Honors Chemistry Semester 2 Final Exam Study Guide Flashcards

Comprehensive vocabulary flashcards covering chemical reactions, stoichiometry, thermochemistry, bonding, gas laws, and solutions based on the Semester 2 Honors Chemistry Final Exam Study Guide.

(s), (l), (g), and (aq)

Symbols in a chemical equation representing solid, liquid, gas, and aqueous (dissolved in water) states.

Reactant vs. Product

A reactant is a starting substance in a chemical reaction, while a product is a substance formed during the reaction.

Chemical Equation Arrow

The symbol in a chemical equation that means "yields" or "produces."

Coefficient

The number that can be changed when balancing a chemical equation; it is the number placed in front of a chemical formula.

Synthesis Reaction

A reaction where two or more substances combine to form a single, more complex product.

Decomposition Reaction

A reaction where a single reactant breaks down into two or more products.

Single Replacement Reaction

A reaction where one element replaces another element in a compound.

Double Replacement Reaction

A reaction where the ions of two compounds exchange places in an aqueous solution to form two new compounds.

Combustion Reaction

A reaction where a substance reacts with oxygen, often producing $CO_2$ and $H_2O$, characterized by the release of energy.

Activity Series

A list that provides information on whether a specific metal will replace another in a single replacement reaction based on reactivity.

Molar Mass

The mass in grams of one mole of a substance, found using the periodic table.

Mole Ratio

A conversion factor derived from the coefficients of a balanced chemical equation.

STP (Standard Temperature and Pressure)

Conditions under which $1$ mole of any gas occupies a volume of $22.4\,L$.

Limiting Reactant

The reactant that is completely consumed first in a reaction, determining the maximum amount of product that can be formed.

Excess Reactant

The reactant that is not completely used up in a chemical reaction.

Theoretical Yield

The maximum amount of product that can be produced from a given amount of limiting reactant, determined by calculation.

Percent Yield Formula

$$\frac{\text{Actual Yield}}{\text{Theoretical Yield}} \times 100$$

Exothermic Reaction

A reaction that releases heat to its surroundings, characterized by a negative $\Delta H$.

Endothermic Reaction

A reaction that absorbs heat from its surroundings, characterized by a positive $\Delta H$.

Heat Equation

$q = mc\Delta T$, where $q$ is heat, $m$ is mass, $c$ is specific heat capacity, and $\Delta T$ is change in temperature.

Heating Curve Flat Sections

Parts of a heating curve representing phase changes, where temperature stays constant while heat is added.

Heating Curve Slanted Sections

Parts of a heating curve representing the heating of a single phase, where temperature increases.

Hess’s Law

A law stating that the total enthalpy change for a reaction is the same regardless of the number of steps it takes.

Bond Enthalpy

The energy required to break a chemical bond (energy absorbed) or the energy released when a bond is formed.

Enthalpy Calculation from Bonds

$$\Delta H = \text{bonds broken} - \text{bonds formed}$$

Valence Electrons

The electrons in the outermost energy level of an atom that are involved in chemical bonding.

Octet Rule

The tendency of atoms to prefer to have eight electrons in the valence shell.

Lone Pair

A pair of valence electrons that are not shared with another atom in a covalent bond.

Kinetic Molecular Theory (KMT)

A model describing gas particles as being in constant, random motion with kinetic energy proportional to Kelvin temperature.

Kelvin Temperature Conversion

$K = ^\circ\text{C} + 273$; gas law temperatures must always be in Kelvin.

Boyle’s Law

$P_1V_1 = P_2V_2$; describes the inverse relationship between pressure and volume.

Charles’s Law

$\frac{V_1}{T_1} = \frac{V_2}{T_2}$; describes the direct relationship between volume and Kelvin temperature.

Gay-Lussac’s Law

$\frac{P_1}{T_1} = \frac{P_2}{T_2}$; describes the direct relationship between pressure and Kelvin temperature.

Combined Gas Law

$$\frac{P_1V_1}{T_1} = \frac{P_2V_2}{T_2}$$

Ideal Gas Law

$PV = nRT$, where $P$ is pressure, $V$ is volume, $n$ is moles, $R$ is the ideal gas constant ($0.0821\,L\cdot\text{atm}/\text{mol}\cdot K$), and $T$ is Kelvin temperature.

Dalton’s Law of Partial Pressures

The total pressure of a mixture of gases is equal to the sum of the partial pressures of the individual gases.

Graham’s Law

A law that compares the rates of effusion or diffusion of two gases based on their molar masses.

Triple Point

The temperature and pressure at which the solid, liquid, and gas phases of a substance coexist in equilibrium.

Solute and Solvent

The solute is the substance being dissolved, and the solvent is the substance (usually liquid) that does the dissolving.

Saturated Solution

A solution containing the maximum amount of solute that can be dissolved at a given temperature.

Supersaturated Solution

A solution that contains more dissolved solute than a saturated solution could normally hold at the same temperature.

Molarity (M)

The concentration of a solution expressed as moles of solute per liter of solution ($M = \frac{n}{V}$).

Neutralization Reaction

A reaction between an acid and a base that produces water and a salt.

pH and pOH Relationship

$\text{pH} + \text{pOH} = 14$ at $25^\circ\text{C}$; pH identifies if a solution is acidic ($<7$), neutral ($7$), or basic ($>7$).

Titration

A laboratory procedure in which a solution of known concentration is used to determine the concentration of an unknown solution.