Comprehensive Study Guide for Science 10 Unit B: Chemistry

Fundamentals of Chemistry and Laboratory Safety

• Core Definitions:

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

  • Matter: Any substance that possesses mass and occupies physical space.

  • Atom: The fundamental building block of all matter.

• Mandatory Laboratory Safety Guidelines:

  • Personal Protective Equipment (PPE): Chemical splash goggles and personal protective equipment must be worn whenever required in the laboratory environment.

  • Attire: Full-length pants and closed-toed shoes are strictly mandatory.

  • Personal Preparedness: Long hair must be securely tied back; loose clothing should be contained.

  • Prohibited Items and Behaviors: No food, beverages, personal listening devices (earbuds), or horseplay are permitted in the laboratory at any time.

  • Protocol and Cleanliness: Laboratory instructions must be followed with precision, and workstations must be kept clean, organized, and free of clutter.

  • Olfactory Technique: Direct smelling of chemical substances is prohibited; wafting must always be used to draw vapors gently toward the nose.

  • Chemical Stock Preservation: Used or excess chemicals must never be returned to their original stock containers to avoid contamination.

  • Incident Reporting: Any chemical spill, cracked apparatus, or broken equipment must be reported immediately to the teacher.

  • Disposal and Hygiene: Chemical and biological waste must be discarded in designated disposal receptacles, followed by thorough handwashing with soap and water.

• Emergency First-Aid Protocols:

  • Chemical Contact on Skin: Flush the affected area immediately under running water for a minimum duration of $10\text{ to }15\,\text{min}$.

  • Chemical Contact in Eyes: Use the emergency eyewash station immediately; force the eyelids open and flush for $10\text{ to }15\,\text{min}$.

  • Chemical Spills: Immediately alert the teacher to ensure appropriate containment and cleanup materials are deployed.

  • Person on Fire: Smother the flames by wrapping the individual in a dedicated fire blanket (stop, drop, and roll).

  • Equipment or Laboratory Fire: Alert the teacher immediately; operate a certified fire extinguisher only if safe to do so.

  • Embedded Glass Shards: If broken glass penetrates the skin or tissue, leave the shard in place to prevent severe hemorrhaging and seek professional medical assistance immediately.

• Workplace Hazardous Materials Information System (WHMIS) Pictograms:

  • Skull and Crossbones: Indicates acute toxicity; substance is fatal or highly toxic upon acute exposure.

  • Person with Star on Chest: Denotes a serious health hazard; substance may cause cancer (carcinogenicity), respiratory sensitization, or reproductive toxicity.

  • Exclamation Mark: Identifies an irritant or less severe health hazard (skin/eye irritation, harmful toxicity).

  • Biohazard: Identifies biohazardous and infectious materials capable of transmitting disease.

  • Dead Tree and Fish: Denotes aquatic and environmental toxicity.

  • Corrosion: Indicates corrosive materials capable of chemically degrading and damaging metals, skin tissue, and eyes.

  • Gas Cylinder: Identifies compressed gases and gases held under pressure.

  • Flame: Denotes flammable substances, pyrophoric materials, and self-heating agents.

  • Flame Over Circle: Identifies oxidizing substances that supply oxygen to intensify fire and combustion.

  • Exploding Bomb: Denotes explosive, self-reactive substances, and organic peroxides.

• Safety Data Sheet (SDS):

  • A standardized technical document providing comprehensive information regarding chemical hazards, safe handling procedures, physical properties, toxicological exposure effects, and specific first-aid countermeasures.

Classification and Properties of Matter

• Taxonomy of Matter:

  • All matter is classified into two overarching categories: pure substances and mixtures.

  • Separation Criterion: Mixtures can be separated into their constituent parts through physical separation methods (such as filtration, distillation, or chromatography), whereas pure substances cannot be separated by physical means.

• Pure Substances:

  • Element: A pure chemical substance consisting entirely of one type of atom that cannot be chemically broken down into simpler substances (e.g., gold).

  • Compound: A pure chemical substance consisting of two or more distinct chemical elements chemically combined in fixed, definite proportions (e.g., water, $\text{H}_2\text{O}$).

• Mixtures:

  • Homogeneous Mixture (Solution): A mixture with uniform composition and properties throughout; distinct phase boundaries cannot be seen (e.g., atmospheric air, vinegar).

  • Heterogeneous Mixture: A mixture characterized by non-uniform composition where individual constituents or phases can be distinguished visually or under magnification (e.g., soil, orange juice containing pulp).

Historical Development of Atomic Theory

• Chronological Milestones:

  • Democritus (circa $400\,\text{BC}$): Conceptualized the existence of fundamental, indivisible particles that constitute matter, coining the term atomos (meaning "indivisible").

  • John Dalton ($1804$): Formulated the modern atomic theory and introduced the "solid sphere model," proposing that matter consists of discrete, solid atomic spheres.

  • J.J. Thomson ($1890\text{s}$): Discovered the negatively charged electron using cathode ray tube experiments; formulated the "plum pudding model" where electrons are embedded in a diffuse sphere of positive charge.

  • Ernest Rutherford ($1911$): Conducted the gold foil experiment; discovered the atomic nucleus, demonstrating that atoms consist primarily of empty space surrounding a tiny, dense, positively charged core.

  • Niels Bohr ($1913$): Proposed the planetary model of the atom, stating that electrons occupy distinct, quantized energy levels (shells) and transition between them by absorbing or emitting photons of discrete energy.

  • James Chadwick ($1932$): Discovered the neutron, the electrically neutral subatomic particle residing within the atomic nucleus.

  • Quantum Mechanical Model: Developed modern atomic physics, proposing that electrons do not follow fixed orbits but rather occupy three-dimensional probability regions (orbitals or electron clouds) around the nucleus.

• Dalton's Five Core Postulates:

  • 1. All matter is composed of extremely small, indivisible particles called atoms.

  • 2. All atoms of a given element were posited to be identical in mass, size, and properties.

  • 3. Atoms of different elements possess different masses and physical/chemical properties.

  • 4. Chemical compounds are formed when atoms of different elements combine in fixed, simple whole-number ratios.

  • 5. In chemical reactions, atoms are rearranged, combined, or separated; they are never created or destroyed.

• Thomson's Cathode Ray Experiment:

  • Cathode rays were observed to bend toward positively charged plates and deflect away from negatively charged plates.

  • Established that cathode rays consist of negatively charged fundamental particles (electrons).

  • Visual Model: A uniform sphere of positive charge dotted with embedded negative electrons, analogous to raisins in plum pudding.

• Rutherford's Gold Foil Experiment:

  • Alpha particles ($\alpha$, positive helium nuclei) were directed at an ultra-thin sheet of gold foil.

  • Observations: The overwhelming majority of alpha particles passed directly through the foil without deflection; a very small fraction was deflected at steep angles, and some bounced directly backward.

  • Conclusions: The atom is mostly empty space; all positive charge and nearly all the atom's mass are concentrated in an extraordinarily small, dense central region named the nucleus.

