Study guide
9.2 Structure and Properties of Matter: Atoms, Elements, Compounds, and Chemical Reactions
California Subject Examinations for Teachers CSET
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1 Chapter 1: The CSET Program and the Multiple Subjects Examination
1.1 The CSET Program: Which Tests a California Teacher Actually Needs 1.2 CSET Multiple Subjects: Test Structure, Scoring, and Test-Day Mechanics
2 Chapter 2: Language Structure, Linguistics, and Non-Written Communication
2.1 Linguistics and Language Structure: Phonology, Morphology, Syntax, Semantics, and Pragmatics 2.2 First and Second Language Acquisition and ELD Instruction 2.3 Non-Written Communication, Listening, Speaking, and Media Analysis
3 Chapter 3: Literacy Development, Reading Comprehension, and Text Complexity
3.1 Foundations of Early Literacy: Print Concepts, Phonemic Awareness, and Phonics 3.2 Reading Fluency and Vocabulary Acquisition 3.3 Reading Comprehension of Literary and Informational Texts 3.4 Text Complexity and Matching Readers to Texts and Tasks
4 Chapter 4: Writing, Conventions, Research, and Literacy Assessment
4.1 The Writing Process, Discourse Modes, and Writing Applications 4.2 Conventions of Language: Grammar, Usage, Mechanics, and Spelling 4.3 Research Practices, Source Credibility, and Presenting Knowledge 4.4 Diagnostic Literacy Assessment, Running Records, and Miscue Analysis
5 Chapter 5: Ancient Civilizations and Medieval World History
5.1 Early River Valley Civilizations and the Foundations of Society 5.2 Classical Mediterranean Civilizations: Ancient Greece and Rome 5.3 Medieval and Early Modern Eras: Byzantium, Islam, Africa, the Americas, and the Renaissance
6 Chapter 6: United States History and Government
6.1 Indigenous North America, Colonial Settlement, and the American Revolution 6.2 Constitutional Foundations, Federalism, and the American System of Government 6.3 Westward Expansion, Sectional Conflict, the Civil War, and Reconstruction 6.4 Industrialization, the Progressive Era, World Wars, and the Civil Rights Movement
7 Chapter 7: California History and Geography
7.1 California Geography, Indigenous Nations, and the Spanish Mission Era 7.2 The Mexican Rancho Period, the Bear Flag Revolt, and the 1849 Gold Rush 7.3 California Statehood, Governance, Water Infrastructure, and Modern Economic Growth
8 Chapter 8: Social Science Skills: Geography, Civics, Economics, and Source Analysis
8.1 Chronological and Spatial Thinking: Timelines, Maps, and Geographic Tools 8.2 Core Concepts in Civics, Economics, and the Social Sciences 8.3 Analyzing Primary and Secondary Sources and Writing Historical Arguments
9 Chapter 9: Scientific Practices and Physical Science
9.1 Scientific Investigation, Engineering Design, and Crosscutting Concepts 9.2 Structure and Properties of Matter: Atoms, Elements, Compounds, and Chemical Reactions 9.3 Forces, Motion, Energy Transformations, and Waves
10 Chapter 10: Life Science and Ecology
10.1 Cell Biology, Cellular Organization, and Metabolic Processes 10.2 Genetics, Heredity, Molecular Biology, and Evolution 10.3 Ecosystems, Ecological Interactions, Energy Flow, and Biogeochemical Cycles
11 Chapter 11: Earth, Water, and Space Science
11.1 Earth's Interior, Plate Tectonics, and Geological Processes 11.2 Atmospheric Structure, Hydrology, Weather Systems, and Climate 11.3 Earth's Water: Oceans, Estuaries, Freshwater Systems, and Tidal Dynamics 11.4 Earth-Sun-Moon Dynamics, the Solar System, and the Universe
12 Chapter 12: Number Sense, Operations, and Fractions
12.1 Number Classifications, Place Value, Properties of Operations, and Integers 12.2 Number Theory, Prime Factorization, GCF/LCM, and Exponents 12.3 Fractions, Decimals, Percentages, and Proportional Problem Solving
13 Chapter 13: Algebra, Functions, and Geometry
13.1 Ratios, Rates, Proportions, and Scaling Relationships 13.2 Algebraic Expressions, Linear Equations, Systems, and Functions 13.3 Geometric Properties, Two-Dimensional Polygons, and Three-Dimensional Solids
14 Chapter 14: Coordinate Geometry, Measurement, Statistics, and Mathematical Practices
14.1 Coordinate Geometry, Linear Graphs, and Geometric Transformations 14.2 Measurement Systems, Dimensional Analysis, and Precision 14.3 Descriptive Statistics, Data Displays, and Probability Theory 14.4 Mathematical Practices: Reasoning, Modeling, Estimation, and Justification
15 Chapter 15: Physical Education: Movement, Fitness, Social Development, and Inclusion
