Comprehensive Study Notes on Chemical Reactions, Acids, Bases, and Laws of Chemical Change

Chemical Changes and Reaction Fundamentals

• Inquiry Lab - Silver Polishing Experiment:

  • Procedure: Crumple a small sheet (about $20\,\text{cm} \times 20\,\text{cm}$) of aluminum foil into a ball. Place the aluminum foil ball into a dishpan filled with soapy water. Position a tarnished piece of silver in the dishpan so that it is submerged and touching the aluminum foil. Wait for about $30\,\text{minutes}$ and observe the silver.
  • Key Concept: The removal of tarnish demonstrates a chemical reaction resulting from contact between the silver, aluminum foil, and soapy water solution.

• Definition of Chemical Change and Reaction:

  • Everyday Example: Popcorn kernels transforming into fluffy popcorn when exposed to extreme heat.
  • Definition: A chemical reaction is a process in which one set of substances, called reactants, is converted into a new set of substances, called products.
  • Reactants: The starting materials in a chemical reaction.
  • Products: The new substances formed as a result of the chemical reaction.
  • Net Chemical Equation Representation:     $\text{reactants} \rightarrow \text{products}$

• Real-World and Industrial Applications of Chemical Reactions:

  • Pharmaceuticals and Manufacturing: Producing medicines and essential materials.
  • Transportation: Powering vehicles through combustion reactions, such as burning gasoline.
  • Energy Generation: Generating electricity in power plants through specialized chemical processes.
  • Culinary Arts: Essential for cooking and baking processes.

• Example Reaction: Formation of Sodium Chloride (Table Salt):

  • Reactants:
  • Sodium ($\text{Na}$): A silver-colored solid metal.
  • Chlorine ($\text{Cl}_2$): A yellowish gas.
  • Process: When sodium metal and chlorine gas react, chemical bonds between sodium atoms and among chlorine atoms are broken. The atoms rearrange, forming new chemical bonds between sodium and chlorine atoms.
  • Product: Sodium chloride ($\text{NaCl}$), a white crystalline solid substance.
  • Energy: A specific amount of energy is produced during this reaction.
  • Balanced Chemical Equation:     $2\text{Na}_{(s)} + \text{Cl}_{2(g)} \rightarrow 2\text{NaCl}_{(s)}$

• Four Key Characteristics of Chemical Reactions:

  1. The atoms of the substances rearrange, but the total number of atoms remains identical before and after the reaction.
  2. The original substances lose their identities and cannot be recognized in the newly formed products.
  3. Energy is involved in the chemical reaction, as heat is either absorbed or released.
  4. The total mass of the reactants is strictly equal to the total mass of the products.

• Distinction Between Chemical and Physical Changes:

  • Chemical Change: Involves the transformation of substances into entirely new ones with different chemical compositions and properties.
  • Physical Change: Does not alter the chemical composition of substances; it only changes their physical appearance or state of matter.
  • Example of Physical Change: The melting of ice is a physical change because the chemical composition remains $\text{H}_2\text{O}$.

Indicators of Chemical Reactions

• Definition: Indicators of chemical reactions are observable clues or changes that suggest a chemical transformation is actively taking place.

• Nine Common Indicators:

  1. Color Change: Alteration of color indicates a reaction. Example: Iron rusting transforms from shiny silver to reddish-brown.
  2. Formation of Precipitate: A precipitate is a cloudy or solid material that settles out of a liquid solution when two clear liquids interact. Example: Mixing vinegar and milk produces solid lumps or curds.
  3. Gas Production: Release of bubbles or evolution of gas. Example: Mixing baking soda and vinegar causes fizzing and releases carbon dioxide ($\text{CO}_2$) gas.
  4. Temperature Change: Absorbing heat energy (endothermic) or releasing heat energy (exothermic) causes significant temperature shifts. Example: Burning wood releases heat energy.
  5. Light Emission: Chemical reactions frequently release light. Example: Fireworks emit vibrant colorful light upon ignition.
  6. Odor Change: A noticeable change in smell. Example: The pleasant aroma produced by freshly baked bread or roasted coffee beans.
  7. Change in Taste: Alterations in flavor indicate chemical shifts. Example: The caramelization of sugar during cooking.
  8. Formation of New Substances: Completely new products with distinct properties are created. Example: Hydrogen gas and oxygen gas combining to form water ($\text{H}_2\text{O}$).
  9. Release or Absorption of Energy: Clear transfer of energy, such as exothermic combustion reactions releasing heat/light, or endothermic reactions like dissolving ammonium nitrate in water absorbing heat.

