Oceans
- Created by: alida
- Created on: 11-05-14 12:05
Storing Carbon Dioxide
Carbon dioxide = dissolves in oceans = soluble in water
Large amounts of carbon dioxide can be stored in oceans
Increasing carbon dioxide levels from human activity = contributes to global warming
Carbon dioxide levels increasing
= more is dissolved in the oceans
= helps slow down global warming
Sea is naturally alkaline, carbon dioxide (acidic) dissolves
= makes sea more acidic
= changes the overall pH of the oceans
= harms species, ecosystems, food chains
Reducing carbon dioxide emission
1. Use less fuel
Increasing efficiency of engines and industrial machines
= Less energy is needed
= less Carbon dioxide is released
2. Use different fuels
- Nuclear energy
produce a lot of energy
leaks of radiation = health problems
carbon dioxide is produced during mining, transporting, processing
- hydrogen
hard to store, only produces clean products, easily available
Removing Carbon dioxide
1. photosynthesis
= reforestation
= more trees
= more photosynthesis
= more carbon dioxide removed from the atmosphere
2. Carbon capture and storage
= collect carbon dioxide from power stations
= pumped into porous rocks under the sea
= stored
Buffers
Buffer = solution that resist changes in pH when small volumes of acid/alkali are added
Formed of:
- Acid buffers = weak acid + one of its salts
- Alkali buffers = weak base + one of its salts
2 assumptions:
1. All the A⁻ ions (anions) come from the salt
- contribution of anions from the acid = negligible compared to the fully ionised salt
2. The concentration of HA in the buffer remains unchanged
Buffers
Two factors which buffers depend on:
1. The Value of Ka
2. The ratio of [salt] : [acid]
- Ratio provides a 'fine tuning' of a buffers pH
- for an effective buffer, the ratio must not be too large or too small
All buffers contain a large amount of proton donors and proton acceptors
= addition of acid/alkali react with the proton donors/acceptor
= maintains the pH within limits
Buffers
HA (aq) <=========> H⁺(aq) + A⁻
Conjugate acid conjugate base
small amounts of acid are added:
= added H+ ions are removed with salt ions present in large amounts
= equilibrium shifts to the left
= maintaining the pH
small amounts of alkali added:
= weak acid dissociates
= produced more H+ ions
= react with OH- ions
= equilibrium shifts to the left
= maintains the pH
Calculations with buffers
Ka = the acidity constant/ acid dissociation constant
greater the value of Ka = stronger the acid
Ka = H⁺ x [salt] / [acid]
= H⁺ x [A⁻] / [HA]
[H+] = Ka x [acid] / [salt]
Buffers in action
carbon dioxide dissolving in water:
CO₂ (aq) + H₂O (l) <== ==> H₂CO₃ (aq)
Carbon dioxide dissolved in oceans/limestone rock
= acts as a buffering system
carbonic acid = H₂CO₃
= acts as a weak acid
limestone = CO₃²⁻
= acts as an anion
= as atomospheric carbon dioxide level increase
= dissolve to form carbonic acid
= pH can be maintained
Entropy
entropy change (ΔS) for a system: difference between the entropies of the reactants and products in an equation for a reaction
ΔSsys = ΣΔS(products) - ΔΣS(reactants)
Entropy change for the surroundings:
ΔS = -ΔH x 1000/T
T = Kelvin = +273 degrees ΔH = Joules
ΔSTotal = ΔSsys + ΔSsurr
ΔS total > 0 = positive = reaction will occur spontaneously
ΔS total = 0 = reaction is at equilibrium
entropy increases = endothermic reactions occur spontaneously
Lattice Enthalpy
Ionic compounds dissolve in water = ionic lattice is broken up = ions are seperated = become hydrated
Energy is used to break up the lattice Energy is given out when the ions are hydrated
Lattice enthalpy: Enthalpy change when 1 mole of a solid is formed from its seprate gaseous ions
The size of ΔHLE = shows the strength of the bonds in the lattice
Bigger charge/smaller ions=stronger ionic bonds=more-ve= ΔHLE
Ionic radius of an element depends on:
- Nuclear charge/ atomic number = bigger the nuclear charge = the smaller the ions are
- Number of full energy levels = more level = bigger the ion
Enthalpy change of hydration
ΔHhyd = enthalpy change when an aqueous solution is formed from 1 mole of gaseous ions
= always negative = involves an ion-dipole forming
Greater the ΔHhyd value:
= the stronger the ion dipole attractions
= greater number of water molecules surrounding the ion
The bigger the charge/smaller the size of the ions
= stronger the ion-dipole bonds
= larger the enthalpy of hydration
Enthlapy change of solution
ΔH solution = enthalpy change when 1 mole of a solution dissolves to form a dilute solution
2 step process : breaking down the lattice, and then hydrating the gaseous ions produced ( water is the solvent)
