Comprehensive Study Notes on Toxic Changes and Chemical Reactions

Unit Introduction: Toxins and Chemical Reactions

• Toxins Unit Overview: The focus of the Toxins Unit is the study of chemical reactions.

• Medicinal Uses of Toxins: Many toxins have medicinal applications.

  • Example: Compounds found in the venom of the African saw-scaled viper are used to treat individuals experiencing heart attacks.

• Living by Chemistry (LBC) Exercises: Worksheets and check-in sheets are provided as packets. These are to be uploaded to Google Classroom individually. Students must include their name, class period, and chapter number ($13$) at the top of each exercise page.

Representing Chemical Reactions and Equations

• Chemical Equation: Defined as a chemical "sentence" describing a reaction using numbers, symbols, and chemical formulas.

• Reaction Arrow ($\rightarrow$): The arrow in a chemical equation means "yields" and indicates the direction of the reaction.

• Reactant: A starting substance in a chemical reaction.

• Product: An ending substance in a chemical reaction.

• States of Matter Abbreviations:

  • $(g)$: Gas

  • $(l)$: Liquid

  • $(s)$: Solid

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

• Coefficients: A number placed in front of a chemical symbol or formula specifying how many atoms or molecules of that substance are involved.

  • Example in $H_2 + Cl_2 \rightarrow 2HCl$: The coefficient for $HCl$ is $2$.

• Subscripts: Numbers that indicate how many atoms of a specific element are present in a single molecule. Subscripts only refer to the atom immediately preceding them.

  • Example: In $H_2$, there are $2$ atoms of $H$ per molecule.

  • In $2HCl$, there is $1$ atom of $H$ and $1$ atom of $Cl$ per molecule.

Chemical Vocabulary and Diatomic Elements

• Molecular Formula: Represents the exact number of each type of atom in a single molecule.

  • Example: Water ($H_2O$) contains $2$ hydrogen atoms and $1$ oxygen atom for a total of $3$ atoms.

• Formula Unit: The simplest chemical formula for network covalent or ionic compounds, representing the smallest whole-number ratio of elements. Examples include $NaCl$ and $Fe_2O_3$.

• Diatomic Gases (BrINClHOF): These elements naturally exist as diatomic molecules in gas form:

  • $Br_2$

  • $I_2$

  • $N_2$

  • $Cl_2$

  • $H_2$

  • $O_2$

  • $F_2$

Physical vs. Chemical Changes

• Physical Change: A change in matter where the substance changes form but not its chemical identity.

  • Examples:

    • Phase changes (e.g., melting sugar: $C_{12}H_{22}O_{11}(s) \rightarrow C_{12}H_{22}O_{11}(l)$).

    • Dissolving (e.g., sugar in water: $C_{12}H_{22}O_{11}(s) \rightarrow C_{12}H_{22}O_{11}(aq)$).

    • Changes in shape.

  • Key Principle: If a substance changes phase or dissolves, its chemical formula does not change.

• Chemical Change (Chemical Reaction): A change in matter resulting in the formation of one or more new substances with new properties.

  • Example: Sugar decomposing: $C_{12}H_{22}O_{11}(s) \rightarrow 12C(s) + 11H_2O(g)$.

The Law of Conservation of Mass and Balancing Equations

• Law of Conservation of Mass: Matter is neither created nor destroyed in physical or chemical changes.

• Equation Balancing Principle: Equations must have an equal number of each type of atom on both sides (reactants and products).

• Balancing Rules:

  • Only coefficients can be changed.

  • Subscripts MUST NOT be changed, as this alters the identity of the chemical substance.

• Problem-Solving Method:

  1. Create an element inventory for both sides.

  2. Adjust coefficients and update the inventory.

  3. Repeat until the inventory for reactants equals the inventory for products.

  4. Coefficients act as counting units. One can double or halve a balanced equation (like a recipe) as long as ratios are maintained.

• Example Reaction Balancing:

  • Unbalanced: $NaOH + H_2SO_4 \rightarrow Na_2SO_4 + H_2O$

  • Step 1 (Inventory): Left: $Na=1$, $O=5$, $H=3$, $S=1$; Right: $Na=2$, $O=5$, $H=2$, $S=1$.

  • Step 2 (Adjust coefficients): Change reactant coefficient to $2NaOH$.

  • Step 3 (Final Balance): $2NaOH + H_2SO_4 \rightarrow Na_2SO_4 + 2H_2O$.

Types of Chemical Reactions

• Combination (Synthesis): Several reactants combine to form a single product.

  • General form: $A + B \rightarrow AB$

• Decomposition: One reactant breaks down into two or more products.

  • General form: $AB \rightarrow A + B$

• Single Exchange (Single Replacement/Displacement): One element reacts with a compound, and a more-reactive element replaces a less-reactive element.

  • General form: $A + BC \rightarrow AC + B$

• Double Exchange (Double Replacement/Displacement): Two compounds exchange atoms (cations or anions) with each other.

  • General form: $AB + CD \rightarrow AD + CB$

• Combustion: Rapid reaction with oxygen ($O_2$), releasing energy as heat and light.

  • General form: $X + O_2(g) \rightarrow \text{product(s)} + \text{heat} + \text{light}$

Predicting Products: Single Exchange and the Activity Series

• Requirement: Predicting a reaction requires understanding element reactivity and product solubility.

• Activity Series: A hierarchical list of elements ordered by reactivity.

  • Elements higher on the list are more reactive.

  • Metals: A metal will only replace a cation in a compound if it is higher (more reactive) on the activity series.

  • Nonmetals: A nonmetal will only replace an anion in a compound if it is higher (more reactive) on the activity series.