• Bohr's Energy Level Model:

  • Electrons are confined to specific, stable orbits (energy levels) at defined distances from the nucleus.

  • Electrons absorb discrete packets of energy to jump to higher levels (excited states) and release energy as electromagnetic radiation when dropping to lower levels (ground states).

  • Experimental Evidence: Successfully explained and matched the discrete lines in hydrogen's emission spectrum.

  • Limitation: The Bohr model failed to accurately calculate and explain the emission spectra and behaviors of multi-electron atoms.

• Quantum Mechanical Probability Model:

  • Electrons exhibit wave-particle duality and occupy diffuse probability clouds around the nucleus.

  • The spatial density of the cloud corresponds mathematically to the probability of locating an electron within that specific volume; regions of higher density correspond to a higher probability of electron presence.

Subatomic Particles, Atomic Structure, and Calculations

• Characteristics of Fundamental Subatomic Particles:

  • Proton: Relative electric charge of $+1$; spatial location is within the nucleus; relative mass is assigned a value of $1$.

  • Neutron: Relative electric charge of $0$ (neutral); spatial location is within the nucleus; relative mass is assigned a value of $1$.

  • Electron: Relative electric charge of $-1$; spatial location is within the electron cloud outside the nucleus; relative mass is $\frac{1}{1840}$.

• Structural and Atomic Definitions:

  • Nucleons: Collective term designating the nuclear particles, namely protons and neutrons.

  • Neutral Atom: An uncharged atom in which the total number of protons strictly equals the total number of electrons.

  • Atomic Number ($Z$): The total quantity of protons inside an atom's nucleus; this value defines the elemental identity. Altering the proton count fundamentally transforms the atom into a different element.

  • Mass Number ($A$): The total combined quantity of protons and neutrons in a nucleus: $A = \text{protons} + \text{neutrons}$.

  • Isotopes: Atoms of the exact same chemical element (identical atomic number) containing differing numbers of neutrons, resulting in different mass numbers.

• Core Atomic Calculation Equations:

  • $\text{Neutrons} = \text{Mass number} - \text{Atomic number}$

  • $\text{Protons} = \text{Atomic number}$

  • $\text{Electrons} = \text{Protons (in a neutral atom)}$

• Calculation Example (Phosphorus-$31$):

  • Atomic number ($Z$) = $15$

  • Protons = $15$

  • Electrons = $15$

  • Neutrons = $31 - 15 = 16$

• Standard Isotopic Notation:

  • Structure: The mass number is placed as an upper-left superscript, and the atomic number is placed as a lower-left subscript preceding the elemental symbol: $\text{}^{A}_{Z}\text{X}$.

  • Example: Phosphorus-$31$ is designated as $\text{}^{31}_{15}\text{P}$.

Modified Bohr Diagrams and Electron Shell Configurations

• Step-by-Step Construction Procedure:

  • 1. Determine the elemental atomic number and mass number from the periodic table.

  • 2. Set the proton count equal to the atomic number.

  • 3. Calculate the neutron count by subtracting the atomic number from the mass number ($\text{neutrons} = A - Z$).

  • 4. For a neutral atom, establish an electron count equal to the proton count.

  • 5. Sketch the central nucleus indicating internal proton and neutron totals, then arrange the electrons in concentric circular energy levels surrounding the nucleus.

• Shell Capacity Rules (First $20$ Elements):

  • The principal energy levels fill progressively from the inside out adhering to the sequence: $2, 8, 8, 2$.

  • Level 1: Maximum $2$ electrons.

  • Level 2: Maximum $8$ electrons.

  • Level 3: Maximum $8$ electrons.

  • Level 4: Maximum $2$ electrons (for elements up to calcium, $Z = 20$).

  • Example — Phosphorus-$31$ ($15$ electrons): Electron configuration across energy levels is $2, 8, 5$.

• Valence Electrons:

  • Defined as the electrons located in the outermost occupied energy level of an atom.

  • Play the primary role in dictating the atom's chemical properties, reactivity, and bonding behavior.

The Periodic Table: Architecture and Elemental Families

• Historical and Modern Organization:

  • Dmitri Mendeleev: Established the foundational framework of the periodic table by organizing elements according to relative atomic mass and successfully predicting the existence and properties of undiscovered elements.

  • Modern Organization: Elements are arranged sequentially in order of increasing atomic number ($Z$).

  • Periods: The horizontal rows of the periodic table; the period number corresponds directly to the total number of occupied electron energy levels.

  • Groups or Families: The vertical columns of the periodic table; elements within the same group share identical valence electron configurations and exhibit similar chemical properties.

• Valence Electron Counts by Group Number:

  • Group 1: $1$ valence electron

  • Group 2: $2$ valence electrons

  • Group 13: $3$ valence electrons

  • Group 14: $4$ valence electrons

  • Group 15: $5$ valence electrons

  • Group 16: $6$ valence electrons

  • Group 17: $7$ valence electrons

  • Group 18: $8$ valence electrons (Helium possesses only $2$ valence electrons)

• Elemental Classifications:

  • Metals:

  • Located to the left of the periodic staircase.

  • Physical Characteristics: Typically lustrous (shiny), malleable (can be hammered into sheets), and ductile (can be drawn into wires).

  • Thermal and Electrical Properties: Highly efficient conductors of thermal energy and electricity.

  • Non-metals:

  • Located to the right of the periodic staircase.

  • Physical Characteristics: Typically dull in appearance, non-ductile, and brittle in the solid state.

  • Thermal and Electrical Properties: Poor conductors of heat and electricity (insulators).

  • Metalloids:

  • Positioned immediately along both borders of the staircase line.

  • Exhibit intermediate properties sharing characteristics of both metals and non-metals.

  • Generally brittle solids; function as semiconductors of electricity.

• Primary Periodic Families:

  • Group 1 — Alkali Metals: Possess $1$ valence electron; extremely reactive chemically, demonstrating vigorous reactivity with water to produce hydrogen gas and basic solutions.

  • Group 2 — Alkaline-Earth Metals: Possess $2$ valence electrons; reactive metallic elements.

  • Groups 3 to 12 — Transition Metals: Metallic elements displaying variable valence states, complex electron configurations, and varied physical and chemical properties.

  • Group 17 — Halogens: Possess $7$ valence electrons; highly reactive non-metallic elements that readily accept electrons.

  • Group 18 — Noble Gases: Possess completely filled outer energy levels ($8$ valence electrons, except Helium which has $2$); chemically inert and non-reactive under standard conditions.

Chemistry in Traditional Practices, Society, and Industry

• Traditional and Historical Chemical Applications:

  • Food Preservation: Chemical processes that halt or dramatically retard microbial spoilage. Techniques include drying (moisture reduction), smoking (depositing phenolic antimicrobial compounds), salting (osmotic dehydration of microbes), and fermentation (acid/alcohol production).

