AF

Thermochemistry and Energy Systems

Event Information

  • Event Title: Purdue Pulse Fall Tailgate

  • Date and Time:

    • Saturday, October 4, 2025, at 10:00 AM EDT

    • Until Saturday, October 4, 2025, at 11:30 AM EDT

  • Location: Stadium Mall

  • Description: Come join us on Stadium Mall before the Purdue vs. Illinois Football game for food, games, gameday gear, and fun!

  • Perks:

    • Free Food

    • Free Stuff

  • Host Organization: NPSUB (Purdue Student Union Board)

Thermochemistry Topics

  • 6.1 Forms of Energy and their Interconversion

  • 6.2 Enthalpy: Changes at Constant Pressure

  • 6.3 Calorimetry

Announcements

  • Office Hours: My office hour is Wednesdays, 11:00 AM – 12:00 PM in WTHR 261 or by appointment (email mille201@purdue.edu with 3 suggested days and times to schedule it).

  • Feasting with Faculty: Friday, October 3, about 12:50 – 1:50 PM, near the desserts

  • Exam 1 Results: Grading is almost completed!

  • Final Exam: Thursday, December 18, 10:30 AM – 12:30 PM, in Elliott Hall of Music for WL students; location TBA for Indianapolis students

Ion Formation and Isoelectronic Species

  • Examples of Ion Formation:

    • Sodium (Na):

      • Electronic Configuration: $1s^2 2s^2 2p^6 3s^1$

    • Sodium Cation (Na+):

      • Electronic Configuration: $1s^2 2s^2 2p^6$

    • Fluorine (F):

      • Electronic Configuration: $1s^2 2s^2 2p^5$

    • Fluoride Anion (F−):

      • Electronic Configuration: $1s^2 2s^2 2p^6$

    • Neon (Ne):

      • Electronic Configuration: $1s^2 2s^2 2p^6$

  • Isoelectronic Species: Na+, F−, and Ne are isoelectronic; they have the same number of electrons (10 electrons).

Ion Formation in Transition Metals

  • General Rule:

    • Remove the 4s electrons before removing the 3d electrons.

  • Electron Configuration Examples:

    • Cobalt (Co):

      • Original: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^7$ or $[Ar] 4s^2 3d^7$

      • Cobalt(II) Ion (Co2+):

        • Configuration after losing two 4s electrons: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^7$ or $[Ar] 3d^7$

      • Cobalt(III) Ion (Co3+):

        • Config.: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^6$ or $[Ar] 3d^6$

Characteristics of Monatomic Ions

  • Main-group Elements: Most form one monatomic ion.

  • Transition Elements: Most form two monatomic ions.

  • Ionic Size Relationships:

    • Cations are smaller than their neutral parent atoms.

    • Anions are larger than their neutral parent atoms.

Ionic vs. Atomic Radii

  • Ionic Radii Trends:

    • As ionic charge increases (positive charge), ionic size decreases.

    • Example Series of Ionic Radii (in pm):

      • N3− > O2− > F− > Na+ > Mg2+ > Al3+

      • 146 pm > 140 pm > 133 pm > 102 pm > 72 pm > 54 pm

Periodic Properties

  • Metals:

    • Characteristics: Lustrous (shiny), malleable, ductile, good conductors of heat and electricity; mostly solids at room temperature.

  • Nonmetals:

    • Characteristics: Dull, brittle, nonconductors (insulators); some are solid, but many are gases (e.g., Br2 is a liquid).

  • Metalloids:

    • Characteristics: Have properties of both metals and nonmetals; shiny but brittle, semiconductors.

The Modern Periodic Table

  • Main Groups:

    • Group 1: Alkali metals

    • Group 2: Alkaline earth metals

  • Transition Elements:

    • Metals (main-group, transition, and inner transition)

    • Metalloids, Nonmetals, Noble gases.

  • Atomic Numbers and Masses: Example elements listed with atomic numbers and masses for clarity (e.g. H, He, Li, Na).

Metallic Behavior

  • Trends in Metals and Nonmetals:

    • Metals lose electrons (low ionization energy [IE]; low electron affinity [EA]) to form positive ions (cations).

    • Nonmetals gain electrons (high IE, high EA) to form negative ions (anions).

  • Zeff Trends:

    • Effective nuclear charge increases with nonconstant n.

Summary of Atomic Properties Trends

  • Trends:

    • Atomic radius, ionization energy, electron affinity, nonmetallic character, metallic character.

  • Periodic Trends Problem Sets:

    • Problems from previous chapters and end-of-chapter questions listed for practice.