• Examples:

  • $Cu(s) + 2AgNO_3(aq) \rightarrow Cu(NO_3)_2(aq) + 2Ag(s)$ (Occurs because $Cu$ is more active than $Ag$).

  • $Cu(s) + AlCl_3(aq) \rightarrow \text{No Reaction}$ (Occurs because $Cu$ is less active than $Al$).

Solutions and Double Exchange Reactions

• Vocabulary:

  • Solution: A uniform mixture of two or more substances.

  • Solvent: The substance (often liquid, especially water) in which the solute dissolves.

  • Solute: The substance that is dissolved.

  • Solubility: The degree to which a substance dissolves.

  • Precipitate: A solid produced during a chemical reaction between two solutions.

• Reaction Conditions: For a double exchange reaction to occur, one product is typically a solid (precipitate), a gas, or a molecular compound like water.

• Ionic Dissociation: When ionic compounds dissolve in water, they separate into individual ions due to the partial charges on water molecules pulling the crystal lattice apart ($NaCl(s) \rightarrow Na^+(aq) + Cl^-(aq)$).

• Polyatomic Ions in Solution: These ions stay together because they are held by strong covalent bonds that water molecules cannot pull apart ($HNO_3(s) \rightarrow H^+(aq) + {NO_3}^-(aq)$).

Solubility Rules

• Soluble Compounds:

  1. Salts of ammonium (${NH_4}^+$) and Group IA metals (Alkali metals: $Li^+$, $Na^+$, $K^+$, etc.).

  2. All chlorates (${ClO_3}^-$), nitrates (${NO_3}^-$), and acetates (${CH_3COO}^-$).

  3. Chlorides ($Cl^-$), Bromides ($Br^-$), and Iodides ($I^-$) EXCEPT when paired with $Ag^+$, ${Hg_2}^{2+}$, or $Pb^{2+}$.

  4. Sulfates (${SO_4}^{2-}$) EXCEPT when paired with $Ca^{2+}$, $Sr^{2+}$, $Ba^{2+}$, ${Hg_2}^{2+}$, $Hg^{2+}$, $Pb^{2+}$, or $Ag_2^{2+}$.

• Insoluble Compounds:

  1. Phosphates (${PO_4}^{3-}$) and Carbonates (${CO_3}^{2-}$) EXCEPT with ${NH_4}^+$ and Group IA.

  2. Metallic oxides ($O^{2-}$) EXCEPT with ${NH_4}^+$ and Group IA.

  3. Metallic hydroxides ($OH^-$) EXCEPT with ${NH_4}^+$ and Group IA, or Group IIA from Calcium ($Ca$) down.

  4. Sulfides ($S^{2-}$) EXCEPT with ${NH_4}^+$ and Groups IA and IIA.

• Transition Metal Note: $Ag^+$, $Cd^{2+}$, and $Zn^{2+}$ have only one cation and do not require Roman numerals in naming.

Net Ionic Equations and Three-Beaker Method

• Complete Ionic Equation: Shows all ions dissolved in a solution.

  • Example: $Pb^{2+}(aq) + 2{NO_3}^-(aq) + 2Li^+(aq) + 2Cl^-(aq) \rightarrow PbCl_2(s) + 2Li^+(aq) + 2{NO_3}^-(aq)$ .

• Spectator Ion: An ion that appears on both sides of a complete ionic equation and does not directly participate in the chemical reaction. These ions are crossed out.

• Net Ionic Equation: The equation remaining after spectator ions are removed, showing only the particles involved in the reaction.

  • Example: $Pb^{2+}(aq) + 2Cl^-(aq) \rightarrow PbCl_2(s)$.

• Three-Beaker Method:

  • Beakers $1$ and $2$ contain the reactants (drawn as separate ions if soluble).

  • Beaker $3$ shows the products. Precipitates are drawn staying together at the bottom; spectator ions remain separate in solution.

Laboratory Activities and Demos

• Observing a Chemical Reaction Lab:

  • Substances: Copper (II) chloride ($CuCl_2$).

  • Safety: Wear goggles and gloves; $CuCl_2$ causes severe skin and eye damage.

  • Observations: Detailed, objective, and multiple per substance.

• Colored Precipitates Lab:

  • Focus: Analyzing solid products formed from various ionic solutions.

  • Requirement: Complete chart and data analysis.

• Snails and Fire Demos:

  • Video: "Beyond the Elements: Killer Snails" and "Beyond the Elements: Reactions" (covering fire and concrete).

Questions & Discussion

• Question: Does the equation $CuCl_2(aq) + Na_2S(aq) \rightarrow CuS(s) + NaCl(aq)$ obey the Law of Conservation of Mass?

  • Answer: No. There is an extra $Cl$ and an extra $Na$ on the reactant side. The balanced version must be $CuCl_2(aq) + Na_2S(aq) \rightarrow CuS(s) + 2NaCl(aq)$.

• Question: What are the names of these polyatomic ions: $OH^-$, ${NO_3}^-$, ${CO_3}^{2-}$, ${SO_4}^{2-}$, ${PO_4}^{3-}$, ${BrO_3}^-$, ${NH_4}^+$?

  • Answer: Hydroxide, Nitrate, Carbonate, Sulfate, Phosphate, Bromate, Ammonium.

• Question: Write a chemical equation for solid magnesium sulfide heated to produce solid magnesium and sulfur gas.

  • Answer: $MgS(s) \rightarrow Mg(s) + S(g)$.

• Question: Predict the solubility of $PbI_2$.

  • Answer: $PbI_2(s)$ (Insoluble).