  • Brain Tanning: An Indigenous technology utilizing emulsified oils from animal brain matter combined with smoke aldehydes to soften animal hides, making them pliable, water-resistant, and flexible.

  • Natural Dyes: Extraction of chemical pigments from plants, berries, and lichens to permanently color fabrics and tools.

  • Medicinal Plants: Utilization of natural chemical compounds for therapeutic treatment. Traditional use of willow bark (containing salicin, the chemical precursor to modern acetylsalicylic acid) treats pain and inflammation; additional herbal treatments utilize yarrow and balsam fir.

  • Annealing: Controlled thermal processing involving heating metal to elevated temperatures followed by slow cooling, altering crystalline microstructure to relieve internal stresses and increase ductility.

  • Alkaline Ash: Sourcing basic potassium carbonate and sodium carbonate from wood ashes to formulate cleansing detergents (traditional lye soap making) and to assist in the dehairing and tanning of animal hides.

  • Chemical Etching: Applying natural acids to selectively dissolve basic minerals, such as calcium carbonate ($\text{CaCO}_3$) found in marine shells, creating decorative patterns.

• Chemically Oriented Professional Careers:

  • Hair Stylist: Employs redox reactions to chemically bleach melanin pigments and utilize oxidation dyes to restructure and color hair.

  • Gas Technician: Manages and services pressurized combustible gas systems, ensuring safe fuel-air stoichiometry.

  • Pharmacist: Formulates, compounds, and dispenses therapeutic chemical compounds and manages pharmaceutical biological interactions.

  • Lifeguard: Regularly measures aqueous pool chemistry, balancing water $pH$ and active chlorine sanitizing levels.

  • Forensic Chemist: Analyzes trace materials, unknown toxins, and biochemical evidence collected from crime scenes.

  • Research and Development Chemist: Synthesizes, optimizes, and analyzes commercial chemical products, such as improved household cleaners and polymers.

Chemical Bonding and Ionic Compounds

• General Principles of Chemical Bonding:

  • Chemical Compound: A distinct substance formed by the chemical combination of two or more different elements.

  • Octet Rule: Atoms transfer or share valence electrons to attain a stable, low-energy electronic configuration comprising $8$ valence electrons (isoelectronic with the nearest noble gas).

• Ionic Bonding:

  • Electrostatic attraction formed between a metal cation and a non-metal anion.

  • Mechanism: Electrons are completely transferred from the metallic element to the non-metallic element.

  • Oxidation and Reduction Behavior: Metals lose valence electrons to form positively charged cations; non-metals gain valence electrons to form negatively charged anions.

  • Structural Outcome: Forms an extended, repeating three-dimensional crystal lattice, never isolated individual molecules.

• Covalent (Molecular) Bonding:

  • Chemical bond formed exclusively between two or more non-metal atoms.

  • Mechanism: Non-metal atoms share pairs of valence electrons to fulfill octet requirements.

  • Structural Outcome: Forms discrete, individual molecules (e.g., $\text{H}_2\text{O}$, $\text{NH}_3$).

• Ion Dynamics and Radii:

  • Cation ($+$

  • Formed when a neutral metal atom loses one or more valence electrons.

  • Retains its elemental nomenclature (e.g., a sodium atom forms a sodium ion).

  • Atomic Radius Trend: Cations are significantly smaller than their parent neutral atoms due to the loss of an outer electron shell and increased effective nuclear attraction on remaining electrons.

  • Anion ($-$

  • Formed when a neutral non-metal atom gains electrons into its valence shell.

  • Nomenclature Rule: The root element name is modified to take an "-ide" suffix (e.g., chlorine gains an electron to become chloride).

  • Atomic Radius Trend: Anions are significantly larger than their parent neutral atoms due to increased electron-electron electrostatic repulsion within the valence shell.

• IUPAC Rules for Binary Ionic Nomenclature:

  • Binary ionic compounds are composed of exactly two elements: a metal and a non-metal.

  • 1. Write the full elemental name of the metal first.

  • 2. Write the non-metal elemental root second, replacing its terminal ending with the suffix "-ide".

  • 3. Chemical compound names are never capitalized in prose unless beginning a sentence.

  • Examples:

  • $\text{NaCl}$: Sodium chloride

  • $\text{Mg}_3\text{N}_2$: Magnesium nitride

  • $\text{K}_2\text{O}$: Potassium oxide

  • $\text{AlCl}_3$: Aluminium chloride

• Deriving Binary Ionic Formulas:

  • The sum of all positive and negative charges in an ionic formula unit must balance to zero.

  • 1. Write the elemental symbols with their respective ionic charge superscripts.

  • 2. Balance total electrical charges by applying whole-number stoichiometric subscripts.

  • 3. Reduce all subscripts to the simplest empirical whole-number ratio.

  • 4. Omit writing the subscript when the value is $1$.

  • Examples:

  • $\text{Mg}^{2+} + \text{Cl}^- \rightarrow \text{MgCl}_2$

  • $\text{Zn}^{2+} + \text{N}^{3-} \rightarrow \text{Zn}_3\text{N}_2$

  • $\text{Ca}^{2+} + \text{O}^{2-} \rightarrow \text{CaO}$ (subscripts reduce from $2:2$ to $1:1$, not $\text{Ca}_2\text{O}_2$)

Polyatomic, Multivalent, and Hydrated Ionic Compounds

• Polyatomic Ions:

  • Defined as a covalently bonded cluster of atoms carrying a net positive or negative electrical charge that functions as a single chemical unit in reactions.

  • Nomenclature Rule: Retain the established, unaltered name of the polyatomic ion within the compound name.

  • Structural Writing Rule: Enclose the polyatomic formula within parentheses whenever a subscript of $2$ or greater is appended to denote multiple polyatomic units.

• Essential Polyatomic Ions:

  • Hydroxide: $\text{OH}^-$

  • Nitrate: $\text{NO}_3^-$

  • Nitrite: $\text{NO}_2^-$

  • Sulfate: $\text{SO}_4^{2-}$

  • Sulfite: $\text{SO}_3^{2-}$

  • Carbonate: $\text{CO}_3^{2-}$

  • Phosphate: $\text{PO}_4^{3-}$

  • Dichromate: $\text{Cr}_2\text{O}_7^{2-}$

  • Cyanide: $\text{CN}^-$

  • Ammonium: $\text{NH}_4^+$

  • Hydrogen carbonate: $\text{HCO}_3^-$

• Polyatomic Compound Examples:

  • $\text{NaClO}_2$: Sodium chlorite

  • $\text{Mg(OH)}_2$: Magnesium hydroxide

  • $\text{NH}_4\text{F}$: Ammonium fluoride

  • $\text{CaCO}_3$: Calcium carbonate

• Oxyanion Naming Conventions:

  • "per-…-ate": Chemical species possessing the greatest number of oxygen atoms in a related family.

  • "-ate": The primary reference base polyatomic ion.

  • "-ite": Species containing one less oxygen atom than the "-ate" form.

  • "hypo-…-ite": Species containing the fewest oxygen atoms (two fewer than the "-ate" form).

