Chemistry 111 Final

Matter Classification

The categorization of matter into pure substances and mixtures.

Pure substances

Material with a fixed composition, which can be elements or compounds.

Mixtures

Combination of two or more substances that retain their individual properties; can be homogeneous or heterogeneous.

Physical properties

Characteristics that can be observed without changing the substance's identity, such as boiling point or color.

Chemical properties

Characteristics that become evident during a chemical reaction, such as reactivity or flammability.

Physical change

A change that alters a substance's appearance but not its composition, e.g., melting ice.

Chemical change

A change that results in the formation of new chemical substances, e.g., burning wood.

Intensive properties

Properties that remain the same regardless of the amount of substance, such as density.

Extensive properties

Properties that depend on the amount of substance present, such as mass or volume.

Scientific notation

A way to express numbers as a product of a number between 1 and 10 and a power of ten.

SI Units

Standard units of measurement used in science, including kilogram (kg) for mass and meter (m) for length.

States of Matter

The distinct forms that different phases of matter take on, primarily solid, liquid, and gas.

Significant Figures

Digits in a number that are significant to its precision, indicating the accuracy of measurement.

Molar Mass

The mass of one mole of a substance, typically expressed in grams per mole (g/mol).

Intermolecular Forces

Forces between molecules that affect physical properties such as boiling and melting points.

Phase Diagram

A graphical representation of the physical states of a substance under different conditions of temperature and pressure.

Gas Laws

Equations that describe the behavior of gases, including the Ideal Gas Law (PV=nRT).

Thermochemistry

The study of heat changes that accompany chemical reactions.

Hess's Law

The principle stating that the total enthalpy change for a reaction is the sum of the enthalpy changes for the individual steps.

Acid-Base Reaction

A chemical reaction that typically occurs when an acid and a base react to form water and a salt.

Limiting Reactant

The reactant that is entirely consumed in a chemical reaction and thus limits the amount of product formed.

What is molecular geometry?

The three-dimensional arrangement of atoms within a molecule.

Linear geometry

A molecular shape where atoms are arranged in a straight line, with a bond angle of 180 degrees.

Trigonal planar geometry

A molecular shape with three atoms positioned around a central atom, forming a planar triangle with bond angles of 120 degrees.

Tetrahedral geometry

A molecular shape with four atoms surrounding a central atom, forming a tetrahedron with bond angles of approximately 109.5 degrees.

Trigonal bipyramidal geometry

A molecular shape where five atoms are arranged around a central atom, with three atoms in a plane and two above and below, resulting in bond angles of 90 and 120 degrees.

Octahedral geometry

A molecular shape with six atoms situated around a central atom, resembling two pyramids base-to-base, with bond angles of 90 degrees.

Bent geometry

A molecular shape that occurs when there are lone pairs on the central atom, resulting in a non-linear shape with bond angles less than 120 degrees for trigonal bent and less than 109.5 degrees for tetrahedral bent.

Trigonal pyramidal geometry

A molecular shape created when a central atom is bonded to three atoms and has one lone pair, resulting in a pyramid-like structure with bond angles of approximately 107 degrees.

Square planar geometry

A molecular shape that occurs when there are four atoms arranged in a square plane around a central atom, with bond angles of 90 degrees.

Seesaw geometry

A molecular shape that arises in trigonal bipyramidal molecules with one lone pair, resulting in a structure where one atom is above and below the central atom, creating varying bond angles.

Periodic Trends

Patterns observed in the properties of elements across periods and down groups in the periodic table.

Atomic Radius

The distance from the nucleus of an atom to the edge of its electron cloud; it generally decreases across a period and increases down a group.

Ionization Energy

The energy required to remove an electron from an atom; it typically increases across a period and decreases down a group.

Electronegativity

A measure of the tendency of an atom to attract a bonding pair of electrons; it generally increases across a period and decreases down a group.

Electron Affinity

The energy change that occurs when an electron is added to a neutral atom; it is usually more negative across a period.

Metallic Character

The tendency of an element to behave like a metal; it decreases across a period and increases down a group.

Reactivity of Metals

The tendency of metals to lose electrons and form positive ions; it generally increases as you move down a group.

Reactivity of Nonmetals

The tendency of nonmetals to gain electrons and form negative ions; it usually decreases as you move down a group.

