Enthalpy

Thermodynamics 101

First Law of Thermodynamics

Energy is conserved in a reaction (it cannot be created or destroyed)---sound familiar???

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) 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

Work in terms of energy change

System DOES work------ POSITIVE work value for system, system is LOSING energy

System has work on ON it----NEGATIVE work value for system, system is GAINING energy

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 Heat from a chemical reaction which is given off or absorbed At constant pressure Units = kJ/mol

Enthalpy (H) cont.

Most chemical reactions happen at constant pressure (atmospheric pressure)—open container

Temperature and pressure are constant Only work is through pressure/volume

Sum of reaction’s internal energy + pressure/volume of system H = U + PV ΔH = ΔU + PΔV

Properties of Enthalpy

Extensive Property Dependent on amount of substance used

State Function Only deals with current condition Focus on initial and final states

Enthalpy changes are unique 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

Calorimetry

How do we find the change in energy/heat transfer that occurs in

chemical reactions???

Calorimetry

Experimentally “measuring” heat transfer for a chemical reaction or chemical compound

Calorimeter Instrument used to determine the heat transfer of a chemical

reaction Determines how much energy is in food Observing temperature change within water around a reaction

container

** assume a closed system, isolated container No matter, no heat/energy lost Constant volume

Specific Heat Capacity

Amount of heat required to increase the temperature of 1g of a chemical substance by 1°C

Units--- J/g°K

Unique to each chemical substance Al(s) = 0.901J/g°K

H2O(l) = 4.18 J/g°K

q = smΔT

Example 1

How much heat is needed to raise the temperature of a 500g iron bar from 25° to 50°C ?

“Coffee Cup” calorimeter

Styrofoam cup with known water mass in calorimeter Assume no heat loss on walls Initial water temp and then chemical placed

inside Final temperature recorded

Any temperature increase has to be from the heat lost by the substance SOOO All the heat lost from the chemical reaction or

substance is transferred to H2O in calorimeter

“Coffee Cup” calorimeter (cont.)

qchemical = -qwater

Example 2: Using the following data,

determine the metal’s specific heat.

Metal mass = 25.0g Water mass = 20.0g

Temperature of large water sample = 95°C

Initial temperature in calorimeter = 24.5°C

Final temperature in calorimeter = 47.2°C

Specific heat of water = 1.00 cal/g°C OR 4.184 J/g°K (KNOW!!!!)

Δqrxn Heat gained/lost in experiment

with calorimeter

ΔHrxn

Heat gained/lost in terms of the balanced chemical equation

Example 3:

A 50.0 ml sample of 0.250M HCl and 50.0 ml sample of 0.250M NaOH react in a cofee cup calorimeter. The temperature increases from 19.50°C to 21.21°C. Calculate the ΔH for this reaction.

Homework

pp. 251-252 #25, 27, 33-35