• Multivalent Cation Systems:

  • Certain transition and post-transition metals possess the ability to form stable ions with multiple oxidation states (charges).

  • Stock System: A capitalized Roman numeral enclosed in parentheses is inserted immediately following the metal name to explicitly indicate its ionic charge.

  • Formulas from Names: Use the Roman numeral to balance total electrical charges to zero.

  • Examples:

  • $\text{FeCl}_2$: Iron(II) chloride ($\text{Fe}^{2+}$)

  • $\text{FeCl}_3$: Iron(III) chloride ($\text{Fe}^{3+}$)

  • $\text{CuSO}_4$: Copper(II) sulfate ($\text{Cu}^{2+}$)

  • $\text{CoCl}_2$: Cobalt(II) chloride ($\text{Co}^{2+}$)

  • $\text{Fe}_2\text{S}_3$: Iron(III) sulfide ($\text{Fe}^{3+}$)

• Hydrated Ionic Salts:

  • Hydrate: An ionic compound possessing a precise, fixed stoichiometric ratio of water molecules incorporated within its solid crystalline matrix.

  • Anhydrous: The salt remaining after water of hydration has been removed via thermal driving or desiccation.

  • Physical Manifestation: Dehydration frequently alters physical properties, particularly compound coloration.

  • Formula Representation: A centered dot ($\bullet$) separates the base salt formula and the stoichiometric water molecule count (e.g., $\text{CuSO}_4\bullet 5\text{H}_2\text{O}$, $\text{FeSO}_4\bullet 7\text{H}_2\text{O}$, $\text{Na}_2\text{CO}_3\bullet 10\text{H}_2\text{O}$).

  • Hydrate Nomenclature Protocol: State the full ionic compound name, append the term "-water", and state the molar hydration ratio as $(1/n)$, where $n$ denotes the coefficient of water.

  • Hydrate Example: $\text{Ba(OH)}_2\bullet 8\text{H}_2\text{O}$ is named barium hydroxide-water $(1/8)$.

Nomenclature and Formulas of Molecular Compounds and Acids

• Characteristics of Molecular Compounds:

  • Consist of two or more non-metal elements bonded via covalent electron sharing.

  • Prefixes dictate the exact quantity of each atom present in the discrete molecule.

  • Formulas are written reflecting actual molecular atom counts and are never simplified or reduced.

• Elemental Molecular Forms:

  • Monatomic: Elemental metals, Group 18 noble gases, and pure elemental carbon ($\text{C}$).

  • Diatomic Elements: Distinct elements that naturally exist as two-atom molecules when uncombined: $\text{H}_2$, $\text{N}_2$, $\text{O}_2$, $\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$.

  • Polyatomic Elements: Elemental phosphorus exists naturally as $\text{P}_4$; elemental sulfur exists naturally as $\text{S}_8$.

• Greek Numerical Prefixes:

  • $1 = \text{mono-}$

  • $2 = \text{di-}$

  • $3 = \text{tri-}$

  • $4 = \text{tetra-}$

  • $5 = \text{penta-}$

  • $6 = \text{hexa-}$

  • $7 = \text{hepta-}$

  • $8 = \text{octa-}$

  • $9 = \text{nona-}$

  • $10 = \text{deca-}$

• Molecular Nomenclature Rules:

  • 1. State the full name of the first non-metal element.

  • 2. The prefix "mono-" is never applied to the first element in the compound name.

  • 3. Change the terminal suffix of the second element to "-ide".

  • 4. Apply appropriate numerical prefixes to both elements to denote exact atom counts.

  • 5. Elide double vowels when "mono-" or "tetra-" through "deca-" is attached to oxygen (e.g., carbon monoxide, not carbon monooxide).

  • Binary Molecular Examples:

  • $\text{PCl}_3$: Phosphorus trichloride

  • $\text{Si}_2\text{H}_6$: Disilicon hexahydride

  • $\text{CO}$: Carbon monoxide

  • $\text{CO}_2$: Carbon dioxide

  • $\text{N}_2\text{O}_4$: Dinitrogen tetroxide

• Molecular Formula Conversions:

  • Sulfur trioxide: $\text{SO}_3$

  • Iodine heptafluoride: $\text{IF}_7$

  • Tetranitrogen hexaoxide: $\text{N}_4\text{O}_6$

  • Dibromine trioxide: $\text{Br}_2\text{O}_3$

  • Disulfur decafluoride: $\text{S}_2\text{F}_{10}$

  • Tetranitrogen monoxide: $\text{N}_4\text{O}$

  • Pentasulfur hexanitride: $\text{S}_5\text{N}_6$

  • Iodine pentaoxide: $\text{IO}_5$

• Memorized Common Molecular Compounds:

  • Water: $\text{H}_2\text{O}$

  • Ammonia: $\text{NH}_3$

  • Methane: $\text{CH}_4$

  • Hydrogen peroxide: $\text{H}_2\text{O}_2$

  • Hydrogen sulfide: $\text{H}_2\text{S}$

  • Ethanol: $\text{C}_2\text{H}_5\text{OH}$

  • Glucose: $\text{C}_6\text{H}_{12}\text{O}_6$

  • Sucrose: $\text{C}_{12}\text{H}_{22}\text{O}_{11}$

  • Acetylene: $\text{C}_2\text{H}_2$

  • Propane: $\text{C}_3\text{H}_8$

• Nomenclature System for Inorganic and Organic Acids:

  • Chemical Identification: Inorganic acids present ionizable hydrogen written first in the chemical formula (e.g., $\text{HCl}(aq)$). Organic carboxylic acids contain the terminal carboxyl group, written as $-\text{COOH}$.

  • Transition Rules:

  • Anion ending in "-ide" converts to: "hydro-" + [anion root] + "-ic acid"

  • Anion ending in "-ate" converts to: [anion root] + "-ic acid"

  • Anion ending in "-ite" converts to: [anion root] + "-ous acid"

  • Acid Examples:

  • $\text{HF}(aq)$: Hydrofluoric acid

  • $\text{HCl}(aq)$: Hydrochloric acid

  • $\text{HBr}(aq)$: Hydrobromic acid

  • $\text{HClO}_4(aq)$: Perchloric acid

  • $\text{HClO}_2(aq)$: Chlorous acid

  • $\text{HNO}_3$: Nitric acid

  • $\text{HNO}_2$: Nitrous acid

  • $\text{H}_2\text{SO}_4$: Sulfuric acid

  • $\text{H}_2\text{SO}_3$: Sulfurous acid

  • $\text{H}_2\text{S}(aq)$: Hydrosulfuric acid

  • Root Spelling Conventions:

  • Phosphorus acids incorporate the root syllable "phosphor-" (e.g., phosphoric acid).

  • Sulfur acids incorporate the root syllable "sulfur-" (e.g., sulfuric acid).

