chapter 7: chemical reactions & chemical quantities

What is the greenhouse effect, and how does it affect Earth's temperature?

Greenhouse gases allow sunlight to enter the atmosphere but prevent some heat from escaping. This keeps Earth warmer than it would otherwise be. Increasing greenhouse gas concentrations can raise Earth's average temperature.

Why are scientists concerned about increasing atmospheric CO₂?

CO₂ is a significant greenhouse gas. Burning fossil fuels increases atmospheric CO₂, strengthening the atmosphere's ability to retain heat and contributing to global warming.

What is a chemical change?

A change in which atoms rearrange and form different substances. Examples include burning gasoline, rusting iron, burning sugar, and photosynthesis.

SIMPLE: THIS IS WHEN THE MOLECULAR FORMULA CHANGES.

eg) rust on a nail

What is a physical change?

A change in which a substance's form or state changes but its chemical identity remains the same. Examples include boiling water, evaporating alcohol, cutting a material, and crushing a solid.

SIMPLE: CHANGE IN STATE OF MATTER (SOLID, LIQUID, GAS)

How can you distinguish a physical change from a chemical change using molecules?

In a physical change, the molecules retain their identities even if their arrangement or spacing changes. In a chemical change, atoms rearrange to form different molecules or substances.

What is a chemical property versus a physical property?

A chemical property describes how a substance behaves when its composition changes, such as flammability or corrosiveness. A physical property can be observed without changing composition, such as color, odor, density, or boiling point.

Classify these changes: natural gas burns; propane evaporates; propane burns; iron rusts.

Natural gas burning: chemical. Propane evaporating: physical. Propane burning: chemical. Iron rusting: chemical. Burning and rusting form new substances; evaporation does not.

Classify these changes: sugar burns; sugar dissolves in water; a platinum ring becomes dull from abrasion; silver tarnishes.

Sugar burning: chemical. Sugar dissolving: physical. Platinum becoming dull from abrasion: physical. Silver tarnishing: chemical.

Classify isopropyl alcohol properties: colorless, flammable, liquid at room temperature, density 0.79 g/mL, mixes with water.

Colorless: physical. Flammable: chemical. Liquid at room temperature: physical. Density: physical. Mixing with water: physical.

Classify ozone properties: bluish color, pungent odor, very reactive, decomposes under ultraviolet light, gas at room temperature.

Color and odor: physical. Being very reactive and decomposing under ultraviolet light: chemical. Being a gas at room temperature: physical.

Which self-assessment change is physical: wood burning, iron rusting, dynamite exploding, or gasoline evaporating?

Gasoline evaporating. It changes from liquid to gas without changing its chemical identity. The other processes involve chemical reactions.

Which property of rubbing alcohol is chemical: density, flammability, boiling point, or melting point?

Flammability. It describes the ability to undergo combustion and form new substances. Density, boiling point, and melting point are physical properties.

What is a chemical equation? What is always in a chemical equation?

A symbolic representation of a chemical reaction showing the reactants, products, and their relative amounts.

1) Molecular Formula/Structures

2) State of Matter (s, l, g, aq)

3) Stoichiometric Coefficients (Smallest whole number integer)

4) Reaction or Equilibrium Arrows

What are reactants and products?

Reactants are the starting substances, written to the left of the arrow. Products are the substances formed, written to the right of the arrow.

What do the state symbols (s), (l), (g), and (aq) mean?

(s) = solid; (l) = liquid; (g) = gas; (aq) = solid or gas COMPLETELY dissolved in water; aqueous.

What is the law of conservation of mass, and how does it apply to chemical equations?

Matter is conserved during a chemical reaction. Atoms are rearranged, not created or destroyed, so each element must have the same number of atoms on both sides of a balanced equation.

“Cannot be created or destroyed”

What is a coefficient in a chemical equation?

A number placed before a formula that specifies how many molecules or formula units participate. MUST BE LOWEST WHOLE NUMBER INTEGER POSSIBLE.

For example, 3H₂O represents three water molecules containing six H atoms and three O atoms.

What is a subscript in a chemical formula?

A small number within a formula showing how many atoms of an element occur in each molecule or formula unit. In CO₂, each molecule contains one C atom and two O atoms.

Why can you change coefficients but not subscripts when balancing equations?

Coefficients change the number of particles without changing their identities. Subscripts change the chemical formula itself; H₂O is water, whereas H₂O₂ is hydrogen peroxide.

How do you balance a chemical equation?