Endothermic and Exothermic Reactions

• Endothermic Reactions:

  • Definition: Chemical reactions that absorb heat or energy from their surrounding environment in order for the reaction to proceed.
  • Key Characteristics:
  • Heat Absorption: Energy is taken in from the surroundings in the form of heat. Reactants possess a lower energy level than the final products.
  • Temperature Decrease: The temperature of the surrounding environment drops as heat is absorbed, causing a sensation of cooling when touching the reaction container.
  • Examples of Endothermic Reactions:
  • Dissolution of ammonium nitrate in water
  • Melting of ice / melting solid salts
  • Evaporation of liquid water
  • Nuclear fission
  • Reaction of water and a strong acid
  • Lighting a candle
  • Cooking an egg
  • Photosynthesis
  • Baking bread
  • Splitting a gas molecule
  • Electrolysis process

• Exothermic Reactions:

  • Definition: Chemical reactions that release heat or energy into their surrounding environment as they proceed.
  • Key Characteristics:
  • Heat Release: Energy is emitted in the form of heat or light. Reactants possess a higher energy level than the final products.
  • Temperature Increase: The temperature of the surrounding environment rises once heat is released, causing a sensation of warmth when touching the reaction container.
  • Examples of Exothermic Reactions:
  • Combustion of fuel / combustion reactions
  • Formation of bonds in chemical compounds
  • Reaction between acids and bases
  • Burning of sugar
  • Snow formation in clouds
  • Rusting of iron
  • Lighting a firework

Chemical Reactions in the Environment

• Overview: Environmental chemical reactions shape natural surroundings, drive ecosystem processes, and can have both beneficial and adverse ecological impacts.

• Five Crucial Environmental Chemical Reactions:

  1. Photosynthesis:
  • Process: Converts carbon dioxide ($\text{CO}_2$) and water ($\text{H}_2\text{O}$) into glucose ($\text{C}_6\text{H}_{12}\text{O}_6$) and oxygen ($\text{O}_2$) in the presence of sunlight.
  • Organisms: Occurs in plants, algae, and certain bacteria.
  • Significance: Provides energy for the plant and releases oxygen into the atmosphere, supporting aerobic life on Earth.
  1. Respiration:
  • Process: Involves the breakdown of glucose and the consumption of oxygen to release carbon dioxide, water, and energy.
  • Organisms: Takes place in living organisms, including plants, animals, and microorganisms.
  • Significance: Maintains the critical balance of oxygen and carbon dioxide within the Earth's atmosphere.
  1. Combustion Reactions:
  • Process: Rapid chemical combination of a fuel source with oxygen to yield heat, carbon dioxide ($\text{CO}_2$), water vapor, and various pollutants.
  • Significance: Responsible for energy production in transportation vehicles, industrial power plants, and natural wildfires.
  1. Ozone Depletion:
  • Process: The release of chlorofluorocarbons (CFCs) and other ozone-depleting pollutants initiates chemical reactions that decompose ozone molecules ($\text{O}_3$) in the stratospheric ozone layer.
  • Consequences: Depletion permits higher levels of harmful ultraviolet (UV) radiation from the sun to penetrate the atmosphere and reach the Earth's surface, posing severe risks to human health and ecosystems.
  1. Biogeochemical Cycles:
  • Process: Transformations of elements and compounds across ecological reservoirs, including the carbon, nitrogen, and phosphorus cycles.
  • Example: Nitrogen fixation converts atmospheric nitrogen ($\text{N}_2$) into usable ammonia ($\text{NH}_3$).

