1 Chemical reactivity Flashcards
Define ‘energy’
the ability to work
Define ‘work’
the distance moved against an opposing force
What are both work and energy measured in?
Joules
What is 1 joule the same as?
the amount of energy required to raise a 1kg substance 1 cm up against the force of gravity
What is 1 calorie the same as?
the amount of heat necessary to raise the temperature of exactly one gram of water by one degree Celsius
What is one food calorie equivalent to?
1 Kcal or 1,000 calories
What is potential energy?
the stored energy of position
State the formula for potential energy.
PE = mgh
State the formula for kinetic energy
KE = 1/2mv^2
What is ‘electromagnetic’ energy?
a form of energy that can be reflected/emitted from objects through electric or magnetic waves travelling through space
What is ‘nuclear’ energy?
the energy released during nuclear fission or fusion, especially when used to generate electricity
What is ‘chemical’ energy?
energy stored in bonds
State the ‘first law of thermodynamics.’
energy cannot be created or destroyed by any physical or chemical changes, it can only be converted from one form to another (except nuclear)
State and explain Einstein’s equation of special relativity related to the ‘first law of thermodynamics.’
e = mc²
if matter is destroyed, energy is created, and if energy is destroyed, mass is created
State 2 energy transformations
chemical, motion, radiant, chemical
How do reactions occur?
By interaction of molecules
Collision theory: reactions occur when particles collide, BUT not all collisions lead to a reaction
Which 2 conditions must reactants meet in order for a chemical reaction to occur?
must be energetic
and orientated correctly
State the 3 main stages of a successful chemical reaction
- collision between reactant particles
- breaking of chemical bonds
- right orientation of reactants will lead to a new product
What happens to energy in chemical reactions?
It is either released or absorbed
What is the ‘enthalpy change of reaction, ΔH?’
change in energy which is the difference between bond energies of reactants and products
What is meant by ‘enthalpy?’
A measure of the heat content of a substance at constant pressure
What type of energy is enthalpy and how can it be measured?
It is internal energy
the actual enthalpy of a substance cannot be measured but the enthalpy change can
Define ‘enthalpy change.’
heat released or absorbed during a chemical reaction at standard conditions
State the formula for enthalpy change.
ΔH = ΔH(products) - ΔH(reactants)
Why might enthalpy values vary?
according to conditions
What are standard conditions for an element?
form the element exists in under conditions of 1 atm and 25 Degrees Celcius
What are standard conditions for a compound if:
a) a gas
b) a pure substance
c) a substance present in solution
a) at a pressure of exactly 1 atm
b) in a condensed state (liquid/solid), at 1 atm
c) at a concentration of exactly 1M
For an exothermic reaction:
a) how does enthalpy of reactants and products compare
b) what is the value
c) what occurs with heat energy
d) draw a graph labelling ΔH and Ea
a) enthalpy of reactants > products
b) ΔH is negative
c) heat is given out/released to surroundings
d) see document
Give 3 examples of an exothermic reaction.
- burning
- respiration
- production of quicklime
State what happens to the temperature of the system in exothermic reactions.
It will be observed to rise
For an endothermic reaction:
a) how does enthalpy of reactants and products compare
b) what is the value
c) what occurs with heat energy
d) draw a graph labelling ΔH and Ea
a) enthalpy of products > reactants
b) ΔH is positive
c) heat is absorbed/taken in from surroundings
d) see document
Give 3 examples of an endothermic reaction.
- photosynthesis
- thermal decomposition of limestone
- an electron being excited from n=1 to n=3
What happens in terms of energy in bond-making and bond-breaking?
bond-breaking = add energy bond-making = release energy
Define ‘entropy.’
measures the amount of disorder of a system
Finish the following sentence:
The greater the amount of disorder, or randomness of the particles in a substance/mixture …
… the higher the entropy
State the formula with units for change in entropy.
ΔS = S(final) - S(initial) S = measured in J/mol/K
State the formula for total entropy change.
ΔTotal(total) = ΔS(system) - ΔS(surroundings)
State the ‘second law of thermodynamics’ and explain why entropy will always decrease
entropy tends towards a maximum
What do all spontaneously-occurring chemical and physical changes involve?
an overall increase in entropy
Give 5 examples where entropy increases.
- solids melting
- liquids boiling
- number of molecules increases
- ionic solids dissolve
- temperature increases
What is ‘Gibbs free energy?’
energy from a reaction free to do work
State the formula for Gibbs free energy (including units).
ΔG = ΔH - TΔS ΔG in J/mol ΔH in KJ/mol T in Kelvin (K) ΔS J/K/mol
Interpret the spontaneity of each ΔG value:
a) ΔG < 0
b) ΔG > 0
c) ΔG = 0
a) reaction will be spontaneous
b) reaction needs energy input to happen
c) the system is in equilibrium
What ΔG value corresponds to exergonic and endergonic reactions?
exergonic = ΔG negative endergonic = ΔG positive
Draw a table to show how ΔH, ΔS and -TΔS affect the value of ΔG.
see document
State the formula for the temperature for a reaction to occur spontaneously and state units required.
T = ΔH/ΔS
T = Kelvin (K)
ΔH and ΔS = must be in same units
(here, ΔG is assumed to be zero)
Define a ‘catabolic’ reaction.
high-energy, complex, compounds are broken down into simple ones
Give 2 examples of catabolic reactions.
- Glucose + 6O^2 → 6CO^2 + 6H^2O
2. ATP + H^2O → ADP + Pi (phosphate transfer)
Define an ‘anabolic’ reaction.
building-up of complex molcules from smaller ones
Give 2 examples of an anabolic reaction.
takes place in steps coupled with ATP hydrolysis or another exergonic reaction
Give 4 reactions that are coupled in metabolism.
glucose + fructose → sucrose
glucose + ATP → glucose-P + ADP
fructose + ATP → Fructose-P + ADP
glucose-P + fructose-P → sucrose + 2Pi