Write correct formulas first. Count each element on both sides, adjust coefficients to equalize atom counts, recount, and reduce coefficients to the smallest whole-number ratio. Never change subscripts.

How do you balance CH₄ + O₂ → CO₂ + H₂O?

CH₄ + 2O₂ → CO₂ + 2H₂O. Four H atoms require 2H₂O. The products then contain four O atoms, requiring 2O₂ on the reactant side.

Balance PCl₃ + H₂O → H₃PO₃ + HCl.

PCl₃ + 3H₂O → H₃PO₃ + 3HCl. Coefficients from left to right: 1, 3, 1, 3.

Balance SO₂ + O₂ → SO₃.

2SO₂ + O₂ → 2SO₃. Both sides contain two S atoms and six O atoms.

Balance Fe + O₂ → Fe₂O₃.

4Fe + 3O₂ → 2Fe₂O₃. This gives four Fe atoms and six O atoms on each side.

Balance Na + H₂O → NaOH + H₂.

2Na + 2H₂O → 2NaOH + H₂. The coefficients balance Na, O, and H atoms.

Balance Ca + H₂O → Ca(OH)₂ + H₂.

Ca + 2H₂O → Ca(OH)₂ + H₂. Both sides contain one Ca, two O, and four H atoms.

Balance N₂ + H₂ → NH₃.

N₂ + 3H₂ → 2NH₃. Two N atoms require 2NH₃, which contains six H atoms, requiring 3H₂.

Balance C₃H₈ + O₂ → CO₂ + H₂O.

C₃H₈ + 5O₂ → 3CO₂ + 4H₂O. Balance C first, then H, then O. The products contain ten O atoms, requiring 5O₂.

How do you balance a combustion equation for a compound containing C and H, or C, H, and O?

For complete combustion, place CO₂ among the products for carbon and H₂O for hydrogen. Balance C, then H, then O. If the fuel contains O, count its oxygen atoms before determining the O₂ coefficient.

What is the balanced equation for complete combustion of ethanol, C₂H₅OH?

C₂H₅OH + 3O₂ → 2CO₂ + 3H₂O. Two C atoms give 2CO₂; six H atoms give 3H₂O; the products contain seven O atoms, one already in ethanol.

What is reaction stoichiometry?

The quantitative relationship between reactants and products in a chemical reaction. A balanced equation lets you calculate how much reactant is needed or how much product can form.

What do coefficients in a balanced equation tell you about moles?

They give mole ratios. For 2H₂ + O₂ → 2H₂O, 2 mol H₂ react with 1 mol O₂ to form 2 mol H₂O.

What is the general setup for a stoichiometry calculation?

Convert the given quantity to moles if needed; multiply by the balanced-equation mole ratio; convert the resulting moles to the requested unit if needed. The mole ratio comes from the balanced equation.

Why must a chemical equation be balanced before stoichiometry calculations?

The coefficients provide the correct quantitative mole ratios. An unbalanced equation gives incorrect ratios and therefore incorrect calculated quantities.

How many moles of C react with 6 mol Co₂O₃ in 2Co₂O₃ + 3C → 4Co + 3CO₂?

6 mol Co₂O₃ × (3 mol C / 2 mol Co₂O₃) = 9 mol C. The equation requires 3 mol C for every 2 mol Co₂O₃.

In 2C₈H₁₈ + 25O₂ → 16CO₂ + 18H₂O, how many moles of CO₂ form per mole of octane?

16 mol CO₂ / 2 mol C₈H₁₈ = 8 mol CO₂ per mole of octane. The balanced coefficients give the ratio.

How many moles of O₂ are needed to react with 2.50 mol CH₄?

CH₄ + 2O₂ → CO₂ + 2H₂O. 2.50 mol CH₄ × (2 mol O₂ / 1 mol CH₄) = 5.00 mol O₂.

How many moles of H₂O form from 3.00 mol O₂ in 2H₂ + O₂ → 2H₂O, assuming sufficient H₂?

3.00 mol O₂ × (2 mol H₂O / 1 mol O₂) = 6.00 mol H₂O.

How do you convert a reactant mass into a product mass?

Convert grams of reactant to moles using its molar mass, multiply by the product-to-reactant mole ratio, then convert product moles to grams using the product's molar mass.

Self-assessment 5: How much Al reacts with 25.0 g MnO₂? 3MnO₂ + 4Al → 3Mn + 2Al₂O₃.

Molar mass MnO₂ ≈ 86.94 g/mol. 25.0 g ÷ 86.94 = 0.287 mol MnO₂. Multiply by 4 mol Al / 3 mol MnO₂ and by 26.98 g/mol: about 10.3 g Al.