  • Common Organic Acids:

  • Ethanoic acid (acetic acid): $\text{CH}_3\text{COOH}$

  • Methanoic acid (formic acid): $\text{HCOOH}$

Comparative Physical and Chemical Properties of Matter

• Ionic versus Molecular Compounds:

  • Structural Makeup:

  • Ionic: Three-dimensional crystal lattice composed of alternating positively charged cations and negatively charged anions.

  • Molecular: Independent, discrete units of neutral molecules bound internally by covalent electron sharing.

  • Standard State at Room Temperature ($25\,^\circ\text{C}$):

  • Ionic: Always crystalline solids.

  • Molecular: Can exist as solids, liquids, or gases.

  • Thermal Transition Points (Melting and Boiling Points):

  • Ionic: Exceptionally high melting and boiling points due to strong electrostatic ionic forces requiring high thermal energy to disrupt.

  • Molecular: Relatively low melting and boiling points due to weak intermolecular attractions.

  • Mechanical Integrity:

  • Ionic: Extremely hard and brittle (impact forces shift like charges adjacent to one another, causing lattice fracture).

  • Molecular: Softer textures and considerably more flammable.

  • Electrical Conductivity:

  • Ionic: In solid state, ions are locked rigidly in position and cannot conduct electricity; when molten (liquid) or dissolved in aqueous solution, the dissociated ions become mobile charge carriers and conduct electricity efficiently.

  • Molecular: Do not dissociate into charged ions in solution; remain neutral molecules and cannot conduct electricity in any state.

• Solution Conductivity Determinations:

  • Pure Water: Extremely poor electrical conductor because of near-zero self-ionization.

  • Electrolytic Behavior:

  • $\text{LiF}(aq)$: Dissociates into $\text{Li}^+$ and $\text{F}^-$ ions; conducts electricity.

  • $\text{C}_6\text{H}_{12}\text{O}_6(aq)$: Molecular compound that dissolves as intact, neutral molecules; does not conduct electricity.

  • $\text{NH}_4\text{Cl}(aq)$: Dissociates into $\text{NH}_4^+$ and $\text{Cl}^-$ ions; conducts electricity.

• Aqueous Solubility Rules:

  • State Labels: $(aq)$ indicates a solute dissolved in aqueous solution; $(s)$ indicates a precipitate or undissolved solid.

  • High Solubility Species: Compounds containing alkali metal cations (Group 1) or ammonium ions ($\text{NH}_4^+$) are consistently soluble in water.

  • General Determination: Reference solubility tables to verify whether ionic combinations form precipitates ($(s)$) or soluble species ($(aq)$).

  • Solubility Examples:

  • Barium sulfate ($\text{BaSO}_4$): Slightly soluble / precipitate, designated as $\text{BaSO}_4(s)$.

  • Iron(III) sulfate ($\text{Fe}_2(\text{SO}_4)_3$): Soluble salt, designated as $\text{Fe}_2(\text{SO}_4)_3(aq)$.

Water Chemistry and the Aqueous State

• Molecular Polarity:

  • A polar molecule features an uneven distribution of electron density across its chemical bonds, resulting in distinct partial charges.

  • Molecular Geometry and Charge of Water ($\text{H}_2\text{O}$):

  • Highly electronegative oxygen bears a partial negative charge ($\delta-$).

  • Hydrogen atoms bear a partial positive charge ($\delta+$).

  • Hydrogen Bonding: Intermolecular electrostatic attractions between the partial positive hydrogen of one water molecule and the partial negative oxygen of a neighboring molecule.

• Seven Distinct Physical Properties of Water:

  • 1. High Melting and Boiling Points: Strong intermolecular hydrogen bonds require significant thermal kinetic energy to sever compared to similar-sized hydrides.

  • 2. Density Inversion of Ice: Liquid water expands as it freezes because hydrogen bonds form an open hexagonal crystalline matrix, causing solid ice to be less dense than liquid water and float.

  • 3. High Specific Heat Capacity: Water absorbs or releases large quantities of thermal energy before changing temperature, moderating environmental temperatures.

  • 4. Cohesion: Attractive intermolecular forces causing water molecules to adhere strongly to one another.

  • 5. Adhesion: Intermolecular attractions enabling water molecules to bind to other polar or charged surfaces.

  • 6. Capillary Action: Combined cohesive and adhesive forces that pull water upward against gravity through narrow spaces and conduits.

  • 7. High Surface Tension: Water molecules along the air-liquid boundary form strong cohesive bonds laterally and downward, creating a resilient, skin-like surface.

  • Thermal Density Maximum: Water attains its maximum absolute density at approximately $4\,^\circ\text{C}$.

Chemical Classification of Solutions: Acids, Bases, and Unknowns

• Acidic Substances:

  • Taste: Characteristic sour taste.

  • Value on the $pH$ Scale: Measured $pH < 7$.

  • Acid-Base Indicator Reaction: Converts blue litmus paper to a distinct red color.

  • Conductivity: Conducts electric current in aqueous media through dissociated hydronium and counter-ions.

  • Metal Reactivity: Reacts vigorously with active metals to generate gaseous hydrogen ($\text{H}_2$).

  • Chemical Identification: Formulas generally start with ionizable hydrogen ($\text{H}$) or contain an organic carboxyl group ($-\text{COOH}$).

• Basic Substances (Alkalis):

  • Taste: Characteristic bitter taste.

  • Texture: Characteristically slippery to the touch.

  • Value on the $pH$ Scale: Measured $pH > 7$.

  • Acid-Base Indicator Reaction: Converts red litmus paper to a distinct blue color.

  • Conductivity: Conducts electric current in aqueous media through dissociated metallic cations and mobile hydroxide anions.

  • Chemical Identification: Formulas frequently pair a metallic cation or ammonium with a hydroxide ion ($\text{OH}^-$, e.g., $\text{NaOH}$, $\text{KOH}$).

• Neutral Substances:

  • Measured $pH$ is exactly $7$.

  • No color transition observed on either red or blue litmus papers.

  • Laboratory Safety Rule: Tasting or directly handling chemical reagents in the laboratory is strictly prohibited.

• Analytical Identification Protocol for Unknown Solutions:

  • Calcium bromide ($\text{CaBr}_2$): Ionic; conducts electricity; neutral on litmus tests.

  • Methanol ($\text{CH}_3\text{OH}$): Molecular neutral; contains covalent hydroxyl group that does not ionize; non-conductive; neutral on litmus tests.

  • Hydrofluoric acid ($\text{HF}(aq)$: Acidic; conducts electricity; converts blue litmus red.

  • Potassium hydroxide ($\text{KOH}$): Basic; conducts electricity; converts red litmus blue.

  • Analytical Workflow:

  • Step 1: Conductivity testing to isolate ionic/acidic/basic solutions from neutral non-conductive molecular species.

  • Step 2: Litmus testing to determine if conductive solutions are acidic (blue to red), basic (red to blue), or neutral salts (no color change).