Acid-Base Theories

• Brønsted-Lowry Theory of Acids and Bases (Proton Theory):

  • Proposers: Formulated independently by Danish chemist Johannes Nicolaus Brønsted and English chemist Thomas Martin Lowry.
  • Core Focus: Proton transfer ($\text{H}^+$).
  • Brønsted-Lowry Acid: A proton ($\text{H}^+$) donor.
  • Brønsted-Lowry Base: A proton ($\text{H}^+$) acceptor.
  • Reaction Example:
  • Dissolution of hydrochloric acid in water:       $\text{HCl}_{(aq)} + \text{H}_2\text{O}_{(l)} \rightleftharpoons \text{H}_3\text{O}^+_{(aq)} + \text{Cl}^-_{(aq)}$
  • Mechanism: Hydrochloric acid ($\text{HCl}$) donates a proton to water ($\text{H}_2\text{O}$). Water accepts the proton to become a hydronium ion ($\text{H}_3\text{O}^+$), acting as a conjugate acid. Hydrochloric acid becomes a chloride ion ($\text{Cl}^-$), which acts as a conjugate base because its negative charge allows it to readily accept a proton in the reverse reaction to establish chemical equilibrium.

• Lewis Theory of Acids and Bases:

  • Proposer: Introduced in 1916 by American chemist Gilbert N. Lewis in his publication The Atom and the Molecule.
  • Core Focus: Electron pair transfer and coordinate covalent bond formation.
  • Lewis Acid: An electron pair acceptor. These species possess an electron-deficient center or an empty atomic orbital capable of accepting an electron pair.
  • Examples: Metal cations such as aluminum ion ($\text{Al}^{3+}$), iron (III) ($\text{Fe}^{3+}$), and neutral compounds like boron trifluoride ($\text{BF}_3$).
  • Lewis Base: An electron pair donor. These species typically feature non-bonding lone pairs of electrons that can be shared.
  • Examples: Ammonia ($\text{NH}_3$), water ($\text{H}_2\text{O}$), and hydroxide anions ($\text{OH}^-$).

Properties and Classification of Acids

• Definition: Chemical substances that, when dissolved in water or aqueous solutions, release hydrogen ions ($\text{H}^+$) or protons, increasing the concentration of hydrogen ions and lowering the solution's pH.

• General Properties of Acids:

  • Taste: Sour or tart flavor.
  • Texture: Concentrated acids are corrosive and exhibit a slippery or oily feel upon contact.
  • pH Level: Possess a pH value strictly below 7 ($\text{pH} < 7$). Lower pH values correspond to stronger acids.
  • Conductivity: Readily conduct electricity in aqueous solution due to the presence of free-moving hydrogen ions ($\text{H}^+$).
  • Reactivity: React vigorously with active metals to produce hydrogen gas ($\text{H}_2$) and with carbonates to produce carbon dioxide gas ($\text{CO}_2$). Neutralize basic solutions.
  • Indicator Behavior: Turn blue litmus paper red.

• Classification of Acids:

  • Strong Acids: Completely ionize in water, releasing a high concentration of hydrogen ions ($\text{H}^+$).
  • Hydrochloric acid ($\text{HCl}$)
  • Sulfuric acid ($\text{H}_2\text{SO}_4$)
  • Nitric acid ($\text{HNO}_3$)
  • Hydrobromic acid ($\text{HBr}$)
  • Perchloric acid ($\text{HClO}_4$)
  • Weak Acids: Partially ionize in water, releasing a significantly lower concentration of hydrogen ions ($\text{H}^+$).
  • Acetic acid ($\text{CH}_3\text{COOH}$)
  • Citric acid ($\text{C}_6\text{H}_8\text{O}_7$)
  • Carbonic acid ($\text{H}_2\text{CO}_3$)
  • Phosphoric acid ($\text{H}_3\text{PO}_4$)
  • Formic acid ($\text{HCOOH}$)

• Acidic Environmental Trivia: Rainwater is naturally slightly acidic because atmospheric carbon dioxide dissolves in atmospheric water to produce carbonic acid ($\text{H}_2\text{CO}_3$).