What is a limiting reactant?

The reactant that runs out first according to the balanced equation. It determines the maximum amount of product that can form.

What is an excess reactant?

A reactant supplied in more than the stoichiometric amount needed to react completely with the limiting reactant. Some of it remains after the reaction.

What is theoretical yield?

The maximum amount of product predicted from the limiting reactant, assuming the reaction proceeds as expected and the product is recovered completely.

What is actual yield?

The amount of product actually obtained from a reaction in practice. It is usually less than theoretical yield because of incomplete reaction, side reactions, or product loss.

What is percent yield, and what is its formula?

Percent yield = (actual yield ÷ theoretical yield) × 100%. It expresses what percentage of the predicted product was actually obtained.

Can percent yield be calculated using moles instead of grams?

Yes, provided actual and theoretical yields use the same units and refer to the same product. Their ratio is unchanged when both quantities are converted consistently.

How do you identify the limiting reactant when given masses?

Convert each reactant mass to moles. Calculate the amount of the same product each reactant could form using the balanced equation. The reactant that produces the smaller amount of product is limiting.

Is the limiting reactant always the reactant with the smallest mass?

No. The limiting reactant depends on the available moles and the equation's mole ratios, not on which reactant has the smallest mass.

How do you calculate the amount of excess reactant remaining?

Identify the limiting reactant. Use its amount and the balanced equation to calculate how much excess reactant is consumed. Subtract the consumed amount from the initial amount of excess reactant.

Self-assessment 4: For 3A + 2B + C → 2D, 3.5 mol A, 5.9 mol B, and 2.2 mol C are mixed. What is limiting?

A is limiting. A could form (3.5 × 2/3) = 2.33 mol D; B could form 5.9 mol D; C could form 4.4 mol D. A forms the least product.

For 2Mg + O₂ → 2MgO, how do you determine the limiting reactant from 42.5 g Mg and 33.8 g O₂?

Convert each mass to moles and calculate possible MgO. Mg: 42.5/24.31 ≈ 1.75 mol, producing 1.75 mol MgO. O₂: 33.8/32.00 ≈ 1.06 mol, producing 2.11 mol MgO. Mg is limiting.

For the magnesium reaction above, what is the theoretical yield of MgO?

Mg is limiting and produces about 1.75 mol MgO. Multiply by 40.31 g/mol: theoretical yield ≈ 70.5 g MgO.

If 55.9 g MgO is obtained from a theoretical yield of 70.5 g, what is the percent yield?

(55.9 g ÷ 70.5 g) × 100% = 79.3%.

Self-assessment 6: What is the theoretical yield of NaCl from 55.0 g Na and 67.2 g Cl₂? 2Na + Cl₂ → 2NaCl.

Na: 55.0/22.99 ≈ 2.39 mol; Cl₂: 67.2/70.90 ≈ 0.948 mol. Cl₂ is limiting and forms 1.896 mol NaCl. Multiply by 58.44 g/mol: 111 g NaCl.

Self-assessment 7: 50.0 g S reacts with 105.0 g F₂ to form SF₆. How much excess reactant remains? S + 3F₂ → SF₆.

F₂: 105.0/38.00 ≈ 2.76 mol, so it consumes 2.76/3 ≈ 0.921 mol S, or about 29.5 g S. Sulfur is in excess; 50.0 − 29.5 = 20.5 g S remains.

Self-assessment 8: The theoretical yield is 45.8 g and actual yield is 37.2 g. What is the percent yield?

(37.2 ÷ 45.8) × 100% = 81.2%.

What does a percent yield below 100% usually indicate?

Less product was collected than theoretically possible. Product may be lost, the reaction may be incomplete, or competing reactions may occur.

What does a calculated percent yield above 100% suggest?

Check for experimental error, impurities, leftover solvent, or an incorrect theoretical yield. The actual amount of pure product should not exceed the theoretical maximum under the stated assumptions.

How does the pizza-recipe analogy explain limiting reactants?

The ingredient that makes the fewest complete pizzas limits production. Similarly, the reactant that can produce the least product determines the theoretical yield.

In N₂ + 3H₂ → 2NH₃, how do you determine which reactant limits production?

Calculate how much NH₃ each starting reactant could form using the equation's ratios. The reactant yielding less NH₃ is limiting; the other reactant is in excess.

Self-assessment 1: Balance PCl₃ + H₂O → H₃PO₃ + HCl. What are the coefficients?

