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Transcript
```Unit 11 Energy, Changes of State, Solids and Liquids
Thermodynamics – the study of energy
Energy – the ability to do work or produce heat
Units of energy:
calorie (cal) – amount of energy required to rise the temperature of 1 gram of water 1 ˚C
Joule (J) – SI unit – 1 calorie = 4.184 J
Potential Energy – energy due to position or composition
Kinetic Energy – energy due to the motion of the object and depends on mass (KE = ½ mv 2)
Law of Conservation of Energy – energy can be converted from one form to another but can be
neither created nor destroyed
Work – force acting over a distance
Temperature – measure of the random motion of the components of a substance (average KE)
Heat – flow of energy due to a temperature difference
Solar System – part of the universe on which we want to focus attention (reactants and products)
Surroundings – everything else in the universe
Calorimeter – device used to determine the heat associated with a chemical reaction
Entropy (S) – measure of disorder and randomness
The amount of energy required to change the temperature of a substance depends on:
1.
Amount of substance (mass)
2. The amount of temperature change (ΔT)
3. Type or identity of the substance - each has a different specific heat (s) or (c)
q = m ΔT s
Specific heat capacity (s or c) – describes the amount of energy required to change the
temperature of one gram of a substance by 1 ˚C, measured in J/g ˚C
q (in joules) = energy to achieve the temperature change
m (in grams) = mass of the substance
ΔT (in ˚C) = change in temperature (Tfinal – Tinitial)
Example 1: You have a 5.63 gram sample of solid gold and heat it from 21 ˚C to 32 ˚C. How much
energy (in Joules) is required? The specific heat of gold is 0.13 J/g ˚C.
q = m ΔT s = (5.63) (11) (0.13) = 8.0 J (round to 2 sig figs)
Example 2: A 2.8 g sample of a pure metal required 10.1 Joules of energy to change its
temperature from 21 ˚C to 36 ˚C. What is the metal’s specific heat capacity? What can we do now
that we know it?
q = m ΔT s
s =
𝒒
𝒎 𝜟𝑻
=
𝟏𝟎.𝟏
(𝟐.𝟖)(𝟏𝟓)
= 0.24 J/g ⁰C (round to 2 sig figs)
Enthalpy (H) – heat for the reaction
EXOTHERMIC
-ΔH = energy released = exothermic reaction

Energy transferred out of system by heat

Products have less PE than reactants

Energy shown as a product
EXOTHERMIC
Examples:

Liquid water freezing to form ice

Water vapor condensing into dew or raindrops

Explosions

Combustion
ENDOTHERMIC
+ΔH = energy absorbed = endothermic reaction

Feels cold to touch (energy absorbed into system from surroundings)

Products have more PE than reactants

Energy shown as a reactant
Examples:

Ice cubes melting into liquid water

Liquid water evaporating and becoming water vapor

A cake baking in the oven

Barium hydroxide and ammonium nitrate (a chemical cold pack)
Review: Intermolecular forces:
 Intramolecular = bonds within a molecule = covalent
 Intermolecular = forces between molecules or ions
1.
Ion-Ion Interactions

Strongest: hold ions in a crystal structure
2. Hydrogen bonding

Hydrogen bonded to highly electronegative ion (“FON”)
3. Dipole-dipole Interactions

Polar molecules
4. London Dispersion Forces

weakest

Noble gases atoms and non-polar molecules

“momentary’ dipoles in neighbor
Heating/Cooling Curve: plot of temperature vs. time for a substance, where energy is added at a
constant rate
Molar Heat of Fusion (∆Hf) – the amount of heat necessary to melt one mole of a substance at its
melting point (KJ/mol)
q = ∆Hf (n)
n = number of moles
Example: 18 g of H2O is being melted at its melting point of 0 ˚C. How many kJ of energy are
required to melt this ice? The ∆H f of water is 6.02 kJ/mol
Use dimensional analysis because sometimes you will see 𝜟H in different units (like J/g)
18 g
1 mol
6.02 kJ
= 6.0 kJ
18.02 g
1 mol
Molar Heat of Vaporization (∆Hv) – the amount of heat necessary to boil (or vaporize) one mole of a
substance at its boiling point
q = ∆Hv (n)
n = number of moles
Example: 2.3 mole of H2O is being boiled at its boiling point of 100 ˚C. How many kJ are required to
boil this water? The ∆Hv of water is 40.7 kJ/mol
Don’t need molar mass because amount is already in moles
2.3 mol
40.7 kJ
= 94 kJ
1 mol
Phase Diagram: areas of stability of various phases in a chemical system at equilibrium
Triple Point: all 3 phases exist in equilibrium
Critical point: above this the gas will be a gas, cannot be liquified
normal boiling point: boiling temperature under one atmosphere of pressure
normal freezing point: the temperature in which a solid is formed under one atmosphere of pressure
freezing point depression: add a solute to lower freezing temperature
boiling point elevation: add a solute to raise boiling temperature
vaporization: (evaporation) the change in state that occurs when a liquid evaporates to form a gas
condensation: the process by which vapor molecules reform a liquid
vapor pressure: the pressure of the vapor over a liquid at equilibrium in a closed container
Changes of State
Can you name the 5 states of matter?
Describe (according to KMT) the states of matter:

Solid: crystals – vibrate around fixed points, high IM forces

Liquid: flow easily, lower IM forces, move more freely

Gas: small volume (point masses); constant, rapid, random motion (extremely low IM forces);
elastic (close to) collisions
What is the name for each change of state?
In the changes:



melting
vaporization
sublimation
Energy is added to the substance (absorbed)
Moving from less → more freely (by overcoming IM forces)
Increase temperature or decrease pressure
In the changes:



freezing
condensation
deposition
Energy is given off to the surroundings
Moving from more → less freely (allow IM forces to “mater” or “take over”
Decrease temperature or increase pressure
```