15.1 Motor Development, Fundamental Movement Skills, and Movement Concepts 15.2 Fitness Principles, Exercise Physiology, Nutrition, and Safety 15.3 Self-Image, Social Development, and Personal and Social Responsibility 15.4 Movement Forms, Cultural Traditions, and Inclusive Adapted Physical Education
16 Chapter 16: Human Development: Cognition, Intelligence, Social and Physical Growth
16.1 Cognitive and Linguistic Development: Piaget, Vygotsky, and Information Processing 16.2 Perspectives on Intelligence, Play, and Individual Differences 16.3 Social-Emotional and Moral Development: Erikson, Kohlberg, and Attachment Theory 16.4 Physical Development, Ecological Systems, and Mandated Reporting
17 Chapter 17: Visual and Performing Arts: Dance, Music, and Theatre
17.1 Fundamental Elements of Dance, Choreographic Principles, and Cultural Styles 17.2 Core Elements of Music, Staff Notation, Rhythm, and Auditory Analysis 17.3 Music History, World Traditions, and Elementary Music Pedagogy 17.4 Elements of Theatre, Dramatic Structures, Acting Techniques, and Stagecraft
18 Chapter 18: Visual Art, Arts Integration, and Constructed-Response Mastery
18.1 Elements of Art, Principles of Design, and Visual Analysis 18.2 Art Media, Techniques, World Art History, and Art Criticism 18.3 California Arts Standards and Cross-Curricular Arts Integration 18.4 CSET Constructed-Response Strategy, Scoring Scale, and Domain-Specific Exemplars
9.2 Structure and Properties of Matter: Atoms, Elements, Compounds, and Chemical Reactions
Key Takeaways Explain This Topic AI
Atoms consist of a dense central nucleus (protons and neutrons) surrounded by an electron cloud; the atomic number (number of protons) defines element identity, while outer valence electrons dictate chemical reactivity.
The Periodic Table arranges elements by increasing atomic number into periods (electron energy levels) and groups/families (shared valence electron counts and similar chemical properties).
Matter is classified into pure substances (elements and chemically bonded compounds in fixed stoichiometric ratios) and mixtures (heterogeneous suspensions/colloids and homogeneous solutions) that can be physically separated.
Phase changes (melting, freezing, vaporization, condensation, sublimation, deposition) are physical transformations driven by thermal energy, during which temperature remains constant due to latent heat.
The Law of Conservation of Mass dictates that atoms cannot be created or destroyed in chemical reactions; total reactant mass equals total product mass, and balanced equations reflect the conservation of matter.
Last updated: August 2026
9.2 Structure and Properties of Matter: Atoms, Elements, Compounds, and Chemical Reactions
CSET Focus: California multiple subjects educators must understand the particulate nature of matter, atomic structure, the organization of the periodic table, phase transitions, and chemical reaction dynamics. Expect exam scenarios requiring you to distinguish between physical and chemical changes, balance chemical equations demonstrating conservation of mass, and evaluate acid-base reactions.
1. Atomic Structure and Subatomic Particles
All matter in the universe is composed of atoms —the fundamental structural units of chemical elements that retain the characteristic properties of that element. [ ELECTRON CLOUD / ORBITALS ]
- Electrons (e⁻) [Charge: -1, Mass: ~0 amu]
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- ┌───────────────────────┐
- │ DENSE ATOMIC NUCLEUS│
- │ • Protons (p⁺) │ [Charge: +1, Mass: 1 amu]
- │ • Neutrons (n⁰) │ [Charge: 0, Mass: 1 amu]
- └───────────────────────┘
- Subatomic Particle Comparison
Subatomic Particle Electrical Charge Relative Mass (amu) Location within Atom Primary Role & Significance
Proton ( p + p^+ p + ) + 1 +1 + 1 (Positive) ≈ 1.0 amu \approx 1.0\text{ amu} ≈ 1.0 amu Inside the central nucleus Defines the Atomic Number ( Z Z Z ) and chemical identity of the element
Neutron ( n 0 n^0 n 0 ) 0 0 0 (Neutral) ≈ 1.0 amu \approx 1.0\text{ amu} ≈ 1.0 amu Inside the central nucleus Stabilizes nucleus via the strong nuclear force; determines the isotope
Electron ( e − e^- e − ) − 1 -1 − 1 (Negative) ≈ 1 1836 amu \approx \frac{1}{1836}\text{ amu} ≈ 1836 1 amu (Negligible) In electron shells/orbitals surrounding nucleus Valence electrons dictate chemical bonding, reactivity, and ion formation
Key Atomic Metrics and Definitions
Atomic Number ( Z Z Z ): The number of protons in the nucleus of an atom. Every carbon atom has exactly 6 protons ( Z = 6 Z=6 Z = 6 ); every gold atom has exactly 79 protons ( Z = 79 Z=79 Z = 79 ). In a neutral atom, the number of electrons equals the number of protons.