ΔH solution= ΔHhyd( cation) + ΔHhyd (anion) - ΔHLE
ΔH solvation = enthalpy change when a solution is formed from 1 mole of gaseous ions using a solvent other than water
enthalpy change of solution
ΔH solution = negative or slighlty positive
= solid will dissolve due to entropy changes being favoured
If ΔHsolution = large and positive
= solid will not dissolve
= too much energy is needed to break the bonds
= occurs in ionic solutes in non-polar solvents
= -Hsolv = tiny
= little attraction between ions and solvent
Acids and Bases
weak acids/bases= partially dissociate in aqueous solutions
Acids = not all acid molecules donate their protons = equilibrium lies to the left (side of the acid)
Bases = not all bases accept protons = equilibrium lies to the left
Stronger the Acids = the more the equilibrium lies to the right-hand side
Stronger the Base = the more the equilibrium lies to the right - hand side
Acids and Bases
Acids = proton donors= donate H+ ions
Bases = proton acceptors = accept H+ ions
Strong acids = fully dissociate in aqueous solution = all acid molecules donate their protons
= we can assume that the conc of acid put into the solution = conc of H+ ions in the solution IF the acid = monoprotic acid
monoprotic acid = release one proton per molecule = only have one hydrogen
Strong bases = fully dissociate in aqueous solution = all base molecule accept H+ ions
For strong bases/ acid = equilibrium lies to the right
Calculations with Acids/Bases
Ka = acid dissociation constant
Greater the value of Ka = stronger the acid
weak acids = smaller Ka value
pKa = -logKa
Ka= 10^-pKa
pH = -log[H+(aq)]
Kw = ionic product of water = [H+][OH-]
Kw = same value at a given temperature
neutral solution = [H+] = [OH-]
Kc = [H+][OH-] / [H2O]
Acids/Base
Assumption for weak acids =
1. Few molecules dissociate = conc of acid = conc given
2. Assume for every 1 mole [H+] = 1 mole of base
HA and A- = conjugate pairs = HA = conjugate acid of A- (vice versa)
Equilibrium with conjugate pairs = set up when a base dissolves in water
In water:
Water and H3O+ = conugate pair
Water and OH- = conjugate pair
Ionisation Enthalpy
Ionisation enthalpy = energy needed to remove one mole of electrons from one mole of gaseous ions to form one mole of gaseous cation
Increases across a period and up a group
Factors:
1. Atomic radius
Further away the outer shell from the +ve nucleus = less attraction=lower IE
2. Nuclear charge
More protons = higher attraction between nucleus and outer e' = higher the IE
3. Shielding
More shells = less attraction = lower the IE
Hydrogen Bonding
Strongest type of intermolecular bonding
- Between H and an electronegative atom (F, O, N)
- electronegative atom attracts electrons away from the H
substances won't dissolve in water unless they can form H- bond with water
The bonds in the solvent and solute break --> new bonds between the solvent and solute form
Won't dissolve if the bonds to be broken are stronger than those that will be formed
Hydrogen Bonding
Ionic substances: won't dissolve in non-polar solvents
= non polar molecules
= do not interact strongly enough with ions to pulls them away from the ionic lattice
= electrostatic forces between ions
= much stronger than the bonds that could form between ion and solvent
Most covalent substance: only dissolve in non-polar solvents
= bonds between covalent molecules
= weak
= can be broken by non-polar solvents
Hydrogen Bonding
Ionic solids = dissolve in polar solvents (water)
- The delta +ve H atom = attracted to -ve ions
- the delta -ve O atom = attracted to +ve ions
= ions seperate from the ionic lattice
= become surrounded by water molecules
= ion- dipole forms = each ion-dipole is weak
= total attraction from all the water molecules
= strong to overcome the strong ionic attraction in the lattice
= hydration ( solvent isn't water = solvation)
Properties of water due to H bonds
1. High boiling point
- H bonds = strong = lot of energy needed to overcome the H bonds in order for them to break (much higher BP than other Gp6 hydrides)
2. High specific heat capacity
- takes a lot of energy to raise its temperature
- large masses of water circulating the ocean = can carry a large amount if energy = distribute it across the worlds oceans
3. High enthalpy of Vaporisation
= amount of energy needed to change a substance from its standard state --> vapour
= water = more energy in order to evaporate
4. Ice is less dense than water
- as water freezes all H bonds that can form are formed = wasted space = less dense
- ice melts = H bonds break = lattice breaks down = water can 'fill the gaps' = increase in density
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