Thermochemistry and Acids/Bases Review

Introduction to Thermochemistry

• Thermochemistry (or Thermodynamics) is the study of heat, work, energy, and temperature changes during chemical reactions.

• Energy is the capacity to produce heat or do work.

  • Potential Energy: Stored in chemical bonds (e.g., gasoline).

  • Kinetic Energy: Energy due to the motion of molecules.

• Law of Conservation of Energy (1st Law of Thermodynamics): The energy of the universe is constant (neither created nor destroyed).

Temperature vs. Heat

• Temperature is the measure of the average kinetic energy of the particles in a substance.

  • Can be measured directly using a thermometer.

• Heat (q) is the energy that is transferred between two objects by collisions between the particles.

  • Cannot be measured directly.

  • Can only detect the changes caused by heat.

  • Changes include temperature and phase changes.

Thermochemical Processes

• System: What we are interested in, usually the reactants and products.

• Surroundings: Everything else, including the reaction container.

• Exothermic:

  • Heat/Energy flows out of the system.

  • Heat/Energy is a product.

  • q is negative.

• Endothermic:

  • Heat/Energy flows into the system.

  • Heat/Energy is a reactant.

  • q is positive.

Measuring Heat - Calorimetry

• The Joule (J) is the SI unit of heat and energy.

• The calorie (cal) is also used to measure heat capacity and is defined by the amount of heat needed to raise the temperature of 1g of pure water by 1°C.

  • $1 \text{ cal} = 4.184 \text{ J}$$$1 \text{ cal} = 4.184 \text{ J}$$

• NOTE: The energy contained in food is measured in Calories (Cal) = 1000 calories = 1 kcal.

Heat Capacity

• Specific heat capacity (C) is the amount of heat needed to raise the temperature of one gram of a substance by 1°C (depends on the type of substance ONLY, not the amount).

  • Has units of J/g°C or J/g K

  • $$\Delta T = (T{\text{final}} – T{\text{initial}})$$

Sample Calculation

• How much energy (heat) is given off when 1kg of sand (0.79 J/g°C) cools from 35°C to 20°C?

Formula Usage

• When 435J of heat is added to 3.4g of olive oil at 21°C, the temperature increases to 85°C. What is the specific heat of the olive oil?

  • $q = +435 \text{ J} = (3.4 \text{ g})(c)(85°\text{C} – 21°\text{C})$$$q = +435 \text{ J} = (3.4 \text{ g})(c)(85°\text{C} – 21°\text{C})$$

  • $c = (435 \text{ J})/(3.4 \text{ g})/(64°\text{C}) = 2.0 \text{ J/g°C}$$$c = (435 \text{ J})/(3.4 \text{ g})/(64°\text{C}) = 2.0 \text{ J/g°C}$$

Potential Energy Diagrams

• Potential Energy Diagrams (PEDs) help us to see how energy changes during a chemical reaction.

Activation Energy

• Where does it come from?

  • Normal molecular collisions

  • Heat (sparks) – which speeds up the collisions.

  • Light – which can affect electrons/bonds

• Activation energy can also be lowered by catalysts by changing the pathway.

Enthalpy of Reaction

• The change in enthalpy, $\Delta H$$$\Delta H$$, is the enthalpy of the products minus the enthalpy of the reactants:

  • $$\Delta H{rxn} = H{products} − H_{reactants}$$

Enthalpy of Reaction (Continued)

• This quantity, $\Delta H_{rxn}$$$\Delta H_{rxn}$$, is called the enthalpy of reaction, or the heat of reaction.

• Represented by the thermochemical equation:

  • $$CH4 (g) + 2 O2 (g) \rightarrow CO2 (g) + 2 H2O (g) + 890 \text{ kJ}$$

  • or, more properly,

  • $$CH4 (g) + 2 O2 (g) \rightarrow CO2 (g) + 2 H2O (g) \quad \Delta H = -890 \text{ kJ}$$

Worksheet Examples

• Exothermic or endothermic? Sketch the PED.

  • $$2 Fe (s) + 3 CO2 (g) + 26.33 \text{ kJ} \rightarrow Fe2O_3 (s) + 3 CO (g)$$

• Write the thermochemical equation for the formation of hydrogen chloride gas from hydrogen gas and chlorine gas that gives off 184 kJ for heat.

