Chapter 7 – Chemical Reactions and Quantities: Study Notes

Chapter 7 – Chemical Reactions and Quantities

Section 7.1: Equations for Chemical Reactions

Chemical reactions involve the transformation of substances into new products with different chemical formulas and properties. Understanding how to represent and balance these reactions is fundamental in chemistry.

• Chemical Change: A process in which a substance is converted into one or more new substances with different formulas and properties.

• Possible Observations of Chemical Change:

  • Formation of bubbles (gas production)

  • Color change

  • Production of a solid (precipitate)

  • Heat produced or absorbed

Example: The rusting of iron: Fe reacts with O2 to form rust (Fe2O3).

Symbols Used in Chemical Equations

• An arrow (→) separates reactants from products.

• Reactants are written on the left, products on the right.

• Multiple reactants or products are separated by a + sign.

• The delta symbol (Δ) above the arrow indicates heat is used to start the reaction.

• Physical states are shown in parentheses:

  • (s): solid

  • (l): liquid

  • (g): gas

  • (aq): aqueous (dissolved in water)

General format: reactant 1 + reactant 2 → product 1 + product 2

Identifying a Balanced Equation

• No atoms are lost or created (Law of Conservation of Matter).

• The number of atoms of each element is the same on both sides of the equation.

Example: $\mathrm{C(s) + O_2(g) \xrightarrow{\Delta} CO_2(g)}$

Reactant atoms = Product atoms

Guide to Balancing a Chemical Equation

1. Write the equation using correct formulas for reactants and products.

2. Count the atoms of each element on both sides.

3. Use coefficients to balance each element (never change subscripts).

4. Check that coefficients are the lowest whole numbers and the equation is balanced.

Example: Formation of aluminum sulfide from aluminum and sulfur: $\mathrm{2Al(s) + 3S(s) \rightarrow Al_2S_3(s)}$

Practice Problems

• Balance: $\mathrm{P_4(s) + Br_2(l) \rightarrow PBr_3(g)}$

• Balance: $\mathrm{Fe_2O_3(s) + H_2(g) \rightarrow Fe(s) + H_2O(l)}$

Section 7.4: The Mole

The mole is a counting unit in chemistry, similar to terms like dozen or gross, but much larger. It allows chemists to count atoms, molecules, or ions by weighing them.

• 1 mole = $6.02 \times 10^{23}$ items (Avogadro's number)

• Used for atoms, molecules, ions, etc.

Examples:

• 1 mole of Na contains $6.02 \times 10^{23}$ atoms of Na

• 1 mole of H2O contains $6.02 \times 10^{23}$ molecules of water

Avogadro's Number and Conversion Factors

• Equality: $1\ \text{mole} = 6.02 \times 10^{23}\ \text{particles}$

• Conversion Factors:

  • $\frac{1\ \text{mole}}{6.02 \times 10^{23}\ \text{particles}}$

  • $\frac{6.02 \times 10^{23}\ \text{particles}}{1\ \text{mole}}$

Guide to Calculating Atoms or Molecules

1. State the given and needed quantities.

2. Use Avogadro's number to write conversion factors.

3. Set up the problem to cancel units and calculate the number of particles.

Example: How many CO2 molecules are in 0.500 mole of CO2?

Moles of Elements in a Formula

Subscripts in a chemical formula indicate the number of atoms of each element in one molecule, and the number of moles of each element in one mole of the compound.

• Example: Aspirin, C9H8O4

  • 1 molecule: 9 C, 8 H, 4 O atoms

  • 1 mole: 9 moles C, 8 moles H, 4 moles O

Possible conversion factors for 1 mole of C9H8O4:

• $\frac{9\ \text{moles C}}{1\ \text{mole C}_9\text{H}_8\text{O}_4}$

• $\frac{8\ \text{moles H}}{1\ \text{mole C}_9\text{H}_8\text{O}_4}$

• $\frac{4\ \text{moles O}}{1\ \text{mole C}_9\text{H}_8\text{O}_4}$

Guide to Calculating Moles of an Element in a Compound

1. State the given and needed quantities.

2. Use Avogadro's number to write conversion factors.

3. Set up the problem to cancel units and calculate the number of particles.

Example: How many atoms of O are there in 0.150 moles of aspirin, C9H8O4?

Section 7.5 & 7.6: Molar Mass and Calculations

The molar mass of an element or compound is the mass of one mole of that substance, expressed in grams. It is numerically equal to the average atomic or molecular mass from the periodic table.

• Molar Mass of Elements:

  • 1 mole C = 12.01 g C

  • 1 mole Li = 6.941 g Li

• Molar Mass of Compounds: Add the molar masses of all atoms in the formula.