Controlled and Hazardous Substances in Society

• Targeted Toxicological Compounds:

  • Fentanyl ($\text{C}_{22}\text{H}_{28}\text{N}_2\text{O}$):

  • An extremely potent synthetic opioid analgesic.

  • Physiological Danger: High risk of fatal respiratory depression and arrest.

  • Societal Burden: Drives epidemic substance addiction, healthcare system strain, and community crises.

  • Nicotine ($\text{C}_{10}\text{H}_{14}\text{N}_2$):

  • A powerful and addictive stimulant alkaloid.

  • Neurological Harm: Directly impairs adolescent neurological development and synaptic wiring.

  • Exposure Hazards: E-cigarette aerosol (vaping) introduces toxic chemical byproducts, aldehydes, and heavy metals.

  • Tetrahydrocannabinol / THC ($\text{C}_{21}\text{H}_{30}\text{O}_2$):

  • The principal psychoactive cannabinoid found in cannabis.

  • Cognitive Impairment: Induces deficits in short-term working memory, executive learning, and psychomotor coordination.

  • Developmental Risk: Early adolescent exposure can permanently alter brain neurochemistry and maturation.

• Workplace and Academic Safety Compliance:

  • SDS Review: Access and understand the specific Safety Data Sheet for any chemical prior to handling.

  • Pictogram Recognition: Recognize and react appropriately to standardized WHMIS symbols.

  • Protective Shielding: Wear calibrated personal protective equipment (PPE) to block inhalation, ingestion, or dermal exposure.

  • Handling and Disposal: Adhere to regulated safe chemical handling, storage, neutralization, and hazardous waste disposal guidelines.

• Analytical Framework for Societal Impact Evaluation:

  • 1. Explicitly identify and define the targeted chemical substance.

  • 2. Formulate and elaborate on two separate real-world societal impacts.

  • 3. Provide sound logical reasoning and supporting scientific evidence for each impact.

  • 4. Accurately cite all primary literature and external references utilized.

Fundamentals of Chemical Reactions and Energy Changes

• General Characteristics of Chemical Reactions:

  • Chemical reactions involve the transformation of original reactants into chemically distinct products through the breaking and forming of chemical bonds.

  • Physical vs. Chemical Changes:

  • Physical Change: A transformation in the state, shape, or outward physical appearance of matter without changing its chemical composition (e.g., boiling liquid water into steam).

  • Chemical Change: A process where one or more new chemical substances are generated with distinct chemical and physical properties (e.g., thermal denaturation and coagulation when cooking an egg).

  • Clarification: Homogeneous mixing of molten metals to manufacture an alloy represents a physical change, not a chemical change.

• Everyday and Industrial Chemical Transformations:

  • Photosynthesis: Plants harness solar radiation to convert carbon dioxide and water into glucose and oxygen gas:

  • $6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$

  • Chemical Leavening Agents:

  • Baking Soda ($\text{NaHCO}_3$): Requires addition of an external dietary acid to trigger production of carbon dioxide ($\text{CO}_2$) gas bubbles that leaven baked goods.

  • Baking Powder: A manufactured blend containing both dry sodium bicarbonate and a dry acid; the leavening reaction is triggered simply by adding moisture.

  • Water Dissociation (Electrolysis): Electrical energy splits water molecules into elemental hydrogen and oxygen:

  • $2\text{H}_2\text{O} + \text{energy} \rightarrow 2\text{H}_2 + \text{O}_2$

  • Silicon Refining: High-temperature chemical extraction reduces silicon dioxide ($\text{SiO}_2$, sand) into high-purity elemental silicon for microchip semiconductor manufacturing.

  • Haber-Bosch Process: High-pressure synthesis of ammonia ($\text{NH}_3$) from elemental nitrogen and hydrogen gases, used in fertilizers, industrial cleaners, and explosives.

  • Culinary Chemistry: Cooking breaks down proteins, creates complex aromatic compounds (Maillard reactions), softens textures, and destroys pathogenic bacteria.

• Chemical Reactions Producing Energy:

  • Electrochemical Batteries: Spontaneous redox reactions generate electron flow, powering portable consumer electronics (e.g., lithium-ion batteries).

  • Coal Combustion: Oxidation of carbon generates thermal energy and carbon dioxide gas:

  • $\text{C} + \text{O}_2 \rightarrow \text{CO}_2 + \text{energy}$

  • Natural Gas Combustion: Methane oxidizes to generate energy, water, and carbon dioxide:

  • $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} + \text{energy}$

  • Cellular Respiration: Aerobic biological oxidation of glucose releases metabolic energy (ATP):

  • $\text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2 \rightarrow 6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy}$

  • Environmental Impact: Industrial combustion of carbon-based fossil fuels releases large amounts of $\text{CO}_2$, contributing to global climate change.

• Seven Empirical Indicators of Chemical Change:

  • 1. Formation of a Chemically Distinct Substance: Confirmed emergence of new chemical bonds and properties (the most essential evidence).

  • 2. Irreversibility: The transformation cannot be readily reversed through physical means.

  • 3. Gas Production: Visible effervescence, bubbling, or vapor generation without boiling.

  • 4. Precipitate Formation: An insoluble solid product crashes out of a mixture of clear aqueous solutions.

  • 5. Odour Change: Generation of a new, distinct scent.

  • 6. Permanent Colour Change: A lasting change in pigment or optical absorption.

  • 7. Temperature / Thermal Energy Change: Spontaneous release or absorption of heat energy.

  • Caveat: A single indicator does not guarantee a chemical change has occurred; for example, boiling water exhibits bubbles and temperature changes but is purely a physical change of state.

• Enthalpy Classifications of Reactions:

  • Exothermic Reactions:

  • Net release of thermal energy into the surrounding environment.

  • Causes temperature increases in surrounding materials.

  • Energy term is written on the product side (right side) of the chemical equation:

    • $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O} + \text{energy}$

  • Primary Examples: Combustion reactions and cellular respiration.

  • Endothermic Reactions:

  • Net absorption of thermal energy from the surrounding environment.

  • Causes temperature drops in surrounding materials.

  • Energy term is written on the reactant side (left side) of the chemical equation:

    • $6\text{CO}_2 + 6\text{H}_2\text{O} + \text{energy} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$

  • Primary Examples: Photosynthesis and the electrolysis/decomposition of water.

Classification and Prediction of Chemical Reaction Types

• The Five Primary Reaction Classifications:

  • 1. Formation (Synthesis):

  • Two or more simple reactants combine to produce a single, more complex compound.

  • Generalized Pattern: $\text{A} + \text{B} \rightarrow \text{AB}$

  • Stoichiometric Example: $2\text{Na} + \text{Cl}_2 \rightarrow 2\text{NaCl}$

  • 2. Decomposition:

  • A single complex reactant breaks down into two or more simpler elemental or molecular substances.

  • Generalized Pattern: $\text{AB} \rightarrow \text{A} + \text{B}$

  • Stoichiometric Example: $2\text{H}_2\text{O} \rightarrow 2\text{H}_2 + \text{O}_2$

  • 3. Single Replacement:

  • A reactive uncombined element displaces another element from an aqueous compound.