  • $$H2 (g) + Cl2 (g) \rightarrow 2 HCl (g) + 184 \text{ kJ}$$

Acids and Bases

• Acid – substance that produces $H^+$$$H^+$$ ions when dissolved in water –`$HX \rightleftharpoons H^+(aq) + X^-(aq)$$$HX \rightleftharpoons H^+(aq) + X^-(aq)$$

• Base – produces $OH^-$$$OH^-$$ ions in water – $MOH \rightleftharpoons M^+(aq) + OH^-(aq)$$$MOH \rightleftharpoons M^+(aq) + OH^-(aq)$$

• These are the Arrhenius’ definitions.

Hydronium Ion

• Technically, you will probably never find a free proton ($H^+$$$H^+$$) floating around in water.

• Often, you will see in acid/base reaction equations either $H^+$$$H^+$$ or $H_3O^+$$$H_3O^+$$.

  • Don’t get confused by this.

Acid-Base Reactions

• In an acid-base reaction, the acid donates a proton ($H^+$$$H^+$$) to the base.

• e two types of definitions for acid-base reactions mentioned in the notes are:

  1. Arrhenius Definition: Acids produce $H^+$$$H^+$$ ions when dissolved in water, while bases produce $OH^-$$$OH^-$$ ions in water.

  2. Brønsted-Lowry Definition: In an acid-base reaction, the acid donates a proton ($H^+$$$H^+$$

Neutralization Reactions

• Generally, when solutions of an acid and a base are combined, the products are a salt and water.

  • $$CH3COOH (aq) + NaOH (aq) \longrightarrow CH3COONa (aq) + H_2O (l)$$

  • $HCl (aq) + NaOH (aq) \longrightarrow NaCl (aq) + H_2O (l)$$$HCl (aq) + NaOH (aq) \longrightarrow NaCl (aq) + H_2O (l)$$

Dissociation of Acids

• Strong acids and bases dissociate completely into ions.

  • Concentration of $H^+/OH^-$$$H^+/OH^-$$ ions from concentration of original acid/base

• Weak acids/bases dissociate only a little, leaving mostly intact molecules.

Selected Acids and Bases

• Acids

  • Strong: hydrochloric acid (HCl), hydrobromic acid (HBr), hydroiodic acid (HI), nitric acid (HNO3), sulfuric acid (H2SO4), (per)chloric acid (HClO3, HClO4)

  • Weak: hydrofluoric acid (HF), phosphoric acid (H3PO4), acetic acid (CH3COOH or HC2H3O2)

• Bases

  • Strong: Alkali metal hydroxides, calcium hydroxide (Ca(OH)2), strontium hydroxide (Sr(OH)2), barium hydroxide (Ba(OH)2)

  • Weak: ammonia (NH3)

Acid and Base Strength

• Strong acids are completely dissociated in water.

• Weak acids only dissociate partially in water.

Acid and Base Strength (Continued)

• Acid/base strengths are quantified by the extent of dissociation

  • e.g., how much product you get from $HF(aq) \rightleftharpoons H^+ (aq) + F^- (aq)$$$HF(aq) \rightleftharpoons H^+ (aq) + F^- (aq)$$

pH

• pH is defined as the negative base-10 logarithm of a solution’s concentration of hydronium ion.

  • $pH = -\log [H_3O^+]$$$pH = -\log [H_3O^+]$$ or $pH = -\log [H^+]$$$pH = -\log [H^+]$$

pH (Continued)

• Therefore, in pure water,

  • $pH = -\log (1.0 \times 10^{-7}) = 7.00$$$pH = -\log (1.0 \times 10^{-7}) = 7.00$$

• Solution Type

  • Acidic:

    • $[H^+] > 1.0 \times 10^{-7}$$$[H^+] > 1.0 \times 10^{-7}$$

    • $[OH^-] < 1.0 \times 10^{-7}$$$[OH^-] < 1.0 \times 10^{-7}$$

    • $pH < 7.00$$$pH < 7.00$$

  • Neutral:

    • $[H^+] = 1.0 \times 10^{-7}$$$[H^+] = 1.0 \times 10^{-7}$$

    • $[OH^-] = 1.0 \times 10^{-7}$$$[OH^-] = 1.0 \times 10^{-7}$$

    • $pH = 7.00$$$pH = 7.00$$

  • Basic:

    • $[H^+] < 1.0 \times 10^{-7}$$$[H^+] < 1.0 \times 10^{-7}$$

    • $[OH^-] > 1.0 \times 10^{-7}$$$[OH^-] > 1.0 \times 10^{-7}$$

    • $pH > 7.00$$$pH > 7.00$$

How Do We Measure pH?

• For less accurate measurements, one can use:

  • Litmus paper

    • “Red” paper turns blue above ~pH = 8

    • “Blue” paper turns red below ~pH = 5

  • Or an indicator.

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