Example: Lithium carbonate, Li2CO3:

• 2 moles Li × 6.941 g/mol = 13.88 g

• 1 mole C × 12.01 g/mol = 12.01 g

• 3 moles O × 16.00 g/mol = 48.00 g

• Total: 73.89 g/mol

Calculations Using Molar Mass for Compounds

• Molar mass can be used as a conversion factor between grams and moles.

• Equality: 1 mole Li2CO3 = 73.89 g Li2CO3

• Conversion factors:

  • $\frac{73.89\ \text{g Li}_2\text{CO}_3}{1\ \text{mole Li}_2\text{CO}_3}$

  • $\frac{1\ \text{mole Li}_2\text{CO}_3}{73.89\ \text{g Li}_2\text{CO}_3}$

Guide to Calculating Moles from Mass or Mass from Moles

1. State the given and needed quantities.

2. Write conversion factors, including molar mass.

3. Set up the problem to cancel units and calculate the quantity needed (grams or moles).

Example: How many moles of NaCl are in 737 g of NaCl?

Map: Mass–Moles–Particles

Mass	Moles	Particles
Grams of element	Moles of element	Atoms (or ions)
Grams of compound	Moles of compound	Molecules (or formula units)

Conversions:

• Grams ↔ Moles: Use molar mass

• Moles ↔ Particles: Use Avogadro's number

• Moles of compound ↔ Moles of element: Use formula subscripts

Section 7.7: Mole-to-Mole Relationships in Chemical Equations

Balanced chemical equations provide the mole ratios needed to relate quantities of reactants and products. These ratios are essential for stoichiometric calculations.

• Example: $2\mathrm{Ag(s)} + \mathrm{S(s)} \rightarrow \mathrm{Ag_2S(s)}$

• For every 2 moles of Ag and 1 mole of S, 1 mole of Ag2S is formed.

Law of Conservation of Mass

• Matter cannot be created or destroyed in a chemical reaction.

• The total mass of reactants equals the total mass of products.

Information Obtained from a Balanced Equation

	Reactants	Products
Equation	2Ag(s) + S(s)	Ag2S(s)
Atoms	2 Ag atoms + 1 S atom	1 Ag2S formula unit
Moles	2 moles Ag + 1 mole S	1 mole Ag2S
Mass (g)	2 × 107.9 (Ag) + 32.07 (S)	247.9 (Ag2S)

Mole-Mole Factors from a Chemical Equation (Stoichiometry)

• Mole-to-mole relationships between reactants and products are written as conversion factors.

• Example: $2\mathrm{Fe(s)} + 3\mathrm{S(s)} \rightarrow 1\mathrm{Fe_2S_3(s)}$

• Possible conversion factors:

  • $\frac{2\ \text{moles Fe}}{3\ \text{moles S}}$

  • $\frac{1\ \text{mole Fe_2S_3}}{2\ \text{moles Fe}}$

  • etc.

Guide to Calculating Quantities of Reactants and Products

1. State the given and needed quantities.

2. Write conversion factors (use coefficients for mole-to-mole ratios).

3. Set up the problem to cancel units and calculate the needed quantity.

Example: How many moles of Fe are needed for the reaction of 12.0 moles of O2 in $4\mathrm{Fe(s)} + 3\mathrm{O_2(g)} \rightarrow 2\mathrm{Fe_2O_3(s)}$?

Section 7.2: Types of Reactions

Chemical reactions are classified into several types based on the patterns of reactants and products.

• Combination (Synthesis) Reactions: Two or more elements or simple compounds combine to form one product. Example: $2\mathrm{Mg(s)} + \mathrm{O_2(g)} \rightarrow 2\mathrm{MgO(s)}$

• Decomposition Reactions: One substance splits into two or more simpler substances. Example: $2\mathrm{HgO(s)} \rightarrow 2\mathrm{Hg(l)} + \mathrm{O_2(g)}$

• Single Replacement Reactions: One element replaces another in a compound. Example: $\mathrm{Zn(s)} + 2\mathrm{HCl(aq)} \rightarrow \mathrm{ZnCl_2(aq)} + \mathrm{H_2(g)}$

• Double Replacement Reactions: Ions in two compounds exchange places. Example: $\mathrm{AgNO_3(aq)} + \mathrm{NaCl(aq)} \rightarrow \mathrm{AgCl(s)} + \mathrm{NaNO_3(aq)}$

• Combustion Reactions: A hydrocarbon reacts with O2 to produce CO2 and H2O, releasing energy. Example: $\mathrm{CH_4(g)} + 2\mathrm{O_2(g)} \rightarrow \mathrm{CO_2(g)} + 2\mathrm{H_2O(g)} + \text{energy}$

Some reactions may fit more than one type.