1, 3, 1, 3. Balanced equation: PCl₃ + 3H₂O → H₃PO₃ + 3HCl.

Self-assessment 9: Which is the balanced complete combustion equation for propane, C₃H₈?

C₃H₈ + 5O₂ → 3CO₂ + 4H₂O. It conserves 3 C atoms, 8 H atoms, and 10 O atoms.

Self-assessment 10: How many moles O₂ form when 55.8 g KClO₃ decomposes?

2KClO₃ → 2KCl + 3O₂. Molar mass KClO₃ ≈ 122.55 g/mol; 55.8/122.55 = 0.455 mol KClO₃. Multiply by 3/2: 0.683 mol O₂.

What is a combustion reaction?

A reaction in which a substance reacts with O₂ to form one or more oxygen-containing compounds, often releasing heat. Complete combustion of a carbon- and hydrogen-containing fuel forms CO₂ and H₂O.

What products form during complete combustion of a compound containing carbon and hydrogen?

Carbon forms CO₂ and hydrogen forms H₂O. For example, propane burns in oxygen to form carbon dioxide and water.

What are the products when elemental carbon burns completely in oxygen?

Carbon dioxide: C(s) + O₂(g) → CO₂(g).

What are the products when hydrogen gas burns in oxygen?

Water: 2H₂(g) + O₂(g) → 2H₂O(g).

What general reaction pattern describes an alkali metal reacting with a halogen?

2M + X₂ → 2MX. M represents an alkali metal and X represents a halogen. Example: 2Na + Cl₂ → 2NaCl.

Why are alkali metals highly reactive?

They have one valence electron that is relatively easy to remove, giving them low first ionization energies. They readily form positive ions.

What happens when an alkali metal reacts with water?

It forms aqueous metal ions, hydroxide ions, and hydrogen gas: 2M + 2H₂O → 2M⁺(aq) + 2OH⁻(aq) + H₂(g).

Why can alkali metal reactions with water be dangerous?

The reaction releases substantial heat and produces hydrogen gas, which can ignite. The reaction becomes more explosive down the alkali-metal group.

Why are halogens highly reactive nonmetals?

They have seven valence electrons and need one more to reach a noble-gas configuration. Their strong tendency to gain an electron makes them reactive.

What is the general equation for a halogen reacting with a metal?

2M + nX₂ → 2MXₙ, where M is the metal, X is the halogen, and n reflects the metal's charge in the product. Example: 2Fe + 3Cl₂ → 2FeCl₃.

What is the general equation for a halogen reacting with hydrogen?

H₂ + X₂ → 2HX. The hydrogen halides formed contain covalent bonds because hydrogen and the halogen are nonmetals.

What is an interhalogen compound, and how does it form?

A compound containing two different halogens. Example: Br₂ + F₂ → 2BrF. Because both elements are nonmetals, the compound contains covalent bonds.

What type of bonding occurs in metal halides versus hydrogen halides or interhalogens?

Metal halides contain ionic bonds between a metal and a nonmetal. Hydrogen halides and interhalogen compounds contain covalent bonds between nonmetals.

Exercise 23: Write the balanced equation for SO₂ reacting with O₂ to form sulfuric acid, H₂SO₄.

As written, SO₂ + O₂ alone cannot form H₂SO₄ because hydrogen is absent. For acid rain formation, the full reaction needs a hydrogen-containing reactant; do not invent a balanced equation without it.

Exercise 25: Solid sodium reacts with water to form aqueous sodium hydroxide and hydrogen gas. Write the balanced equation.

2Na(s) + 2H₂O(l) → 2NaOH(aq) + H₂(g).

Exercise 26: Iron reacts with oxygen to form iron(III) oxide. Write the balanced equation.

4Fe(s) + 3O₂(g) → 2Fe₂O₃(s).

Exercise 27: Nitrogen triiodide decomposes into nitrogen gas and iodine. Write the balanced equation.

2NI₃(s) → N₂(g) + 3I₂(s). Balance N first, then I.

Exercise 30: Write the balanced equation for photosynthesis from carbon dioxide and water forming glucose and oxygen.

6CO₂ + 6H₂O → C₆H₁₂O₆ + 6O₂. This balances six C, twelve H, and eighteen O atoms on each side.

Exercise 39: What is the general method for solving a reaction stoichiometry problem when given an unbalanced equation?

First balance the equation. Convert the given quantity to moles, apply the coefficient ratio, then convert to the requested unit. Never use coefficients from an unbalanced equation.

Exercise 41–42: How do you find the moles of multiple reactants needed for a specified amount of a substance?

