Matter, Atoms, and Molecules

Definitions and Distinctions

This section introduces the foundational concepts of chemistry, focusing on the basic building blocks of matter and their organization.

• Atom: The smallest unit of an element that retains the properties of that element.

• Molecule: A group of two or more atoms bonded together, representing the smallest fundamental unit of a chemical compound.

• Chemistry: The scientific study of matter, its properties, and the changes it undergoes.

• Difference between atoms and molecules: Atoms are single units; molecules are combinations of atoms bonded together.

• Drawing molecules: Simple molecules such as carbon monoxide (CO), carbon dioxide (CO2), water (H2O), and hydrogen peroxide (H2O2) can be represented using spheres for atoms and sticks for bonds.

Example: Water (H2O) consists of two hydrogen atoms bonded to one oxygen atom.

The Scientific Approach to Knowledge

Scientific Method and Key Terms

The scientific method is a systematic approach to understanding natural phenomena.

• Hypothesis: A tentative explanation for an observation, which can be tested by experiments.

• Theory: A well-substantiated explanation of some aspect of the natural world, based on a body of evidence.

• Scientific Law: A statement that describes a consistently observed phenomenon, often expressed mathematically.

• Experiment: A controlled procedure carried out to test a hypothesis.

• Falsifiable: Capable of being disproved by experimental results.

• Law of Conservation of Mass: Mass is neither created nor destroyed in a chemical reaction.

• Principle vs. Law: A principle is a general scientific concept; a law is a concise statement of a relationship.

• Dalton’s Atomic Theory: States that all matter is composed of atoms, atoms of a given element are identical, and atoms combine in simple ratios to form compounds.

• Scientific Approach Steps: Observation → Hypothesis → Experiment → Theory/Law.

Example: The law of conservation of mass can be demonstrated by weighing reactants and products in a closed system.

The Classification of Matter

States and Types of Matter

Matter can be classified by its physical state and composition.

• Matter: Anything that has mass and occupies space.

• States of Matter: Solid, liquid, gas.

• Properties: Solids have definite shape and volume; liquids have definite volume but no definite shape; gases have neither definite shape nor volume and are compressible.

• Crystalline vs. Amorphous Solids: Crystalline solids have an ordered structure; amorphous solids lack long-range order.

• Mixture: A combination of two or more substances not chemically bonded.

• Pure Substance: Matter with a fixed composition (element or compound).

• Element: A substance that cannot be broken down into simpler substances.

• Compound: A substance composed of two or more elements chemically combined.

• Heterogeneous Mixture: Composition varies throughout.

• Homogeneous Mixture (Solution): Uniform composition throughout.

Example: Salt water is a homogeneous mixture; sand and iron filings form a heterogeneous mixture.

Physical and Chemical Changes and Properties

Types of Changes and Properties

Understanding the difference between physical and chemical changes is essential in chemistry.

• Physical Change: Alters the state or appearance but not composition (e.g., melting ice).

• Chemical Change: Alters the composition, producing new substances (e.g., rusting iron).

• Physical Property: Can be observed without changing the substance (e.g., color, melting point).

• Chemical Property: Describes the ability to undergo a chemical change (e.g., flammability).

Example: Boiling water is a physical change; burning wood is a chemical change.

Energy in Physical and Chemical Change

Forms and Conservation of Energy

Energy is a fundamental concept in understanding chemical and physical changes.

• Energy: The capacity to do work.

• Work: The result of a force acting over a distance.

• Kinetic Energy: Energy of motion.

• Potential Energy: Stored energy due to position or composition.

• Thermal Energy: Energy associated with temperature.

• Law of Conservation of Energy: Energy cannot be created or destroyed, only transformed.

Example: When gasoline burns, chemical potential energy is converted to thermal energy and work.

Units of Measurement

Measurement Systems and Conversions

Accurate measurement is crucial in chemistry, requiring understanding of units and conversions.

• System of Measurement: Standardized set of units (English, metric, SI).

• SI Base Units: Length (meter, m), mass (kilogram, kg), time (second, s), temperature (kelvin, K).

• Temperature Scales: Fahrenheit (°F), Celsius (°C), Kelvin (K).

• Freezing/Boiling Points of Water:

  • Fahrenheit: 32°F / 212°F

  • Celsius: 0°C / 100°C

  • Kelvin: 273.15 K / 373.15 K

• Temperature Conversions:

• Prefixes: Used to denote powers of ten (e.g., kilo-, centi-, milli-).