  • Generalized Pattern: $\text{A} + \text{BC} \rightarrow \text{B} + \text{AC}$

  • Stoichiometric Example: $\text{Cu} + 2\text{AgNO}_3 \rightarrow 2\text{Ag} + \text{Cu(NO}_3)_2$

  • 4. Double Replacement:

  • The ionic species of two distinct aqueous compounds exchange partners to yield two new compounds.

  • Generalized Pattern: $\text{AB} + \text{CD} \rightarrow \text{AD} + \text{CB}$

  • Stoichiometric Example: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

  • 5. Hydrocarbon Combustion:

  • An organic compound composed of hydrogen and carbon (or carbon, hydrogen, and oxygen) reacts exothermically with oxygen gas to yield carbon dioxide, water vapor, and energy.

  • Stoichiometric Example: $\text{C}_3\text{H}_8 + 5\text{O}_2 \rightarrow 3\text{CO}_2 + 4\text{H}_2\text{O}$

• Standard Structural Components of a Chemical Equation:

  • Reactants: Original starting substances placed to the left of the reaction arrow.

  • Products: Resultant substances generated, written to the right of the reaction arrow.

  • Reaction Arrow ($\rightarrow$): Represents the process of chemical change, read as "produces" or "yields".

  • Stoichiometric Coefficients: Whole numbers placed in front of chemical formulas to balance the equation according to conservation of mass.

  • States of Matter Identifiers:

  • $(s)$: Solid state

  • $(l)$: Pure liquid state

  • $(g)$: Gaseous state

  • $(aq)$: Aqueous state (substance dissolved in water)

  • Full Equation Example: $2\text{KClO}_3(s) \rightarrow 2\text{KCl}(s) + 3\text{O}_2(g)$

• Chemical Equation Formats:

  • Word Equation: Identifies reactants and products strictly by full chemical name:

  • $\text{Zinc} + \text{hydrobromic acid} \rightarrow \text{zinc bromide} + \text{hydrogen}$

  • Skeleton Equation: Replaces chemical names with corresponding formulas but does not balance atom counts:

  • $\text{Zn} + \text{HBr} \rightarrow \text{ZnBr}_2 + \text{H}_2$

  • Balanced Chemical Equation: Applies lowest whole-number coefficients to ensure atom counts on each side are equal:

  • $\text{Zn} + 2\text{HBr} \rightarrow \text{ZnBr}_2 + \text{H}_2$

Writing, Balancing, and Predicting Chemical Equations

• Law of Conservation of Mass:

  • In any closed system, matter can neither be created nor destroyed during a chemical reaction.

  • Quantitative Consequence: The total combined mass of the initial reactants must equal the total mass of the products formed.

  • Microscopic Meaning: Chemical transformations simply break, form, and rearrange chemical bonds; the total count of each atomic species remains unchanged.

• Method for Balancing Chemical Equations:

  • 1. Formulate the accurate, charge-balanced skeleton equation.

  • 2. Inventory each atom of every element on both reactant and product sides.

  • 3. Adjust stoichiometric coefficients in front of formulas until atom counts match on both sides.

  • 4. Recount and verify the balance across all elements.

  • 5. Reduce coefficients to the lowest possible whole-number ratio.

  • Critical Rule: Never alter chemical subscripts to balance an equation; changing subscripts changes the chemical identity of the compound.

• Step-by-Step Balancing Example:

  • Unbalanced Skeleton: $\text{H}_2 + \text{O}_2 \rightarrow \text{H}_2\text{O}$

  • Step 1 (Balance Oxygen): Oxygen counts are $2$ on the left and $1$ on the right. Insert coefficient $2$ before water: $\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

  • Step 2 (Balance Hydrogen): Hydrogen counts are now $2$ on the left and $4$ on the right. Insert coefficient $2$ before elemental hydrogen: $2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O}$

  • Final Verification: Reactants ($4\,\text{H}$, $2\,\text{O}$); Products ($4\,\text{H}$, $2\,\text{O}$). The equation is balanced.

• Handling Polyatomic Groups During Balancing:

  • Treat polyatomic ions as single intact units if they appear unchanged on both sides of the reaction arrow.

  • Unbalanced: $\text{AlI}_3 + \text{AgNO}_3 \rightarrow \text{AgI} + \text{Al(NO}_3)_3$

  • Balancing Nitrate: Nitrate ($\text{NO}_3^-$) has $1$ unit on the left and $3$ on the right; add coefficient $3$ to $\text{AgNO}_3$: $\text{AlI}_3 + 3\text{AgNO}_3 \rightarrow \text{AgI} + \text{Al(NO}_3)_3$

  • Balancing Silver and Iodine: Add coefficient $3$ to $\text{AgI}$:

  • $\text{AlI}_3 + 3\text{AgNO}_3 \rightarrow 3\text{AgI} + \text{Al(NO}_3)_3$

• Elemental Formulas in Isolation:

  • Diatomic Elements: When written uncombined as pure elements, seven elements must always be written with subscript $2$:

  • $\text{H}_2$, $\text{N}_2$, $\text{O}_2$, $\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$

  • Mnemonic Device: HOBrFINCl

  • Polyatomic Elements: Pure phosphorus is written as $\text{P}_4$; pure sulfur is written as $\text{S}_8$.

• Predictive Workflow for Chemical Reactions:

  • 1. Classify the reactant types (formation, decomposition, single replacement, double replacement, or combustion).

  • 2. Apply the general reaction pattern to determine product identities.

  • 3. Deduce correct chemical formulas for all products by balancing ionic charges and accounting for diatomic elements.

  • 4. In double replacement reactions, consult solubility guidelines to identify if an insoluble solid precipitate ($(s)$) forms.

  • 5. Apply coefficients to balance the final equation.

The Mole Concept, Molar Mass, and Stoichiometry

• Definition of the Mole:

  • The mole ($\text{mol}$) is the SI base unit used to measure the chemical amount of a substance.

  • Avogadro's Number: $1\,\text{mol} = 6.02 \times 10^{23}\,\text{particles}$.

  • Particle Types: Can refer to individual atoms, covalent molecules, or ionic formula units.

  • Practical Purpose: Bridges submicroscopic particle counts with measurable macroscopic masses in grams.

  • Particle Count Scaling: $1\,\text{mol}$ of sodium contains $6.02 \times 10^{23}$ sodium atoms; $2\,\text{mol}$ of sodium contains $1.204 \times 10^{24}$ sodium atoms.

• Molar Equivalence and Mass Disparities:

  • Exactly one mole of any pure substance always contains the exact same number of constituent particles ($6.02 \times 10^{23}$).

  • One mole of different elements or compounds possess different masses because their constituent atoms have different atomic masses.