Balance the equation, convert the given amount to moles if needed, then multiply by each requested reactant-to-given-substance mole ratio. Cancel units to check the setup.

Exercise 43–44: A problem asks you to balance an equation and calculate moles of product. What is the correct sequence?

Balance the equation first, identify the given and desired substances, use the coefficient ratio to calculate product moles, and check that the units and atom balance are correct.

Exercise 47: Iron reacts with hydrobromic acid. How do you calculate the HBr needed and H₂ formed from 2.80 g Fe?

Use the balanced equation Fe + 2HBr → FeBr₂ + H₂. Convert 2.80 g Fe to moles, use 2 mol HBr per mol Fe and 1 mol H₂ per mol Fe, then convert each result to grams if requested.

Exercise 48: Aluminum reacts with sulfuric acid. How do you calculate acid needed and hydrogen produced?

Balance 2Al + 3H₂SO₄ → Al₂(SO₄)₃ + 3H₂. Convert the given Al mass to moles, use the ratios 3 mol H₂SO₄/2 mol Al and 3 mol H₂/2 mol Al, and convert to grams as needed.

Exercise 49–50: How do you calculate the mass of a requested substance in a reaction?

Balance the equation, convert the given mass to moles, apply the mole ratio, and convert the answer to the desired substance's mass using its molar mass.

Exercise 51: How do you calculate the mass of water formed when a compound containing carbon, hydrogen, and oxygen burns?

Write and balance the complete-combustion equation. Convert the given compound mass to moles, use its coefficient ratio to find moles of H₂O, then multiply by 18.02 g/mol.

Exercise 52: How do you calculate product mass when nitrogen triiodide reacts?

Use the balanced equation 2NI₃ → N₂ + 3I₂. Convert the given NI₃ amount to moles, apply the appropriate coefficient ratio, and convert the product moles to mass if requested.

Exercises 53–54: How do you find the limiting reactant when the initial amounts are given?

Convert quantities to moles, calculate the possible amount of the same product from each reactant, and identify which produces the least. That reactant is limiting.

Exercises 55–56: How do you find how many molecules of a product form from given reactant molecules?

Use the balanced equation's coefficient ratio directly on the molecule counts. The ratio works for molecules just as it does for moles.

Exercises 57–58: How do you calculate theoretical yield in moles?

Identify the limiting reactant, then multiply its moles by the product-to-limiting-reactant coefficient ratio. The result is the theoretical yield in moles.

Exercises 59–60: How do you determine the limiting reactant when reactants are given in moles?

Calculate the possible product from each reactant using the balanced equation. For ZnS + O₂ → ZnO + SO₂, balance first: 2ZnS + 3O₂ → 2ZnO + 2SO₂. For FeS + HCl, balance first as well before comparing product amounts.

Exercises 61–62: How do you calculate theoretical yield when reactant quantities are provided?

Balance the equation, convert given quantities to moles, calculate the product possible from each reactant, select the smaller amount, and convert to requested units.

Exercises 63–65: What should you do when a reaction problem involves several reactants or ions?

Use the balanced equation and the stated quantities to calculate how much product each reactant can form. The smallest possible product amount determines the limiting reactant and theoretical yield.

Exercises 67–68: How do you solve a reaction problem that asks for a product mass from given reactant masses?

Balance the equation, convert each given mass to moles, compare possible product yields if there are multiple reactants, and use the limiting reactant to calculate the theoretical product mass.

Exercises 69–70: How do you complete and balance an unfamiliar combustion reaction?

For complete combustion, place CO₂ and H₂O among the products when the fuel contains carbon and hydrogen. Balance C, then H, then O, and use whole-number coefficients.

Exercises 71–76: How do you write equations for alkali metals and halogens?

Use the reaction patterns: 2M + X₂ → 2MX; 2M + 2H₂O → 2M⁺ + 2OH⁻ + H₂; 2M + nX₂ → 2MXₙ; and H₂ + X₂ → 2HX. Balance each equation.

Exercise 77: How do you calculate CO₂ released by gasoline combustion?

Represent gasoline with the stated fuel formula, balance its combustion equation, convert fuel mass to moles, use the CO₂-to-fuel mole ratio, and convert CO₂ moles to mass.

Exercise 78: What is the strategy for calculating CO₂ from propane combustion?

Use C₃H₈ + 5O₂ → 3CO₂ + 4H₂O. Convert propane mass to moles, multiply by 3 mol CO₂ per mol propane, then convert CO₂ moles to the requested mass units.