Mass Number ( A A A ): The total sum of protons and neutrons in an atom's nucleus ( A = Protons + Neutrons A = \text{Protons} + \text{Neutrons} A = Protons + Neutrons ). Electrons contribute virtually no mass.
Isotopes: Atoms of the same element that have the same atomic number (same number of protons) but different numbers of neutrons , resulting in different mass numbers. For example, Carbon-12 ( 6 p + , 6 n 0 6p^+, 6n^0 6 p + , 6 n 0 ) and Carbon-14 ( 6 p + , 8 n 0 6p^+, 8n^0 6 p + , 8 n 0 ) share identical chemical reactivity, but Carbon-14 is radioactive.
Ions: Electrically charged atoms formed when an atom gains or loses valence electrons:
Cation: A positively charged ion formed when an atom loses one or more electrons (e.g., Na → Na + + e − \text{Na} \to \text{Na}^+ + e^- Na → Na + + e − ). Formed primarily by metals.
Anion: A negatively charged ion formed when an atom gains one or more electrons (e.g., Cl + e − → Cl − \text{Cl} + e^- \to \text{Cl}^- Cl + e − → Cl − ). Formed primarily by nonmetals.
2. The Periodic Table of Elements
The Periodic Table arranges all 118 known elements in order of increasing atomic number ( Z Z Z ), revealing repeating periodic patterns in atomic radius, electronegativity, and chemical reactivity. [Periods (1 to 7)] ──> Horizontal Rows (Reflect the number of electron energy shells)
- [Groups (1 to 18)] ──> Vertical Columns (Reflect the number of valence electrons & chemical behavior)
- Major Elemental Groups / Families
Group / Family Periodic Location Valence Electrons Oxidation State Chemical and Physical Characteristics
Alkali Metals Group 1 (except H) 1 e − 1\text{ }e^- 1 e − + 1 +1 + 1 Extremely reactive, soft, silvery metals; react violently with water to release hydrogen gas and basic hydroxides
Alkaline Earth Metals Group 2 2 e − 2\text{ }e^- 2 e − + 2 +2 + 2 Highly reactive metals; form basic alkaline oxides; burn brightly (e.g., Magnesium in flares)
Transition Metals Groups 3–12 Variable ( 1 – 2 1\text{--}2 1 – 2 ) Variable ( + 1 to + 6 +1\text{ to }+6 + 1 to + 6 ) Hard, dense, lustrous metals with high melting points; excellent thermal/electrical conductors; form colorful compounds
Halogens Group 17 7 e − 7\text{ }e^- 7 e − − 1 -1 − 1 Highly reactive nonmetals; readily gain 1 e − 1\text{ }e^- 1 e − to form halide salts with metals (e.g., NaCl , KBr \text{NaCl}, \text{KBr} NaCl , KBr )
Noble Gases Group 18 8 e − 8\text{ }e^- 8 e − (He has 2) 0 0 0 (Inert) Completely unreactive, stable monatomic gases; full valence octet prevents spontaneous chemical bonding under standard conditions
Metals, Nonmetals, and Metalloids
Metals (Left and Center): Shiny (metallic luster), malleable (can be hammered into thin sheets without shattering), ductile (can be drawn into wires), high density, high melting points, excellent conductors of electricity and thermal energy. Readily lose electrons to form positive cations.
Nonmetals (Upper Right): Dull appearance, brittle in solid state, low density, low melting points, poor electrical and thermal conductors (insulators). Readily gain or share electrons.
Metalloids / Semimetals (Staircase Boundary: B, Si, Ge, As, Sb, Te): Exhibit intermediate physical and chemical properties. Crucially, they function as semiconductors of electricity (their electrical conductivity increases with temperature), providing the material foundation for modern computer microprocessors and solar cells.