Valence Electrons

Electrons in the outermost shell of an atom that participate in chemical reactions; these determine an element's chemical properties.

Atomic Mass Trend

Atomic mass generally increases from left to right across a period and down a group due to the addition of protons and neutrons.

Hybridization

The concept of mixing atomic orbitals to form new hybrid orbitals for bonding in molecules.

sp Hybridization

A type of hybridization that occurs when one s and one p orbital mix to form two equivalent sp hybrid orbitals, typically seen in linear molecular geometries.

sp2 Hybridization

A type of hybridization where one s and two p orbitals combine to form three equivalent sp2 hybrid orbitals, typically found in trigonal planar geometries.

sp3 Hybridization

A type of hybridization involving one s and three p orbitals that creates four equivalent sp3 hybrid orbitals, associated with tetrahedral geometries.

Strong Acid

An acid that completely dissociates in water, releasing all its protons (H+), e.g., hydrochloric acid (HCl).

Weak Acid

An acid that partially dissociates in water, establishing an equilibrium between the undissociated acid and the ions produced, e.g., acetic acid (CH3COOH).

Strong Base

A base that completely dissociates in water to produce hydroxide ions (OH-), e.g., sodium hydroxide (NaOH).

Weak Base

A base that partially dissociates in water to produce hydroxide ions, establishing an equilibrium, e.g., ammonia (NH3).

Bronsted-Lowry Acid

A substance that donates a proton (H+) in a chemical reaction.

Bronsted-Lowry Base

A substance that accepts a proton (H+) in a chemical reaction.

pKa

A measure of the strength of an acid; the lower the pKa value, the stronger the acid.

Conjugate Acid-Base Pair

A pair of compounds that differ by the presence or absence of a proton; the acid donates a proton to form its conjugate base and vice versa.

Solubility Rules

Guidelines used to predict the solubility of ionic compounds in water.

Soluble Salts

Nitrates (NO3-), acetates (CH3COO-), and most chlorides, bromides, and iodides are generally soluble in water.

Insoluble Salts

Most carbonates, phosphates, sulfides, and hydroxides are generally insoluble, except those of alkali metals and ammonium.

Hydrogen Bonding

A strong intermolecular force that occurs when hydrogen is bonded to highly electronegative atoms like N, O, or F.

Dipole-Dipole Interactions

Intermolecular forces between polar molecules due to the attraction between positive and negative ends.

London Dispersion Forces

Weak intermolecular forces caused by temporary fluctuations in electron distribution, present in all molecules.

Intermolecular Force Strength Ranking

The order of strength from weakest to strongest is London dispersion < dipole-dipole < hydrogen bonds.

Factors Influencing Solubility

Temperature, pressure, and nature of solute and solvent determine the solubility of a substance.

Like Dissolves Like

A principle stating that polar solvents dissolve polar solutes, and nonpolar solvents dissolve nonpolar solutes.

Salt Effect on Solubility

The presence of salts can affect the solubility of other compounds due to common ion effects

Enthalpy (H)

The total heat content of a system, often measured in joules (J) or kilojoules (kJ).

Exothermic Reaction

A chemical reaction that releases heat to the surroundings, resulting in a temperature increase.

Endothermic Reaction

A chemical reaction that absorbs heat from the surroundings, resulting in a temperature decrease.

Standard Enthalpy Change (ΔH°)

The change in enthalpy for a process measured under standard conditions (1 atm, 25°C).

Heat of Formation

The change in enthalpy when one mole of a compound is formed from its elements in their standard states.

Common Compounds: Water

Formula: H₂O; the most abundant compound on Earth, essential for life.

Common Compounds: Sodium Chloride

Formula: NaCl; common table salt, used in food and as a preservative.

Common Compounds: Ammonia

Formula: NH₃; a compound used in fertilizers and cleaning products.

Common Compounds: Ethanol

Formula: C₂H₅OH; a common alcohol used in beverages and as a solvent.

Common Compounds: Glucose

Formula: C₆H₁₂O₆; a simple sugar and important energy source for living organisms.

Calorimetry

The measurement of heat transfer during chemical reactions or physical changes.

Hess's Law Application

Method for calculating enthalpy changes by summing the enthalpy changes for individual steps of a reaction.

Thermodynamic Equilibrium

A state where a system's macroscopic properties are stable over time, with no net change.