($P$) is directly proportional to its Kelvin temperature ($T$), expressed as $\frac{P_1}{T_1} = \frac{P_2}{T_2}$.

Partial pressure

The independent pressure exerted by a particular component gas within a mixture as if it alone occupied the entire volume.

Dalton’s law of partial pressures

The gas law establishing that total pressure of a gas mixture equals the sum of the individual partial pressures of every component gas ($P_T = P_1 + P_2 + P_3 + \dots$).

Solution

A homogeneous mixture comprising two or more distinct substances that can exist stably in solid, liquid, or gaseous physical states.

Solute

The dissolved component within a solution, typically present in a smaller proportion than the solvent.

Solvent

The substance in a solution present in the largest quantity that acts as the dissolving medium for solutes.

“Like dissolves like”

The chemical solubility rule establishing that solutions form between substances having similar polarities (e.g., polar solutes in polar solvents, nonpolar solutes in nonpolar solvents).

Electrolytes

Ionic compounds or acids and bases that dissociate into free ions upon dissolving in water, producing a solution capable of conducting electricity.

Strong electrolyte

A solute that dissociates $100\%$ into positive and negative ions in aqueous solution, conducting electric current strongly enough to illuminate a light bulb brightly.

Weak electrolyte

A solute that only slightly dissociates into ions in water, leaving the majority of solute as intact neutral molecules and conducting electricity poorly.

Nonelectrolyte

A solute that dissolves purely as intact molecules in water without generating ions, rendering the resulting solution nonconductive to electric current.

Concentration

A quantitative description of a solution reflecting the specific ratio of solute quantity to the total quantity of the solution.

Mass percent ($\%\text{ m/m}$)

A concentration unit determined by calculating $\frac{\text{g of solute}}{\text{g of solute} + \text{g of solvent}} \times 100\%$.

Volume percent ($\%\text{ v/v}$)

A concentration unit determined by calculating $\frac{\text{mL of solute}}{\text{mL of solution}} \times 100\%$.

Mass/volume percent ($\%\text{ m/v}$)

A concentration expression determined by calculating $\frac{\text{g of solute}}{\text{mL of solution}} \times 100\%$.

Molarity ($M$)

A concentration metric expressing the number of moles of solute dissolved per liter of solution ($M = \frac{\text{moles of solute}}{\text{liter of solution}}$).

Osmosis

The net spontaneous movement of water through a semipermeable membrane from an area of lower solute concentration toward an area of higher solute concentration.

Isotonic solution

A fluid environment possessing an equal solute concentration relative to red blood cells ($5\%\,\text{(m/v) glucose}$ or $0.9\%\,\text{(m/v) NaCl}$), maintaining steady cell volume without net fluid shift.

Hypotonic solution

A solution that has a lower solute concentration than the cytoplasm of a red blood cell ($< 5\%\text{ glucose}$ or $< 0.9\%\text{ NaCl}$), causing osmotic water entry into the cell.

Hemolysis

The swelling and bursting of red blood cells caused by excess water uptake when suspended in a hypotonic medium.

Hypertonic solution

A solution with a higher solute concentration than the cellular fluid of red blood cells ($> 5\%\text{ glucose}$ or $> 0.9\%\text{ NaCl}$), prompting water to exit the cell via osmosis.

Crenation

The shrinking and distortion of red blood cells that occurs when placed into a hypertonic solution as water leaves the cell.

Arrhenius acid

A chemical substance that dissociates in water to yield hydrogen ions ($\text{H}^+$ or protons), exhibits a sour taste, and acts as an electrolyte.

Arrhenius base

An ionic compound that dissociates in water to yield hydroxide ions ($\text{OH}^-$), feels slippery or tastes bitter and chalky, and functions as an electrolyte.

Brønsted–Lowry acid

Any chemical substance capable of donating a hydrogen ion ($\text{H}^+$, or proton) to a base during a reaction.

Brønsted–Lowry base

Any chemical species that accepts a hydrogen ion ($\text{H}^+$, or proton) from an acid in a chemical reaction.

Conjugate acid–base pair

Two molecular or ionic formulas that differ from one another solely by the loss or gain of a single hydrogen ion ($\text{H}^+$).

Conjugate base

The chemical species produced after a Brønsted–Lowry acid releases and loses its proton ($\text{H}^+$).

Conjugate acid

The chemical species generated once a Brønsted–Lowry base gains and bonds to an incoming proton ($\text{H}^+$).

Strong acid

An acid that ionizes completely ($100\%$) in water into hydronium ions and conjugate base ions (examples include $\text{HI}$, $\text{HBr}$, $\text{HClO}_4$, $\text{HCl}$, $\text{H}_2\text{SO}_4$, and $\text{HNO}_3$).

Weak acid

An acid that ionizes only slightly in water, maintaining an equilibrium composed primarily of undissociated molecular acid.

Strong base

A metal hydroxide formed from Group 1A or Group 2A metals that dissociates completely into metal cations and hydroxide ions in aqueous solution.

Weak base

A base formed from transition metals or aluminum and hydroxide, or molecular bases like ammonia, that dissociate or ionize only partially in water.

$\text{pH}$ scale

A quantitative scale spanning from $0$ to $14$ indicating acidity, wherein $\text{pH} < 7.0$ is acidic, $\text{pH} = 7.0$ is neutral, and $\text{pH} > 7.0$ is basic.

Acidosis

A clinical condition where blood $\text{pH}$ drops below its normal physiological window ($< 7.35$), associated with conditions such as diabetic ketoacidosis or respiratory failure.

Alkalosis

A clinical condition where blood $\text{pH}$ rises above its normal physiological window ($> 7.45$), caused by excessive loss of stomach acid from vomiting or hyperventilation.

Buffer solution

An aqueous system consisting of a weak acid and a salt of its conjugate base that resists drastic shifts in $\text{pH}$ when moderate amounts of acid or base are introduced.