3. Classification of Matter: Pure Substances vs. Mixtures
All matter can be classified based on its structural composition and whether it can be physically separated: [ ALL MATTER ]
- │
- ┌─────────────────────────┴─────────────────────────┐
- ▼ ▼
- [ PURE SUBSTANCES ] [ MIXTURES ]
- (Fixed chemical composition) (Physically blended components)
- │ │ │ │
- ▼ ▼ ▼ ▼
- [ ELEMENTS ] [ COMPOUNDS ] [ HETEROGENEOUS ] [ HOMOGENEOUS ]
- (Single type atom) (Bonded elements) (Visible phases) (Uniform solutions)
- (e.g., O₂, Fe) (e.g., H₂O, NaCl) (e.g., Sand+Water) (e.g., Saltwater, Air)
- Pure Substances
Elements: Substances composed of only one type of atom that cannot be chemically broken down into simpler substances (e.g., pure iron Fe \text{Fe} Fe , oxygen gas O 2 \text{O}_2 O 2 , gold Au \text{Au} Au ).
Compounds: Substances composed of two or more different elements chemically bonded in fixed, definite stoichiometric ratios (e.g., water H 2 O \text{H}_2\text{O} H 2 O , table salt NaCl \text{NaCl} NaCl , glucose C 6 H 12 O 6 \text{C}_6\text{H}_{12}\text{O}_6 C 6 H 12 O 6 ). Compounds have unique chemical properties completely distinct from their constituent elements (e.g., toxic green chlorine gas and explosive sodium metal combine to form harmless culinary salt NaCl \text{NaCl} NaCl ). Compounds can only be separated via chemical reactions.
Mixtures
Heterogeneous Mixtures: Components are not uniformly distributed and exhibit visible distinct phases or particulate boundaries.
Suspensions: Heterogeneous mixtures containing large solid particles that settle out upon standing and can be separated by filtration (e.g., muddy pond water, flour in water).
Colloids: Heterogeneous mixtures containing medium-sized particles that remain permanently suspended and scatter light (the Tyndall Effect ; e.g., milk, fog, gelatin, mayonnaise).
Homogeneous Mixtures (Solutions): Components are uniformly mixed at the molecular scale, exhibiting a single uniform phase throughout.
Solute: The substance that is dissolved (present in smaller amount, e.g., sucrose sugar).
Solvent: The dissolving medium (present in greater amount, e.g., liquid water; water is the "universal solvent").
Solutions can exist as liquids (saltwater), gases (ambient atmospheric air: 78 % N 2 , 21 % O 2 78\%\text{ N}_2, 21\%\text{ O}_2 78% N 2 , 21% O 2 ), or solids (alloys like brass: copper + zinc).
Physical Separation Techniques
Because mixtures are formed by physical blending rather than chemical bonding, their components can be separated using differences in physical properties:
Filtration: Separates undissolved solid precipitates from liquids based on particle size.
Distillation: Separates miscible liquids based on differences in their boiling points (e.g., distilling ethanol from water).
Evaporation / Crystallization: Recovers dissolved solid solutes by heating the solution to evaporate the volatile liquid solvent.
Chromatography: Separates pigment or molecular components based on their differential solubility and migration rates across a stationary medium.
Magnetic Separation: Extracts ferromagnetic materials (iron, nickel, cobalt) using magnetic attraction.
4. States of Matter and Phase Changes
Matter exists in four fundamental physical states determined by the balance between intermolecular attractive forces and the kinetic energy (thermal motion) of its particles:
State of Matter Shape Volume Particle Arrangement & Kinetic Motion Compressibility
Solid Definite Definite Tightly packed in fixed geometric lattice; vibrate in place; lowest kinetic energy Incompressible
Liquid Indefinite (Takes container shape) Definite Close together but free to slip, slide, and flow past one another; intermediate kinetic energy Practically Incompressible
Gas Indefinite Indefinite (Expands to fill entire volume) Separated by vast empty distances; move randomly at high velocities; high kinetic energy Highly Compressible
Plasma Indefinite Indefinite Superheated, ionized gas of free electrons and positive nuclei; conducts electricity Highly Compressible
Thermodynamics of Phase Changes
Phase transitions are physical transformations driven by the addition or removal of thermal energy : Solid ⇌ Melting Freezing Liquid ⇌ Vaporization Condensation Gas ⇌ Ionization De-ionization Plasma \text{Solid } \underset{\text{Freezing}}{\overset{\text{Melting}}{\rightleftharpoons}} \text{ Liquid } \underset{\text{Condensation}}{\overset{\text{Vaporization}}{\rightleftharpoons}} \text{ Gas } \underset{\text{De-ionization}}{\overset{\text{Ionization}}{\rightleftharpoons}} \text{ Plasma} Solid Freezing ⇌ Melting Liquid Condensation ⇌ Vaporization Gas De-ionization ⇌ Ionization Plasma
Sublimation: Direct phase transition from solid to gas without entering the liquid state (e.g., dry ice / solid CO 2 \text{CO}_2 CO 2 at room temperature, mothballs).