Learning Objectives - Chapter 6

Fundamental Concepts

  • Explain the distinction between a system and its surroundings. (§6.1)

  • Categorize the transfer of energy to or from a system as heat and/or work. (§6.1)

  • Relate internal energy change, heat, and work. (§6.1)

  • State the meaning of energy conservation. (§6.1)

  • Explain the meaning of state functions and why ΔE is a state function while q and w are not. (§6.1)

  • Explain the meaning of enthalpy, and describe the relationship between ΔE and ΔH. (§6.2)

  • Explain the difference between exothermic and endothermic processes. (§6.2)

  • Relate specific heat capacity and heat. (§6.3)

  • Describe how constant-pressure (coffee-cup) calorimeters are constructed, and how they can be used to experimentally determine values of ΔH. (§6.3)

Skills

  • Determine the change in a system's internal energy (SP 6.1)

  • Solve problems involving specific heat capacity and heat transferred in a reaction. (SPs 6.4-6.7)

Introduction to Energy

  • Definition: Energy is the capacity to do work or transfer heat.

  • Focus: Mechanical energy, due to an object’s motion, position, or both.

Work and Heat Transfer

  • Work (w):

    • Example 1: Person pushing a car, w (+).

    • Example 2: A moving car hitting a person, w (−).

  • Heat Transfer (q):

    • Example 1: Heat transferred from a hand to an ice cube (q (+)).

    • Example 2: Bowl of cold water left in the freezer (q (−)).

Internal Energy

  • First Law of Thermodynamics:

    • Internal energy change relationship expressed as:
      \Delta E = q + w

  • Interpretation of Signs:

    • Positive \Delta E indicates an increase in the system's internal energy.

    • Negative \Delta E indicates a decrease.

Understanding State Functions

  • Definition of a State Function:

    • Depends only on the state's current condition and not on how that state was reached.

    • Internal energy (ΔE) is a state function, while q and w are not.

PV Work and Constant Pressure

  • PV Work:

    • As given by the equation:
      \Delta E = q + w

    • w = -P\Delta V

  • Constant Pressure Work:

    • Study reactions to determine if the system does work on the surroundings or vice versa.

Enthalpy Definition

  • Enthalpy (H):

    • Defined as H = E + PV

    • Derived relation: \Delta H = \Delta E + P\Delta V

Enthalpy in Reactions

  • Enthalpy Change Relation:

    • \Delta H = \Delta E at constant pressure.

  • Situations:

    • ΔH ≈ ΔE for specific conditions:

      • Chemical reactions without gases.

      • Reactions where moles of gas remain unchanged.

Exothermic vs Endothermic Processes

  • Exothermic Process:

    • Heat released (ΔH < 0).

    • Example: CH4(g) + 2 O2(g) \rightarrow CO2(g) + 2 H2O(g)

  • Endothermic Process:

    • Heat absorbed (ΔH > 0).

    • Example: H2O(s) \rightarrow H2O(l)

Everyday Connection to Thermochemistry

  • Hand Warmers:

    • Reaction between iron powder and oxygen, producing iron oxide and releasing heat (exothermic).

  • Cold Packs:

    • Dissolving ammonium nitrate absorbs heat from surroundings (endothermic).

Specific Heat Capacity

  • Formula: q = m \times c \times ΔT

    • Where:

      • m = mass of substance,

      • c = specific heat of substance,

      • ΔT = temperature change in °C or K.

  • Unit of Measurement:

    • Specific heat is typically in J/(g·°C).

Specific Heats of Selected Substances

  • Examples of Specific Heats (J/g·°C):

    • CO2: 0.852

    • H2O(l): 4.184

    • Cu: 0.385

    • Fe: 0.442

  • Implication: Materials with higher specific heats resist temperature changes more.

Calorimetry - Coffee Cup Calorimeter

  • Measurement Method:

    • For a solid object in liquid:

      • $q = m \times c \times ΔT$

    • Objective: Calculate heat transferred at constant pressure.

  • Examples:

    • When an object is added to a liquid, apply heat conservation:

      • q{sample} + q{water} = 0

Practice Problem with Calorimetry

  • Heat Calculation:

    • Solve for specific heat using experimental data from reactions.

  • Example Calculations:

    • Use volumetric conversions and specific heat formulas to process measured temperature changes.

Application of ΔH Calculation of Reactions

  • Example Reaction Data:

    • Mixing HCl and NaOH in calorimetry to measure thermal changes and result in enthalpy calculations.

  • Effect of Solution Changes:

    • Track heat absorption or release signs and determine endothermic vs. exothermic character during dilution or dissolution processes.

Exam Review and Practice

  • Key Problems:

    • Identify signs of energy transfer in calorimetric practices and solve for unknowns based on heat capacities and reaction outputs, ensuring thorough understanding through diverse practice problems.

Conclusion and Summary

  • The study explores core concepts in thermochemistry by interlinking definitions, examples, and real-world implications.

  • Understanding of enthalpy, heat transfer, and calorimetry allows for accurate predictions and calculations in chemical systems.

  • Proficiency in solving specific heat and heat transfer problems is vital for academic success in thermochemistry and physical chemistry courses.

Additional Thoughts

  • Regular practice of previous exam problems enhances understanding and improves application in practical scenarios regarding heat and energy transformations in chemical reactions.