.4 Energetics Flashcards
Draw graph for exothermic + endothermic and give basic info and examples
Endothermic Heat absorbed Enthalpy of reactants < products ΔH = positive Examples: Respiration Combustion of fuels
Exothermic: Heat given out Enthalpy of reactants > products ΔH = negative Examples: Photosynthesis Thermal decomposition of calcium carbonate
Introduction to Energetics and 1st Law of Thermodynamics
All chemical reactions are accompanied with some form of energy change. Usually heat energy is either absorbed or given out*. The products and reactants in a chemical reaction have two main forms (or classes) of energy:
Chemical potential energy
Kinetic energy
As the 1st Law of Thermodynamics states ‘Energy can be neither created nor destroyed but it can be converted from one form to another’.
ΔH (enthalpy change)- why can’t measure directly, how to measure indirectly
We cannot measure enthalpy directly because we cannot measure absolute gravitational or kinetic energy. However we can calculate ΔH indirectly.
The symbol used to indicate change in enthalpy is ΔH. ΔH is the heat energy change at constant pressure when:
Reactants → Products
ΔH = Hproducts – Hreactants
(same as: ΔH = Hfinal – Hinitial)
Standard conditions and symbol
Enthalpy values vary according to the conditions.
Standard state conditions are said to be: 1. Pressure: 100 kPa (100 000 Pa) (1 atmosphere)
2. A stated temperature usually 298K (= 25oC)
3. Any solutions are of concentration 1 mol dm-3.
Standard state standard conditions are indicated by ΔHϴ298.
Key definitions of diff standard enthalpies
Standard enthalpy of combustion. Enthalpy change when 1 mole of a substance is completely burned in oxygen with all reactants and products in standard states under standard conditions.
Standard enthalpy of formation. Enthalpy change when 1 mole of a substance is formed from its constituent elements with all reactants and products in standard states under standard conditions.
Standard enthalpy of neutralisation. Enthalpy change when 1 mole of water is formed in a reaction between an acid and alkali under standard conditions
Key definitions of diff standard enthalpies and examples with CH4
Standard enthalpy of combustion. Enthalpy change when 1 mole of a substance is completely burned in oxygen with all reactants and products in standard states under standard conditions.
e.g. CH4 CH4(g) + 2O2(g) 🡪 CO2(g) + 2H2O(l)
Standard enthalpy of formation. Enthalpy change when 1 mole of a substance is formed from its constituent elements with all reactants and products in standard states under standard conditions.
e.g. CH4 C(s) + 2H2(g) 🡪 CH4(g)
Standard enthalpy of neutralisation. Enthalpy change when 1 mole of water is formed in a reaction between an acid and alkali under standard conditions
e.g. HCl + NaOH HCl(aq) + NaOH(aq) 🡪 NaCl(aq) + H2O(l)
Flame calorimeter
an improved version of the simple calorimeters used for measuring enthalpy changes of combustion. Itincorporates the following features that are designed to reduce heat loss even further:
- the spiral chimney is made of copper
- the flame is enclosed
- the fuel burns in pure oxygen, rather than air.
Simple Calorimeter and draw apparatus
an apparatus for measuring the amount of heat involved in a chemical reaction or other process.
Follows the basic principle of heat being transferred to the water. As the specific heat capacity of water is known, we can work backwards to calculate how much heat was given out by the burning fuel.
Calorimetry- give examples for 1g of water hotter by 1C, 3g of water hotter by 10C
The science of measuring the heat of chemical reactions, i.e. calculating enthalpy change.
1 ºC hotter 1 g of water
Energy required = 4.18J
10 ºC hotter 3 g of water
Energy required = 3 x 10 x 4.18J = 125.4 J
Energy required equation, what they all stand for, units, how to use to work out ΔH
Energy required (q) = mass heated (m) x energy needed to make 1g of substance 1ºC hotter x temperature rise (ΔT)
q = mcΔT
Units: q = J; m = g; T = K
ΔH= Q/mol
Reactions in Solution (Cup Calorimeter)
The reaction is carried out in an insulated beaker and the temperature change measured. The reaction must be fast so that the maximum temperature is reached quickly (before it starts to cool). The specific heat capacity of the solution is taken as 4.18 Jg-1K-1 (the same as water) and we usually use the mass of the water (not the solution) in the calculation.
Biggest source of error= heat loss (for exothermic)
50 cm3 of 1.0 mol dm-3 hydrochloric acid was added to 50 cm3 of 1.0 mol dm-3 sodium hydroxide solution. The temperature rose by 6.8ºC. Calculate the standard enthalpy of neutralisation for this reaction. Assume that the density of the solution is 1.00 g cm-3; the specific heat capacity of the solution is 4.18 Jg-1K-1.
q = mc∆T m = 100 c = 4.18 ∆T = 6.8 q
= 100 x 4.18 x 6.8 = 2842 J
∆H = q / mol
Mol HCl = conc x vol = 1.0 x 50/1000 = 0.050
(Mol NaOH = conc x vol = 1.0 x 50/1000 = 0.050)
HCl + NaOH → NaCl + H2O
∆H = –2.842 / 0.050 = -56.8 kJ mol-1 (3 sig fig)
Consider this reaction: CH4(g) + 2O2(g) -> CO2(g) + 2H2O(l) ΔHθ298=-890kJmol-1 What does the minus sign indicate?
Reaction is EXOTHERMIC- products lower energy than reactants
Pros and Cons of Cup Calorimetry
PROS
1. Heat generated is from the solutions themselves so only has to be kept in calorimeter.
- Polystyrene cups are very good insulators (often you will use a cup within another cup).
- They have a low heat capacity, so absorb very little heat.
CONS
1. Inaccuracy in measuring mass and temperature
- Solution is assumed to be pure water (therefore c is assumed to be 4.18 Jg-1K-1)
- Some heat is absorbed by the cup.
- Heat is lost (or gained) from the surroundings. Not a completely closed system.
Bomb Calorimeter
The most accurate type of calorimeter.
A known mass of a fuel is placed inside a steel container (the “bomb”) and the container filled with oxygen under
pressure. The fuel is electrically ignited and heat evolved used to heat surrounding water. The heat capacity of the calorimeter is found by burning a substance with accurately known ΔHc.