Deposition: Direct phase transition from gas to solid without passing through a liquid state (e.g., sub-freezing water vapor forming frost on windows).
Boiling vs. Evaporation: Evaporation is a surface-level vaporization occurring at any temperature below the boiling point; boiling occurs throughout the entire bulk liquid when the internal vapor pressure equals the external atmospheric pressure.
Heating Curve Analysis & Latent Heat
When a pure solid is continuously heated, its temperature rises until it reaches its melting point. During the phase change (the flat plateau on a heating curve), the temperature remains completely constant even though thermal energy continues to be added. This energy ( latent heat of fusion or latent heat of vaporization ) is consumed to break intermolecular bonds rather than increasing the kinetic speed (temperature) of the molecules.
5. Physical vs. Chemical Properties and Changes
A central distinction on CSET Subtest II is differentiating between physical and chemical phenomena:
Classification Scientific Definition Key Examples & Markers
Physical Property A characteristic observable or measurable without altering the chemical composition of the substance Density ( D = m / V D = m/V D = m / V ), boiling point, melting point, color, hardness, electrical conductivity, solubility, magnetic permeability
Chemical Property The inherent potential of a substance to undergo a chemical transformation that changes its molecular identity Flammability, reactivity with water or acid, toxicity, susceptibility to oxidation (rusting), heat of combustion
Physical Change A change in form, appearance, or physical state that does not break or create chemical bonds Shredding paper, dissolving sugar in iced tea, melting ice cubes, breaking glass, boiling water into steam
Chemical Change (Reaction) A process where chemical bonds are broken and reformed, rearranging atoms into entirely new substances Iron rusting ( 4 Fe + 3 O 2 → 2 Fe 2 O 3 4\text{Fe} + 3\text{O}_2 \to 2\text{Fe}_2\text{O}_3 4 Fe + 3 O 2 → 2 Fe 2 O 3 ), baking bread, combusting wood, neutralising acid with baking soda
Cardinal Signs of a Chemical Reaction
Elementary students can identify chemical reactions by observing five characteristic physical markers:
Production of Gas: Visible bubbling, fizzing, or effervescence in the absence of external boiling (e.g., dropping antacid tablets in water).
Precipitate Formation: An insoluble solid suddenly separates and falls out of a mixture when two clear liquid solutions are combined.
Unprovoked Temperature Change: The mixture spontaneously releases heat ( exothermic reaction , feels hot) or absorbs heat ( endothermic reaction , feels cold) without external heating or cooling.
Permanent Color Change: An unexpected, irreversible alteration in color (e.g., copper turning green patina or sliced apples browning due to enzyme oxidation).
Emission of Light or Odor: Production of a flame, luminescence (glow sticks), or distinct pungent smell.
- 6. Chemical Reactions, Conservation of Mass, and Acid-Base Chemistry
- The Law of Conservation of Mass
Formulated by Antoine Lavoisier, the Law of Conservation of Mass states that in any closed chemical reaction, matter cannot be created or destroyed. The total mass of the reactants must exactly equal the total mass of the products: ∑ Mass of Reactants = ∑ Mass of Products \sum \text{Mass of Reactants} = \sum \text{Mass of Products} ∑ Mass of Reactants = ∑ Mass of Products Number of Atoms of each Element (Reactants) = Number of Atoms of each Element (Products) \text{Number of Atoms of each Element (Reactants)} = \text{Number of Atoms of each Element (Products)} Number of Atoms of each Element (Reactants) = Number of Atoms of each Element (Products)
CSET Scenario Warning: If an open beaker containing vinegar and baking soda loses mass after foaming, mass was not destroyed. Gaseous carbon dioxide ( CO 2 \text{CO}_2 CO 2 ) escaped into the room. If conducted in a sealed, airtight flask, the measured mass before and after the reaction remains completely identical.
