Chemical Reactions

Unit 5: Chemical Reactions

Concept 1: Writing Reactions

• Vocabulary:

  • Chemical reaction: Process by which substances collide with enough energy to form new bonds between atoms, thus creating new substances.

  • Reactants: Starting substances in a chemical reaction.

  • Products: Ending substances in a chemical reaction.

  • Aqueous: Dissolved in water.

• Objectives:

  • Differentiate between physical and chemical changes:

    • Physical changes: Do not impact a substance’s identity (e.g., change in state of matter).

    • Chemical changes: Result in a new substance from a chemical reaction occurring (e.g., burning).

  • Explain the collision theory:

    • Reacting particles must collide with enough energy for a chemical reaction to occur.

  • List several signs that can provide evidence that a chemical reaction has occurred:

    • Release of light or heat.

    • Sudden color change.

    • Odor change.

    • Gas released.

    • Formation of a precipitate.

  • Explain the Law of Conservation of Mass and how it relates to chemical reactions:

    • Matter cannot be created or destroyed in a system; it can only change forms.

    • In a chemical reaction, the mass of the reactants should equal the mass of the products.

  • Differentiate between coefficients and subscripts in a chemical reaction:

    • Coefficients: Placed in front of a substance and represent ratios of reactants to products; used for balancing reactions and can be changed.

    • Subscripts: Small numbers within each chemical formula that show the ratio of atoms in a compound; they cannot be changed because changing them changes the identity of the substance.

  • Objectives:

    • Label, interpret, and write equations for chemical reactions with correct notation.

    • Balance chemical reactions according to the Law of Conservation of Mass.

    • Write the balanced chemical reaction for a given written description of a reaction.

• Practice:

  • List the number of atoms of each element in the compounds below.

    • Mg3(PO4)2

      • Mg – 3, P – 2, O – 8

    • 3CO2

      • C – 3, O – 6

    • 2Ca(OH)2

      • Ca – 2, O – 4, H – 4

  • Balance the following reactions:

    • NF3 à N2 + F2

      • $2NF<em>3 à N</em>2 + 3F_2$

    • AlCl3 + K2SO4 à KCl + Al2(SO4)3

      • $2AlCl<em>3 + 3K</em>2SO<em>4 à 6KCl + Al</em>2(SO<em>4)</em>3$

    • C4H10 + O2 à CO2 + H2O

      • $2C<em>4H</em>{10} + 13O<em>2 à 8CO</em>2 + 10H_2O$

    • Pb(OH)2 + HCl à H2O + PbCl2

      • $Pb(OH)<em>2 + 2HCl à 2H</em>2O + PbCl_2$

  • Write the balanced chemical equation for the following reaction: Solid aluminum reacts with oxygen gas to create solid aluminum oxide.

    • $4Al(s) + 3O<em>2(g) à 2Al</em>2O_3(s)$

Concept 2: Classifying Reactions

• Vocabulary:

  • Oxide: A binary compound with at least 1 atom of oxygen

  • Salt: A compound made of cations and anions

  • Electrolysis: The decomposition of a substance by an electric current

  • Hydrocarbon: A compound made of carbon and hydrogen

  • Precipitate: An insoluble solid that forms from the ions of two aqueous compounds

  • Activity series: A list of elements in the order in which they will easily undergo certain chemical reactions

• Objectives:

  • Create a chart to summarize and organize the characteristics of the types of chemical reactions.

    • Synthesis:

      • General Equation: $A + B à AB$

      • Other notes: 2 or more reactants combine to form a new compound; React with O2 to form oxide; Metal and nonmetal oxides form salts.

    • Decomposition:

      • General Equation: $AB à A + B$

      • Other notes: 1 reactant breaks down into 2 or more products; Usually needs heat or electricity to happen

    • Combustion:

      • General Equation: reactant + $O_2$ à product

      • Other notes: When a substance reacts (burns) with $O<em>2$; When the reactant is a hydrocarbon, $CO</em>2$ and $H_2O$ are products.

    • Single Replacement:

      • General Equation: $A + BC à B + AC$

      • Other notes: 1 element replaces a like element in a compound; The most reactive metals can react with $H<em>2O$ to form metal hydroxide and $H</em>2$.

    • Double Replacement:

      • General Equation: $AB + CD à AD + CB$

      • Other notes: When ions in 2 compounds swap and make 2 new compounds; Often makes a precipitate, insoluble gas, or molecular compound.

  • Classify the 5 reactions from #10 and #11 in Concept 1.

    • #10: $NF<em>3 à N</em>2 + F_2$: Decomposition

    • $AlCl<em>3 + K</em>2SO<em>4 à KCl + Al</em>2(SO<em>4)</em>3$: Double

    • $C<em>4H</em>{10} + O<em>2 à CO</em>2 + H_2O$: Combustion

    • $Pb(OH)<em>2 + HCl à H</em>2O + PbCl_2$: Double (neutralization)

    • #11: $4Al(s) + 3O<em>2(g) à 2Al</em>2O_3(s)$: Synthesis

  • Be able to classify reactions based on an equation or description.