• Derived Units: Volume (liter, L), density ().

Example: 1 mL = 1 cm3; 1 kg = 1000 g.

Reliability of Measurement

Significant Figures, Precision, and Accuracy

Proper reporting of measurements ensures scientific reliability.

• Significant Figures: Digits in a measurement that are known with certainty plus one estimated digit.

• Exact Numbers: Numbers with no uncertainty (e.g., counting numbers, defined quantities).

• Precision: How close repeated measurements are to each other.

• Accuracy: How close a measurement is to the true value.

• Rules for Significant Figures: Nonzero digits are always significant; zeros between nonzero digits are significant; leading zeros are not significant; trailing zeros are significant if there is a decimal point.

• Significant Figures in Calculations: For multiplication/division, use the least number of significant figures; for addition/subtraction, use the least number of decimal places.

Example: 0.00450 has three significant figures.

Solving Chemical Problems

Dimensional Analysis and Problem-Solving Strategies

Dimensional analysis is a systematic approach to solving problems involving unit conversions.

• Dimensional Analysis: Using conversion factors to move from one unit to another.

• Conversion Factor: A ratio that expresses how many of one unit are equal to another unit.

• Problem-Solving Plan: Sort (identify knowns/unknowns), strategize (plan steps), solve (do calculations), check (verify answer).

• Unit Conversions: Use conversion factors to change units (e.g., inches to centimeters).

• Powers and Equations: Apply exponents and algebraic manipulation as needed.

Example: To convert 2.5 inches to centimeters:

Modern Atomic Theory and Laws

Atomic Theory and Fundamental Laws

Modern atomic theory is based on several key scientific laws and experimental evidence.

• Law of Conservation of Mass: Mass is conserved in chemical reactions.

• Law of Definite Proportions: A given compound always contains the same proportion of elements by mass.

• Law of Multiple Proportions: When two elements form more than one compound, the ratios of the masses of the second element that combine with a fixed mass of the first element are ratios of small whole numbers.

• Dalton’s Atomic Theory: All matter is made of atoms; atoms of the same element are identical; atoms combine in simple ratios to form compounds; atoms are rearranged in chemical reactions.

Example: Water (H2O) always has a mass ratio of 8:1 oxygen to hydrogen.

Discovery of Subatomic Particles

Electrons, Protons, and Neutrons

The structure of the atom was elucidated through key experiments.

• Cathode Ray: A stream of electrons observed in vacuum tubes.

• J. J. Thomson’s Experiment: Showed that cathode rays are composed of negatively charged particles (electrons).

• Millikan’s Oil Drop Experiment: Measured the charge of the electron.

• Radioactivity: Spontaneous emission of particles or energy from unstable nuclei (alpha, beta, gamma particles).

• Plum Pudding Model: Early model of the atom with electrons embedded in a positively charged sphere.

• Rutherford’s Gold Foil Experiment: Demonstrated the existence of a small, dense, positively charged nucleus.

• Nuclear Theory: Atom consists of a nucleus (protons and neutrons) surrounded by electrons.

Example: Alpha particles were deflected by the nucleus in Rutherford’s experiment, disproving the plum pudding model.

Subatomic Particles in Atoms

Atomic Structure and Isotopes

Atoms are composed of protons, neutrons, and electrons, each with specific properties.

• Atomic Mass Unit (amu): Standard unit for atomic mass; 1 amu ≈ mass of one proton or neutron.

• Atomic Number (Z): Number of protons in the nucleus; defines the element.

• Chemical Symbol: One- or two-letter abbreviation for an element.

• Isotopes: Atoms of the same element with different numbers of neutrons.

• Mass Number (A): Total number of protons and neutrons ().

• Natural Abundance: Relative percentage of each isotope in a natural sample.

• Ions: Atoms or molecules with a net charge due to loss or gain of electrons.

• Anion: Negatively charged ion (gained electrons).

• Cation: Positively charged ion (lost electrons).

Example: Carbon-12 has 6 protons and 6 neutrons; carbon-14 has 6 protons and 8 neutrons.

Atomic Mass and Isotopic Calculations

Average Atomic Mass

The atomic mass of an element is a weighted average of the masses of its isotopes.

• Calculation:

Example: If an element has two isotopes, 75% at 10 amu and 25% at 11 amu: amu.

The Periodic Table and Periodic Law

Organization and Classification

The periodic table arranges elements by increasing atomic number and recurring chemical properties.