Properties and Classification of Bases

• Definition: Chemical substances that, when dissolved in water or aqueous solutions, release hydroxide ions ($\text{OH}^-$) or negatively charged ions capable of accepting hydrogen ions ($\text{H}^+$), thereby increasing the solution's pH.

• General Properties of Bases:

  • Taste: Bitter flavor.
  • Texture: Slippery or soapy feel to the touch in concentrated forms.
  • pH Level: Possess a pH value strictly above 7 ($\text{pH} > 7$). Higher pH values correspond to stronger bases.
  • Conductivity: Conduct electricity in aqueous solution due to the presence of free-moving hydroxide ions ($\text{OH}^-$).
  • Reactivity: React with fats and oils through a chemical process termed saponification. Neutralize acidic solutions.
  • Indicator Behavior: Turn red litmus paper blue.

• Classification of Bases:

  • Strong Bases: Completely dissociate in water, releasing a high concentration of hydroxide ions ($\text{OH}^-$).
  • Sodium hydroxide ($\text{NaOH}$)
  • Potassium hydroxide ($\text{KOH}$)
  • Barium hydroxide ($\text{Ba(OH)}_2$)
  • Calcium hydroxide ($\text{Ca(OH)}_2$)
  • Weak Bases: Partially dissociate in water, releasing a lower concentration of hydroxide ions ($\text{OH}^-$).
  • Ammonia ($\text{NH}_3$)
  • Methylamine ($\text{CH}_3\text{NH}_2$)
  • Aniline ($\text{C}_6\text{H}_5\text{NH}_2$)
  • Aluminum hydroxide ($\text{Al(OH)}_3$)

Properties and Uses of Salts and Saline Solutions

• Definition: Chemical compounds formed when an acid reacts with a base in a neutralization reaction. They consist of positively charged cations and negatively charged anions bound by ionic bonding, resulting in an electrically neutral compound.

• General Properties of Salts:

  • Taste: Mostly tasteless, though several exhibit a distinctly salty taste.
  • Texture: Typically crystalline solids.
  • pH Level: Neutral, possessing a pH value of 7 ($\text{pH} = 7$) when dissolved in pure water.
  • Conductivity: Conduct electricity in aqueous solution due to free-moving cations and anions.
  • Solubility: Varies widely; some salts are highly soluble in water, whereas others display low solubility.

• Neutralization Reaction Pattern:   $\text{acid} + \text{base} \rightarrow \text{salt} + \text{water}$

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

• Common Salts, Formulas, Common Names, and Uses:

  • Sodium chloride ($\text{NaCl}$)
  • Common Name: Table salt
  • Uses: Food preparation, industrial production of chemical compounds.
  • Potassium chloride ($\text{KCl}$)
  • Common Name: Potassium chloride
  • Uses: Salt substitute for dietary restrictions, specialized medical applications.
  • Magnesium sulfate ($\text{MgSO}_4$)
  • Common Name: Epsom salt
  • Uses: Therapeutic bath salts, clinical laxative.
  • Calcium carbonate ($\text{CaCO}_3$)
  • Common Name: Calcite, chalk
  • Uses: Naturally occurring constituent of limestone and marble; ingredient in antacids.
  • Sodium bicarbonate ($\text{NaHCO}_3$ / listed as $\text{CaCO}$ in table excerpt)
  • Common Name: Baking soda
  • Uses: Culinary cooking, baking, antacid applications.

• Saline Solution Characteristics and Medical Applications:

  • Composition: A precise concentration of sodium chloride ($\text{NaCl}$) dissolved in purified water.
  • Biological Compatibility: Matches the physiological environment of biological cells and human tissue.
  • pH Standard: Features a neutral pH around 7 ($\text{pH} = 7$), ensuring it is neither acidic nor alkaline.
  • Medical Uses: Serves as intravenous (IV) fluids, wound washing/cleansing, contact lens storage solutions, and medical eye drops.

The pH Scale and Hydrogen Ion Concentration

• Definition: The term "pH" stands for "potential of hydrogen". It provides a precise metric for measuring the concentration of hydrogen ions ($\text{H}^+$) in an aqueous solution.

