Enthalpy

Thermodynamics 101 First Law of Thermodynamics o Energy is conserved in a reaction (it cannot be created or destroyed)--- sound familiar??? o Math representation: ΔE total = ΔE sys + ΔE surr = 0 Δ= “change in” ΔΕ= positive (+), energy gained by system ΔΕ= negative (-), energy lost by system Total energy = sum of the energy of each part in a chemical reaction

Mg+ 2HCl  MgCl 2 + H 2

Exothermic Temperature increase (--isolated system) Heat is released to surroundings (--open/closed system) q = - value Chemical  Thermal Energy

Endothermic Temperature decrease (--isolated system) o All energy going into reaction, not into surroundings Heat absorbed by system, surroundings have to put energy into reaction q = + value Thermal  Chemical Energy

Heat of Reaction Amount of heat exchange happening between the system and its surroundings for a chemical reaction. Temperature remains constant Usually reactions happen at constant volume or constant pressure

How does work factor into heat of reaction? W = -PΔV If volume is constant (ΔV), PΔV = 0 and no other work sooooo If pressure (P) is constant so volume can change, work is being done soooo

Enthalpy (H) Measures 2 things in a chemical reaction: 1)Energy change 2)Amount of work done to or by chemical reaction 2 types of chemical reactions: 1) Exothermic —heat released to the surroundings, getting rid of heat, -ΔΗ 2) Endothermic —heat absorbed from surroundings, bringing heat in, +ΔΗ ** Enthalpy of reaction —heat from a chemical reaction which is given off or absorbed, units = kJ/mol Enthalpy of reaction o Heat from a chemical reaction which is given off or absorbed o At constant pressure o Units = kJ/mol

Enthalpy (H) cont. Most chemical reactions happen at constant pressure (atmospheric pressure)—open container Temperature and pressure are constant o Only work is through pressure/volume Sum of reaction’s internal energy + pressure/volume of system o H = U + PV o ΔH = ΔU + PΔV

Properties of Enthalpy Extensive Property o Dependent on amount of substance used State Function o Only deals with current condition o Focus on initial and final states Enthalpy changes are unique o Each condition has specific enthalpy value SO enthalpy change (ΔH) also has specific value

Example 1 CH 4 + 2O 2  CO 2 + 2H 2 O ΔH = kJ

Example 2 2HgO  2Hg + O 2 ΔH = kJ HgO  Hg + ½ O 2 ΔH = kJ

More Enthalpy The reverse of a chemical reaction will have an EQUAL but OPPOSITE enthalpy change HgO  Hg + ½ O 2 ΔH = kJ Hg + ½ O 2  HgO ΔH = kJ SOOO-----total ΔH = 0

Example 1: Based on the following: 2Ag 2 S + 2H 2 O  4Ag + 2H 2 S + O 2 ΔH = kJ Find the ΔH for the reaction below: Ag + ½ H 2 S + ¼ O 2  ½ Ag 2 S + ½ H 2 O ΔH = ?

Example 2: Write a chemical equation for ice melting at 0°C through heat absorption of 334 kJ per gram.

Stoichiometry Returns

Example 1: H 2 + Cl 2  2HCl ΔH = kJ

Example 2: Calculate the ΔH for the following reaction when 12.8 grams of hydrogen gas combine with excess chlorine gas to produce hydrochloric acid. H 2 + Cl 2  2HCl ΔH = kJ

Example 3: Pentaborane (B 5 H 9 ) burns to produce B 2 O 3 and water vapor. The ΔH for this reaction is kJ/mol at 298°K. What is the ΔH with the consumption of mol B 5 H 9 ? 2B 5 H O 2  5B 2 O 3 + 9H 2 O

Homework Study for intermolecular quiz-----Tuesday Problems p. 251 #27, 29-31, due Wednesday