• Periodic Law: Properties of elements recur periodically when arranged by atomic number.

• Groups/Families: Columns with similar properties (e.g., noble gases, alkali metals, alkaline earth metals, halogens).

• Metals, Nonmetals, Metalloids: Metals are conductive and malleable; nonmetals are insulating and brittle; metalloids have intermediate properties.

• Main-Group vs. Transition Elements: Main-group elements are in groups 1, 2, and 13-18; transition elements are in groups 3-12.

• Predictable Ion Charges: Elements in the same group often form ions with the same charge.

Example: Sodium (Na) is an alkali metal and forms a +1 ion.

Molar Mass and Counting Atoms

The Mole and Avogadro’s Number

The mole is a counting unit in chemistry, relating mass to number of particles.

• Mole (mol): Amount of substance containing entities (Avogadro’s number).

• Molar Mass: Mass of one mole of a substance (g/mol).

• Conversions:

  • Number of moles =

  • Number of atoms = moles Avogadro’s number

• Dimensional Analysis: Used to convert between mass, moles, and number of particles.

Example: 18 g of water () is 1 mole, containing molecules.

Chemical Bonds

Ionic and Covalent Bonds

Chemical bonds hold atoms together in compounds.

• Chemical Bond: Force that holds atoms together in a compound.

• Ionic Bond: Transfer of electrons from one atom to another, forming ions (typically between metals and nonmetals).

• Covalent Bond: Sharing of electrons between atoms (typically between nonmetals).

• Mixtures vs. Compounds: Mixtures are physical combinations; compounds are chemically bonded.

• Ionic Compound Formation: Metal loses electrons to become a cation; nonmetal gains electrons to become an anion; oppositely charged ions attract.

• Covalent Bond Formation: Atoms share electrons to achieve stable electron configurations.

Example: Sodium chloride (NaCl) is an ionic compound; water (H2O) is a covalent compound.

Representing Compounds

Chemical Formulas and Models

Compounds can be represented in various ways to convey different information.

• Empirical Formula: Simplest whole-number ratio of elements in a compound.

• Molecular Formula: Actual number of each type of atom in a molecule.

• Structural Formula: Shows how atoms are connected.

• Models: Ball-and-stick and space-filling models represent molecular geometry.

Example: Hydrogen peroxide: Empirical formula HO, molecular formula H2O2, structural formula H–O–O–H.

Classification of Elements and Compounds

Atomic and Molecular Elements, Ionic and Molecular Compounds

Pure substances can be classified based on their composition and structure.

• Atomic Element: Exists as single atoms (e.g., Ne, Ar).

• Molecular Element: Exists as molecules (e.g., O2, N2).

• Ionic Compound: Composed of cations and anions in a lattice (formula units).

• Molecular Compound: Composed of molecules formed by covalent bonds.

• Polyatomic Ion: A charged group of covalently bonded atoms (e.g., SO42–).

Example: NaCl is an ionic compound; CO2 is a molecular compound.

Ionic Compounds: Formulas and Names

Naming and Writing Formulas

Ionic compounds are named and written according to specific rules.

• Formula Writing: Combine cations and anions in ratios that yield a neutral compound.

• Naming: Name the cation first, then the anion (e.g., sodium chloride).

• Transition Metals: May have multiple charges; indicate charge with Roman numerals (e.g., iron(III) chloride).

• Polyatomic Ions: Ions composed of multiple atoms (e.g., NO3–, SO42–).

• Hydrates: Ionic compounds with water molecules; named with prefixes (e.g., copper(II) sulfate pentahydrate).

Example: CaCl2 is calcium chloride; FeCl3 is iron(III) chloride.

Molecular Compounds: Formulas and Names

Naming Molecular Compounds and Acids

Molecular compounds and acids have distinct naming conventions.

• Naming Molecular Compounds: Use prefixes to indicate number of atoms (e.g., dinitrogen tetroxide).

• Binary Acids: Contain hydrogen and one other element; named as "hydro-" + base name + "-ic acid" (e.g., hydrochloric acid).

• Oxyacids: Contain hydrogen, oxygen, and another element; named based on the polyatomic ion (e.g., sulfuric acid from sulfate).

Example: CO2 is carbon dioxide; HNO3 is nitric acid.

Table: Common Polyatomic Ions (Inferred)

Purpose: Classification and identification of common polyatomic ions.

Additional info: Table content inferred as common polyatomic ions are referenced in the objectives.