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: ΔEtotal = ΔEsys + ΔEsurr = 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 MgCl2+ H2

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,

-ΔΗ2) 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 absorbedo At constant pressureo Units = kJ/mol

Enthalpy (H) cont. • Most chemical reactions happen at constant

pressure (atmospheric pressure)—open container

• Temperature and pressure are constanto Only work is through pressure/volume

• Sum of reaction’s internal energy + pressure/volume of systemo H = U + PVo 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 • CH4 + 2O2 CO2 + 2H2O ΔH = -890.3

kJ

Example 2 • 2HgO 2Hg + O2 ΔH = + 181.66 kJ

• HgO Hg + ½ O2 ΔH = + 90.83 kJ

More Enthalpy • The reverse of a chemical reaction will have an

EQUAL but OPPOSITE enthalpy change

• HgO Hg + ½ O2 ΔH = + 90.83 kJ

• Hg + ½ O2 HgO ΔH = - 90.83 kJ

• SOOO-----total ΔH = 0

Example 1: • Based on the following: • 2Ag2S + 2H2O 4Ag + 2H2S + O2 ΔH =

+595.5 kJ

• Find the ΔH for the reaction below: • Ag + ½ H2S + ¼ O2 ½ Ag2S + ½ H2O 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: • H2 + Cl2 2HCl ΔH = -184.6

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.

• H2 + Cl2 2HCl ΔH = -184.6 kJ

Example 3: • Pentaborane (B5H9) burns to produce B2O3 and

water vapor. The ΔH for this reaction is -8686.6 kJ/mol at 298°K. What is the ΔH with the consumption of 0.600 mol B5H9 ?

• 2B5H9 + 12O2 5B2O3 + 9H2O

Homework • Study for intermolecular quiz-----Tuesday

• Problems p. 251 #27, 29-31, 33-35 due Wednesday