• Scale Scale Ranges and Interpretation:

  • Range: Standard scale spans from 0 to 14.
  • Neutral Point ($\text{pH} = 7$): Indicates that the concentration of hydrogen ions ($\text{H}^+$) is strictly equal to the concentration of hydroxide ions ($\text{OH}^-$).
  • Acidic Region ($\text{pH} < 7$): Indicates a higher concentration of hydrogen ions ($\text{H}^+$).
  • Basic / Alkaline Region ($\text{pH} > 7$): Indicates a lower concentration of hydrogen ions ($\text{H}^+$) and a higher concentration of hydroxide ions ($\text{OH}^-$).

• Logarithmic Nature of the pH Scale:

  • A change of one whole pH unit signifies a tenfold ($10\times$) difference in hydrogen ion ($\text{H}^+$) concentration.
  • Comparative Concentration Scale Examples:
  • A solution of $\text{pH} = 4$ is $10$ times more acidic than a solution of $\text{pH} = 5$
  • A solution of $\text{pH} = 3$ is $100$ times more acidic than a solution of $\text{pH} = 5$
  • A solution of $\text{pH} = 2$ is $1,000$ times more acidic than a solution of $\text{pH} = 5$
  • Directional Changes:
  • Decreasing from $\text{pH} = 7$ to $\text{pH} = 6$ means the solution is $10$ times more acidic.
  • Increasing from $\text{pH} = 7$ to $\text{pH} = 8$ means the solution is $10$ times less acidic ($10$ times more basic).

Chemical Indicators and Their Applications

• Definition: Specialized substances or analytical tools used in chemistry to detect the presence of chemical species, measure acidity/basicity, or determine the endpoint of a chemical reaction.

• Five Primary Types of Chemical Indicators:

  1. Litmus Paper:
  • Prepared using paper impregnated with natural dyes extracted from lichens.
  • Blue Litmus: Turns red under acidic conditions ($\text{pH} < 7$).
  • Red Litmus: Turns blue under basic conditions ($\text{pH} > 7$).
  1. Universal pH Paper:
  • Impregnated with a mixture of dyes displaying distinct color transformations across a broad pH spectrum ranging from 1 to 14.
  1. Phenolphthalein:
  • Remains colorless in acidic solutions ($\text{pH} < 7$).
  • Shifts to a vibrant pink or deep red color in basic solutions ($\text{pH} > 8.2$).
  1. Bromothymol Blue:
  • Appears yellow in acidic solutions.
  • Appears green in neutral solutions.
  • Appears blue in basic solutions.
  1. Methyl Orange:
  • Appears red in acidic solutions, shifts to yellow in neutral solutions, and displays orange/red shades in basic solutions.

• Five Essential Practical Applications of Indicators:

  1. Acid-Base Titrations: Pinpoint the precise stoichiometric endpoint of acid-base neutralization reactions.
  2. Environmental Testing: Evaluate the acidity and overall quality of soil, natural water bodies, and industrial wastewater.
  3. Medical and Biological Research: Measure the pH of blood and fluid samples to diagnose and monitor physiological conditions.
  4. Quality Control: Validate the correct pH level to ensure the safety, efficacy, and quality of pharmaceutical and food products.
  5. Education: Serves as a foundational tool for teaching concepts of chemical reactions, pH mechanics, and laboratory techniques.

Questions & Discussion

• Question: Why are acid-base indicators important?
• Response & Kitchen Chemistry Application:
  • Acid-Base indicators are crucial because they offer a visual tool to confirm the presence of acids or bases and monitor dynamic chemical reactions.
  • DIY Cabbage Indicator Protocol:
  1. Extract a pH-sensitive natural dye from red cabbage by cooking chopped cabbage in boiling water.
  2. Soak coffee filter paper thoroughly inside the extracted red cabbage liquid solution.
  3. Remove and allow the filter paper to dry completely, then cut it into uniform test strips.
  4. Color Indicators: Testing household solutions causes the cabbage paper to change color, turning reddish when exposed to acids and greenish when exposed to bases.

Laboratory Procedures and Chemical Changes

• Safety Protocol: Protective safety goggles must be worn at all times throughout laboratory procedures.