Balancing Chemical Equations
Chemical equations use chemical formulas and stoichiometric coefficients to satisfy the conservation of mass: 1 CH 4 (Methane) + 2 O 2 (Oxygen) ⟶ 1 CO 2 (Carbon Dioxide) + 2 H 2 O (Water) 1\text{CH}_4\text{ (Methane)} + 2\text{O}_2\text{ (Oxygen)} \longrightarrow 1\text{CO}_2\text{ (Carbon Dioxide)} + 2\text{H}_2\text{O}\text{ (Water)} 1 CH 4 (Methane) + 2 O 2 (Oxygen) ⟶ 1 CO 2 (Carbon Dioxide) + 2 H 2 O (Water)
Reactants: 1 C , 4 H , 4 O 1\text{ C}, 4\text{ H}, 4\text{ O} 1 C , 4 H , 4 O ⟷ \longleftrightarrow ⟷ Products: 1 C , 4 H , 4 O 1\text{ C}, 4\text{ H}, 4\text{ O} 1 C , 4 H , 4 O . Atoms are conserved.
Five Primary Reaction Types
Synthesis (Combination): A + B → A B A + B \to AB A + B → A B (e.g., 2 Mg + O 2 → 2 MgO 2\text{Mg} + \text{O}_2 \to 2\text{MgO} 2 Mg + O 2 → 2 MgO )
Decomposition: A B → A + B AB \to A + B A B → A + B (e.g., 2 H 2 O → electrolysis 2 H 2 + O 2 2\text{H}_2\text{O} \xrightarrow{\text{electrolysis}} 2\text{H}_2 + \text{O}_2 2 H 2 O electrolysis 2 H 2 + O 2 )
Single Replacement: A + B C → A C + B A + BC \to AC + B A + BC → A C + B (e.g., Zn + 2 HCl → ZnCl 2 + H 2 \text{Zn} + 2\text{HCl} \to \text{ZnCl}_2 + \text{H}_2 Zn + 2 HCl → ZnCl 2 + H 2 )
Double Replacement: A B + C D → A D + C B AB + CD \to AD + CB A B + C D → A D + CB (e.g., AgNO 3 + NaCl → AgCl ( s ) + NaNO 3 \text{AgNO}_3 + \text{NaCl} \to \text{AgCl}_{(s)} + \text{NaNO}_3 AgNO 3 + NaCl → AgCl ( s ) + NaNO 3 )
Combustion: Hydrocarbon + O 2 → CO 2 + H 2 O + Heat \text{Hydrocarbon} + \text{O}_2 \to \text{CO}_2 + \text{H}_2\text{O} + \text{Heat} Hydrocarbon + O 2 → CO 2 + H 2 O + Heat (e.g., cellular respiration or burning propane)
Acids, Bases, and the pH Scale
Acids: Substances that release hydrogen ions ( H + \text{H}^+ H + or hydronium H 3 O + \text{H}_3\text{O}^+ H 3 O + ) in aqueous solutions. Taste sour, corrosive to metals, turn blue litmus paper red ( pH < 7.0 \text{pH} < 7.0 pH < 7.0 ). Examples: hydrochloric acid ( HCl \text{HCl} HCl in stomach), citric acid in citrus fruits, acetic acid in vinegar.
Bases (Alkaline): Substances that produce hydroxide ions ( OH − \text{OH}^- OH − ) in aqueous solutions. Taste bitter, feel slippery, turn red litmus paper blue ( pH > 7.0 \text{pH} > 7.0 pH > 7.0 ). Examples: sodium hydroxide ( NaOH \text{NaOH} NaOH in drain cleaner), bleach, ammonia, antacids.
The pH Scale: A logarithmic scale from 0 to 14 measuring hydrogen ion concentration. Each single-unit shift represents a 10-fold change in acidity (a solution with pH 3 \text{pH } 3 pH 3 is 100 times more acidic than pH 5 \text{pH } 5 pH 5 ). Neutral water has a pH \text{pH} pH of 7.0.
Neutralization Reactions: An acid and a base react stoichimetrically to produce a neutral salt and water: Acid + Base ⟶ Salt + Water \text{Acid} + \text{Base} \longrightarrow \text{Salt} + \text{Water} Acid + Base ⟶ Salt + Water HCl + NaOH ⟶ NaCl + H 2 O \text{HCl} + \text{NaOH} \longrightarrow \text{NaCl} + \text{H}_2\text{O} HCl + NaOH ⟶ NaCl + H 2 O