Chapter 7 Exam

Chemical Reactions: Types and Equations

Overview of Chemical Equations

Chemical equations represent the transformation of reactants into products. Understanding the different types of reactions and how to write their equations is fundamental in general chemistry.

• Molecular Equation: Shows all reactants and products as compounds.

• Total Ionic Equation: Shows all strong electrolytes as ions.

• Net Ionic Equation: Shows only the species that actually change during the reaction.

Major Types of Chemical Reactions

• Combustion Reactions: A substance reacts with oxygen, releasing energy as heat and light. Typically, hydrocarbons produce CO2 and H2O. Example: $\mathrm{CH_4 + 2O_2 \rightarrow CO_2 + 2H_2O}$

• Precipitation Reactions: Two aqueous solutions combine to form an insoluble solid (precipitate). Example: $\mathrm{AgNO_3(aq) + NaCl(aq) \rightarrow AgCl(s) + NaNO_3(aq)}$

• Acid-Base Neutralization Reactions: An acid reacts with a base to produce water and a salt. Example: $\mathrm{HCl(aq) + NaOH(aq) \rightarrow NaCl(aq) + H_2O(l)}$

• Gas-Evolution Reactions: Reactions that produce a gas as a product. Example: $\mathrm{Na_2CO_3(aq) + 2HCl(aq) \rightarrow 2NaCl(aq) + H_2O(l) + CO_2(g)}$

• Redox (Oxidation-Reduction) and Single-Displacement Reactions: Involve the transfer of electrons. In single-displacement, an element replaces another in a compound. Example: $\mathrm{Zn(s) + CuSO_4(aq) \rightarrow ZnSO_4(aq) + Cu(s)}$

• Synthesis Reactions: Two or more substances combine to form one product. Example: $\mathrm{2H_2(g) + O_2(g) \rightarrow 2H_2O(g)}$

• Decomposition Reactions: A single compound breaks down into two or more simpler substances. Example: $\mathrm{2H_2O_2(l) \rightarrow 2H_2O(l) + O_2(g)}$

Identifying Reaction Types

• Precipitation: Look for the formation of an insoluble solid.

• Acid-Base: Look for water and a salt as products.

• Redox: Identify changes in oxidation states.

Thermochemistry & Calorimetry (Chapter 7)

State Functions

State functions depend only on the initial and final states, not the path taken. Key state functions include:

• Internal Energy (\(\Delta E\))

• Enthalpy (\(\Delta H\))

• Pressure (P)

• Volume (V)

For state functions, only the starting and ending values matter.

First Law of Thermodynamics

The first law states that energy cannot be created or destroyed. The change in internal energy is the sum of heat and work:

• Equation: $\Delta E = q + w$

• q: Heat exchanged

• w: Work done on or by the system

Work and Heat Flow

• Work (w): For gases, work is related to volume changes at constant pressure: $w = -P_{\text{ext}} \Delta V$ If a gas expands (\(\Delta V > 0\)), work is negative (system does work on surroundings).

• Heat Flow (q): The amount of heat transferred is calculated by: $q = m s \Delta T$ Where:

  • m = mass (g)

  • s = specific heat capacity (J/g·°C)

  • \(\Delta T\) = temperature change (°C)

Calorimetry: Measuring Heat

• Coffee-Cup Calorimeter (Constant Pressure): Used for reactions in solution. Measures \(\Delta H\) (enthalpy change).

• Bomb Calorimeter (Constant Volume): Used for combustion reactions. Measures \(\Delta E\) (internal energy change).

Enthalpy Calculations

• Step-by-Step Heat Tracking:

  1. Calculate the heat absorbed or released by the water/solution using $q = m s \Delta T$.

  2. Reverse the sign to find the heat of the reaction (system vs. surroundings).

  3. Scale the heat value according to the coefficients in the balanced equation.

• Hess’s Law: The enthalpy change for a reaction is the same, no matter how many steps it takes. Manipulate equations and their \(\Delta H\) values to sum to the target reaction. Rule: If you reverse a reaction, reverse the sign of \(\Delta H\). If you multiply a reaction, multiply \(\Delta H\) by the same factor.

• Standard Enthalpies of Formation (\(\Delta H_f^\circ\)): The enthalpy change when one mole of a compound forms from its elements in their standard states. Formula: $\Delta H_{\text{rxn}}^\circ = \sum n \Delta H_f^\circ (\text{products}) - \sum n \Delta H_f^\circ (\text{reactants})$

Endothermic vs. Exothermic Reactions

• Exothermic (\(\Delta H < 0\)): Heat is released to the surroundings; the system loses energy. The container feels warm. Examples: Combustion, condensation of steam.

• Endothermic (\(\Delta H > 0\)): Heat is absorbed from the surroundings; the system gains energy. The container feels cold. Examples: Melting ice, boiling water.

Practice Focus Areas

• Calculating \(\Delta H\) from laboratory data.

• Determining if a process is endothermic or exothermic.

• Applying Hess’s Law and using standard enthalpy tables.

• Converting between units of work: $1\ \text{L} \cdot \text{atm} = 101.3\ \text{J}$

Additional info: For more complex reactions, always ensure equations are balanced before performing thermochemical calculations. When using calorimetry, remember that the heat lost by the system is gained by the surroundings and vice versa.