• Required Materials:

  • 1 candle
  • 1 match or lighter
  • Small piece of magnesium ribbon (approx. $2\,\text{cm} - 3\,\text{cm}$)
  • Safety goggles
  • Test tubes
  • Baking soda
  • Vinegar
  • Steel wool or iron nail

• Step-by-Step Reaction Procedures:

  • Synthesis Reaction Procedure:
  1. Place $1\,\text{teaspoon}$ of baking soda into a test tube.
  2. Carefully add $\frac{1}{4}\,\text{cup}$ of vinegar to the test tube.
  3. Observe and document all visual changes.
  • Decomposition Reaction Procedure:
  1. Place a small piece of steel wool or an iron nail into a test tube.
  2. Add $\frac{1}{4}\,\text{cup}$ of vinegar to the test tube.
  3. Observe and document all visual changes.
  • Combustion Reaction Procedure:
  1. Light a candle using a match or lighter.
  2. Observe and describe the physical attributes of the flame.
  3. Record physical and chemical changes in the candle as it burns.
  • Displacement Reaction Procedure:
  1. Place a small piece ($2\,\text{cm} - 3\,\text{cm}$) of magnesium ribbon into a test tube.
  2. Add $\frac{1}{4}\,\text{cup}$ of vinegar to the test tube.
  3. Observe and document all visual changes.

• Five Primary Processes of Chemical Change:

  1. Oxidation: The chemical combination of oxygen with other substances to produce oxides.
  • Example Reaction:        $4\text{Fe} + 3\text{O}_2 + 6\text{H}_2\text{O} \rightarrow 4\text{Fe(OH)}_3$
  1. Reduction: The chemical process opposite to oxidation, characterized by the extraction or removal of oxygen from a compound.
  2. Neutralization: A chemical reaction occurring between an acid and a base that produces a salt and water.
  3. Hydrolysis: The chemical process opposite to neutralization, wherein a salt reacts with water to form an acid and a base.
  4. Fermentation: A chemical conversion producing alcohol caused by enzymes secreted by bacteria or other microorganisms acting on organic substrates.

Laws of Chemical Change

• Overview: The laws of chemical change represent fundamental scientific generalizations that govern all substances undergoing chemical transformation.

• The Three Fundamental Laws:

  • A. Law of Conservation of Mass (Lomonosov-Lavoisier Law):
  • Statement: In a closed system, the total mass of substances present before a chemical reaction occurs is strictly equal to the total mass of the products formed after the reaction.
  • Implication: Mass is neither created nor destroyed during any chemical process; it is conserved.
  • B. Law of Definite Composition (Law of Constant Composition):
  • First Observed By: French chemist Joseph Louis Proust.
  • Statement: Elements combine to form chemical compounds in exact, constant proportions by mass, regardless of the compound's source or method of preparation.
  • C. Law of Multiple Proportions (Dalton's Law):
  • Named After: English chemist Charles Dalton.
  • Statement: When two elements combine to form more than one distinct compound, the ratios of the masses of one element that combine with a fixed mass of the second element can always be reduced to small whole numbers.

Types of Chemical Reactions

• Overview: Chemists classify chemical reactions into specific structural categories to systematically predict products.

• 1. Combination or Synthesis Reaction:

  • Definition: A reaction where two or more substances (elements or compounds) combine to yield a single product.
  • General Equation:     $A + B \rightarrow AB$
  • Five Key Examples:
  1. Formation of Water:        $2\text{H}_{2(g)} + \text{O}_{2(g)} \rightarrow 2\text{H}_2\text{O}_{(l)}$
  2. Formation of Sodium Chloride:        $2\text{Na}_{(s)} + \text{Cl}_{2(g)} \rightarrow 2\text{NaCl}_{(s)}$
  3. Formation of Magnesium Oxide:        $2\text{Mg}_{(s)} + \text{O}_{2(g)} \rightarrow 2\text{MgO}_{(s)}$
  4. Formation of Ammonia:        $3\text{H}_{2(g)} + \text{N}_{2(g)} \rightarrow 2\text{NH}_{3(g)}$
  5. Formation of Carbon Dioxide:        $\text{C}_{(s)} + \text{O}_{2(g)} \rightarrow \text{CO}_{2(g)}$