  • Be able to identify if a chemical reaction will occur using the activity series.

  • Be able to predict the products that will result from a chemical reaction when given the reactants.

• Practice:

  • Write out each chemical reaction described below. Then balance and classify:

    • Solid calcium and liquid water react to form aqueous calcium hydroxide and gaseous hydrogen.

      • $Ca(s) + 2H<em>2O(l) à Ca(OH)</em>2(aq) + H_2(g)$

      • Single

    • Copper (II) nitrate and magnesium react to form magnesium nitrate and copper.

      • $Cu(NO<em>3)</em>2 + Mg à Mg(NO<em>3)</em>2 + Cu$

      • Single

    • Propane gas ($C<em>3H</em>8$) is burned.

      • $C<em>3H</em>8(g) + 5O<em>2(g) à 3CO</em>2(g) + 4H_2O(g)$

      • Combustion

  • Predict the products for the chemical reactions below. Then, balance the equations, as needed:

    • Combustion: $C<em>6H</em>{12}O<em>6 + O</em>2 à$

      • $C<em>6H</em>{12}O<em>6 + 6O</em>2 à 6CO<em>2 + 6H</em>2O$

    • Synthesis: $K + Cl_2 à$

      • $2K + Cl_2 à 2KCl$

    • Decomposition: $H_2O à$

      • $2H<em>2O à 2H</em>2 + O_2$

    • Single: $LiF + Mg à$

      • No reaction. Magnesium will not replace lithium

    • Double: $Na<em>3PO</em>4 + CaCl_2 à$

      • $2Na<em>3PO</em>4 + 3CaCl<em>2 à Ca</em>3(PO<em>4)</em>2 + 6NaCl$

Concept 3: Chemical Equilibrium

• Vocabulary:

  • Chemical equilibrium: A dynamic process where there is no net change occurring in the amount of reactants and products in the system, thus no visible change.

  • Le Chatelier’s Principle: When conditions change for a system at equilibrium, the system responds by reducing the effect of the change.

  • Exothermic: When a reaction releases heat.

  • Endothermic: When a reaction absorbs heat.

• Objectives:

  • Explain what chemical equilibrium is and what it is not.

    • Equilibrium is when the forward and reverse reactions are occurring at the same rate.

    • It does NOT mean the amount of reactants and products are equal, or that the reaction has ended.

  • Describe an example of Le Chatelier’s principle at work.

    • $H<em>2CO</em>3(aq) <br>ightharpoons CO<em>2(aq) + H</em>2O(l)$

    • In the human body, as we intake more $CO<em>2$ when exercising, it causes the reaction to shift to make more $H</em>2CO_3$.

  • Describe the response to a change in concentration, temperature, and pressure for a chemical reaction in equilibrium.

    • Concentration:

      • Adding reactant or removing product favors the forward reaction.

      • Adding product or removing reactant favors the reverse reaction.

    • Temperature:

      • Adding heat to an endothermic reaction or removing heat from an exothermic reaction favors the forward reaction.

      • Adding heat to an exothermic reaction or removing heat from an endothermic reaction favors the reverse reaction.

    • Pressure (gases only!!):

      • Increasing pressure (or decreasing volume) causes a shift to favor whichever direction makes less gas.

      • Decreasing pressure (or increasing volume) causes a shift to favor whichever direction makes more gas.

  • Be able to predict the response to a specific change in condition on a chemical reaction in equilibrium when given a model, description, or example.

  • Be able to interpret diagrams and graphs that represent reactions in chemical equilibrium.

• Practice:

  • Consider the following reaction: $N<em>2O</em>4(g) + heat ightharpoons 2NO_2(g)$.

    • Is the reaction endothermic or exothermic? Explain how you know.

      • Endothermic, because heat is a reactant and endothermic means that, overall, the reaction absorbs heat.

    • Predict what would happen to the reaction once it has reached equilibrium if each of the following conditions is changed:

      • $N<em>2O</em>4$ is added

        • Shifts right (favors forward reaction)

      • $N<em>2O</em>4$ is removed

        • Shifts left (favors reverse reaction)

      • $NO_2$ is added

        • Shifts left (favors reverse reaction)

      • $NO_2$ is removed

        • Shifts right (favors forward reaction)

  • Consider the following reaction: $N<em>2O</em>4(g) + heat ightharpoons 2NO_2(g)$.

    • Predict what would happen to the reaction once it has reached equilibrium if each of the following conditions is changed:

      • Heat is added

        • Shifts right (favors forward reaction)

      • Temperature is decreased

        • Shifts left (favors reverse reaction)

      • Pressure is increased

        • Shifts left (favors reverse reaction)

      • Pressure is decreased

        • Shifts right (favors forward reaction)

      • Volume is decreased

        • Shifts left (favors reverse reaction)

      • Volume is increased

        • Shifts right (favors forward reaction)