  • Elemental Molar Mass Comparisons:

  • Hydrogen ($\text{H}$): $1.01\,\text{g/mol}$

  • Argon ($\text{Ar}$): $39.95\,\text{g/mol}$

  • Copper ($\text{Cu}$): $63.55\,\text{g/mol}$

• Macroscopic Interpretation of Chemical Equations:

  • Balanced reaction: $2\text{C}(s) + \text{O}_2(g) \rightarrow 2\text{CO}(g)$

  • Microscopic scale: $2$ carbon atoms react with $1$ oxygen molecule to yield $2$ carbon monoxide molecules.

  • Macroscopic molar scale: $2\,\text{mol}$ of carbon atoms reacts with $1\,\text{mol}$ of diatomic oxygen molecules to yield $2\,\text{mol}$ of carbon monoxide molecules.

• Calculating Molar Mass ($M$):

  • Molar mass ($M$) is the mass of one mole of a substance, expressed in units of $\text{g/mol}$.

  • 1. Identify and count every atom in the chemical formula.

  • 2. Multiply the atomic mass of each element by its atom count.

  • 3. Sum these mass values to determine total molar mass.

  • Calculation 1 — Water ($\text{H}_2\text{O}$):

  • $\text{H} = 2 \times 1.01 = 2.02\,\text{g/mol}$

  • $\text{O} = 1 \times 16.00 = 16.00\,\text{g/mol}$

  • Total Molar Mass = $2.02 + 16.00 = 18.02\,\text{g/mol}$

  • Calculation 2 — Aluminium Oxide ($\text{Al}_2\text{O}_3$):

  • $\text{Al} = 2 \times 26.98 = 53.96\,\text{g/mol}$

  • $\text{O} = 3 \times 16.00 = 48.00\,\text{g/mol}$

  • Total Molar Mass = $53.96 + 48.00 = 101.96\,\text{g/mol}$

  • Calculation 3 — Calcium Phosphate ($\text{Ca}_3(\text{PO}_4)_2$):

  • Subscripts outside parentheses multiply all atoms inside the parentheses.

  • $\text{Ca} = 3 \times 40.08 = 120.24\,\text{g/mol}$

  • $\text{P} = 2 \times 30.97 = 61.94\,\text{g/mol}$

  • $\text{O} = 8 \times 16.00 = 128.00\,\text{g/mol}$

  • Total Molar Mass = $120.24 + 61.94 + 128.00 = 310.18\,\text{g/mol}$

  • Calculation 4 — Diatomic Chlorine ($\text{Cl}_2$):

  • Molar mass = $2 \times 35.45 = 70.90\,\text{g/mol}$

• Mathematical Conversions Between Mass and Moles:

  • Defined Mathematical Variables:

  • $m$: Sample mass expressed in grams ($\text{g}$

  • $n$: Amount of chemical substance expressed in moles ($\text{mol}$

  • $M$: Substance molar mass expressed in grams per mole ($\text{g/mol}$

  • Operational Formulas:

  • Converting grams to moles: $n = \frac{m}{M}$

  • Converting moles to grams: $m = n \times M$

  • Worked Conversion Example 1 (Mass to Moles):

  • Determine the moles present in $150\,\text{g}$ of calcium chloride ($\text{CaCl}_2$).

  • Step 1: Calculate molar mass: $M = 40.08 + 2(35.45) = 110.98\,\text{g/mol}$

  • Step 2: Compute moles: $n = \frac{150\,\text{g}}{110.98\,\text{g/mol}} = 1.35159\dots\,\text{mol}$

  • Final Result: Approximately $1.35\,\text{mol}$

  • Worked Conversion Example 2 (Moles to Mass):

  • Determine the mass in grams of $0.25\,\text{mol}$ of calcium hydroxide ($\text{Ca(OH)}_2$).

  • Step 1: Calculate molar mass: $M = 40.08 + 2(16.00 + 1.01) = 74.10\,\text{g/mol}$

  • Step 2: Compute mass: $m = 0.25\,\text{mol} \times 74.10\,\text{g/mol} = 18.525\,\text{g}$

  • Final Result: Rounding to two significant digits yields $19\,\text{g}$

• Mass Conservation Stoichiometric Verification:

  • In a chemical reaction, the total mass of the reactants must equal the total mass of the products:

  • $\text{Total mass of reactants} = \text{Total mass of products}$

  • Verification Example (Methane Combustion):

  • Reaction: $\text{CH}_4 + 2\text{O}_2 \rightarrow \text{CO}_2 + 2\text{H}_2\text{O}$

  • Reactant masses: $\text{CH}_4 = 16.05\,\text{g}$, $2\text{O}_2 = 64.00\,\text{g}$; total reactant mass = $16.05 + 64.00 = 80.05\,\text{g}$

  • Product masses: $\text{CO}_2 = 44.01\,\text{g}$, $2\text{H}_2\text{O} = 36.04\,\text{g}$; total product mass = $44.01 + 36.04 = 80.05\,\text{g}$

  • Both sides possess an identical mass of $80.05\,\text{g}$.

  • Practice Example:

  • A sample of $28.02\,\text{g}$ of nitrogen reacts completely with $32.00\,\text{g}$ of oxygen.

  • Predicted total yield of nitrogen monoxide: $28.02\,\text{g} + 32.00\,\text{g} = 60.02\,\text{g}$, assuming complete reaction without loss.

• Significant Figures and Mathematical Rigor:

  • Final calculated values must reflect the proper number of significant digits.

  • For multiplication and division calculations, the final value is rounded to the same number of significant figures as the measurement with the fewest significant figures.

  • Retain unrounded intermediate values throughout multi-step calculations, rounding only the final reported value.

  • Measurement units ($\text{g}$, $\text{mol}$, $\text{g/mol}$) must accompany every quantitative value.

Review Questions and Answers

• Question 1: What is a mole?

  • Answer: A standardized scientific unit representing an amount equal to exactly $6.02 \times 10^{23}$ particles.

• Question 2: How many particles are in one mole?

  • Answer: $6.02 \times 10^{23}$ particles (Avogadro's number).

• Question 3: What is molar mass?

  • Answer: The mass of one mole of a pure substance, quantitatively expressed with units of $\text{g/mol}$.

• Question 4: How do you calculate the molar mass of a compound?

  • Answer: Multiply the atomic mass of each constituent element by its respective atom count indicated by the subscripts, then sum all values together.

• Question 5: How do you convert grams to moles?

  • Answer: Divide the measured mass in grams by the substance's molar mass in grams per mole ($n = \frac{m}{M}$

• Question 6: How do you convert moles to grams?

  • Answer: Multiply the amount of moles by the substance's molar mass in grams per mole ($m = n \times M$

• Question 7: Why do one mole of sodium and one mole of zinc have different masses?

  • Answer: While both samples contain an identical count of atoms ($6.02 \times 10^{23}$ atoms), individual zinc atoms have a greater mass than individual sodium atoms.

• Question 8: What is the Law of Conservation of Mass?

  • Answer: Matter can neither be created nor destroyed during a chemical reaction, meaning the total combined mass of the initial reactants must equal the total combined mass of the products formed.