• 2. Analysis or Decomposition Reaction:

  • Definition: A reaction where a single reactant compound breaks down into two or more smaller compounds or constituent elements. It represents the exact reverse of a synthesis reaction.
  • General Equation:     $AB \rightarrow A + B$
  • Five Key Examples:
  1. Thermal Decomposition of Calcium Carbonate:        $\text{CaCO}_{3(s)} \rightarrow \text{CaO}_{(s)} + \text{CO}_{2(g)}$
  2. Electrolysis of Water:        $2\text{H}_2\text{O}_{(l)} \rightarrow 2\text{H}_{2(g)} + \text{O}_{2(g)}$
  3. Decomposition of Hydrogen Peroxide (catalyzed by manganese dioxide, $\text{MnO}_2$):        $2\text{H}_2\text{O}_{2(aq)} \rightarrow 2\text{H}_2\text{O}_{(l)} + \text{O}_{2(g)}$
  4. Decomposition of Ammonium Nitrate:        $\text{NH}_4\text{NO}_3 \rightarrow \text{N}_{2(g)} + 2\text{H}_2\text{O}_{(g)} + \text{O}_{2(g)}$
  5. Decomposition of Hydrogen Chloride Gas:        $2\text{HCl}_{(g)} \rightarrow \text{H}_{2(g)} + \text{Cl}_{2(g)}$

• 3. Single-Replacement Reaction (Substitution Reaction):

  • Definition: A reaction in which a free element displaces another element that is bound within a compound. Reactants always consist of one free element and one compound.

  • General Equation:     $A + BX \rightarrow AX + B$(Where $\text{BX}$ and $\text{AX}$ are ionic compounds, and $\text{A}$ and $\text{B}$ are elements in their free uncombined states).

  • Activity Series Guidelines for Metals:

  • Order of Reactivity (from most to least reactive):       $\text{K} > \text{Na} > \text{Ca} > \text{Mg} > \text{Al} > \text{Zn} > \text{Fe} > \text{Pb} > \text{Hg} > \text{Ag} > \text{Pt} > \text{Au}$

  • Rules:

    • Any free metal situated higher on the activity series displaces a metal lower on the series from its aqueous compound.
    • Any metal positioned above hydrogen ($\text{H}$) on the activity series displaces hydrogen gas ($\text{H}_2$) from an acid.
    • Highly active metals (such as sodium, $\text{Na}$) displace hydrogen gas from water ($\text{H}_2\text{O}$), releasing $\text{H}_2$ and producing hydroxide ions:         $2\text{Na} + 2\text{H}_2\text{O} \rightarrow \text{H}_2 + 2\text{NaOH}$

  • Activity Series Guidelines for Nonmetals (Halogens):

  • Halogen Series Order: $\text{F}_2 > \text{Cl}_2 > \text{Br}_2 > \text{I}_2$

  • Rules: Nonmetals in free states exist as diatomic molecules. A free halogen displaces a less active halogen situated below it in the series from a compound.

  • Examples:

    • Free chlorine displaces bromine:         $\text{Cl}_2 + 2\text{NaBr} \rightarrow 2\text{NaCl} + \text{Br}_2$
    • Iodine cannot displace bromine:         $\text{I}_2 + 2\text{NaBr} \rightarrow \text{no reaction}$

  • Specific Metal Replacement Reaction Examples:

  1. Iron with Copper Sulfate:        $\text{Fe} + \text{CuSO}_4 \rightarrow \text{FeSO}_4 + \text{Cu}$(Iron is higher in activity than copper, successfully displacing Cu to form iron (II) sulfate).
  2. Copper with Iron Sulfate:        $\text{Cu} + \text{FeSO}_4 \rightarrow \text{no reaction}$(Copper is lower in activity than iron, yielding no reaction).
  3. Zinc with Hydrochloric Acid:        $\text{Zn} + 2\text{HCl} \rightarrow \text{ZnCl}_2 + \text{H}_2$(Zinc is situated above hydrogen in the activity series, displacing hydrogen gas).