The Oceans Flashcards by Kiara Tooms

Carbon dioxide dissolves in water giving the overall equation

CO2(aq) + H2O(l) 2H+(aq) + CO3^2- (aq)

Product is carbonic acid

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Methods to reduce CO2 in atmosphere

More economical use of fossil fuels
Alternatives to fossil fuels
Capture and storage of CO2
Increased levels of photosynthesis

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Entropy is a measure of

Order

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Entropy change for a chemical system is

The difference between the entropies of the reactants and products. P-R

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Entropy change of the surroundings

Depends on the transfer of heat to and from the surroundings I.e enthalpy change.

Delta Ssurr = delta H (enthalpy change) / T

T is in kelvin. 0 –> +273

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Total entropy change is…

Sum of delta system add delta surroundings.

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Total entropy predictions:

Delta total is positive then a reaction will occur spontaneously
Delta total = 0 then reaction is in equilibrium

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Ionic compounds in water

Dissolve in water, ionic lattice is broken up and the ions seperate and become hydrated. Energy is used to break up lattice but energy is given out when ions are hydrated, ion- dipole bonds are formed.

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Lattice (formation) enthalpy definition :

Enthalpy change when one mole of a solid is formed from its separate gaseous ions.
Magnitude of delta LE is an indication for bond strength in the lattice
ALWAYS NEGATIVE - bonds forming

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Bigger charges and or smaller ions…._____ ionic bonds_______ lattice enthalpy

Stronger

Larger (or more negative )

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Ionic radius of an element depends on

Nuclear charge (atomic number) - the bigger the nuclear charge the smaller the ion ( for atoms with the same number of filled energy levels) 
The number of full energy levels - the more levels, the bigger the ion

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Enthalpy change of hydration definition

Enthalpy change when an aqueous solution is formed from one mole of gaseous ions.
Ion-dipole bonds form - ALWAYS NEGATIVE.

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The greater the enthalpy change of hydration the…

Stronger the ion-dipole attractions and the greater the number of water molecules surrounding the ion.

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Strength of ion - dipole attraction depends on the charge and size of the ions:

Bigger the charge&/or small the size of ions the stronger the ion -dipole bonds - larger the enthalpy of hydration.

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Enthalpy change of solution definition

Is the enthalpy change when one mole of a solute dissolves to form a dilute solution. Can be thought of as a two step process - breaking down the lattice and then hydrating the gaseous ions produced (water is solvent)

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Enthalpy change of solution =

Delta Hhyd (cation) + delta Hhyd(anion) - deltaHle

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Enthalpy change of solvation

Enthalpy change when a solution is formed from 1 mole of gaseous ions using a solvent OTHER than water.

If delta H solution is negative or slightly positive

Then solid will dissolve as entropy change will be favorable

If delta H solution is large and positive then

Solid will not dissolve even tho enthalpy change is favorable because too much energy is required.

Ionic solutes in non polar solvents

-Delta H solvation is tiny because there is little attraction between the ions and solvent.

Strong acids in water

Almost donate all their protons - said to have undergone complete dissociation
Eg HCL, sulfuric acid and nitric acid (HNO3)

Weak acids in water ..

Only a small proportion of the acid molecules donate their protons. Said to have undergone incomplete dissociation
Eg Carboxylic acids and carbonic acid

Acidity constant or acid dissociation constant. Greater the value of Ka the Stronger the acid.

Equation for Ka

= (H+) (acid-) / (acid)

All aqueous

For very weak acids we can use something instead of Ka for ease

pKa = -log (Ka)

Calculation to work out pH

pH= -log(H+ (aq))

Monoprotic acids

Have one acidic hydrogen - they release one proton per molecule

Strong acid assumptions

Because they fully dissociate we can assume that concentration of acid is the same as concentration of hydrogen ions in that solution if acid is Monoprotic

Weak acids assumptions

Because it hardly dissociates we say that acid concentration remains same. Also for every one mole of H+ there's one mole of acid-. conc =

What's needed to calculate weak acid

Concentration of the acid and its dissociation constant

Strong bases in water

Gives (OH-) is fully dissociated

Ionic product of water:

Used to calc pH of strong bases | Kw = (H+)(OH-) = 1X10-14 mol2dm-6 at 298K

Calculating pH of strong base

Fully dissociates so conc of base = conc of OH- Kw =(H+)(OH-) So (H+) = Kw/(OH-) Take -log(H+)

Buffer solution definition

Solutions that have almost constant pH, despite dilution of small additions of acid or alkali

Buffer solutions contain

Either a weak acid and on of its salts - e.g ethanoic acid and sodium ethanoate Or a weak base and one of its salts - ammonia and ammonium chloride

All buffers contain large amounts of...

Proton donor - weak acid or conjugate acid And large amounts of proton acceptor - weak base or conjugate base. Any additions react with these large amounts and maintain the stable pH

Two assumptions made in buffer calculations

All the anions (A-) Have come from the salt, so contribution from the acid is negligible Concentration of the acid in solution (HA(aq)) is the same as conc of acid put into solution.

Calculations with buffers

Ka= (H+)(A-) / (HA) Using assumptions can be shown that Ka=(H+) X ((salt)/(acid))

Finding pH of a buffer solution

If Ka of the weak acid is known and conc of both salt and weak acid then (H+) can be calculated and hence the pH can be also. Whenever solutions are mixed, the individual concentrations are reduced in proportion to their volumes in the mixture. Equal amounts = halved

Making buffer solution of a specific pH

Use Ka= (H+) X (salt)/acid | Choose right weak acid with correct Ka or pKa and then calculate ratio of acid : salt

Buffers in action

CO2 dissolved in oceans and limestone rocks (Calcium carbonate) Carbonic acid produced acts as weak acid Limestone acts as as anion sink (CO3^2-) Means as CO2 atmospheric levels increase and dissolve to form carbonic acid the pH can be maintained

Unique properties of water

High bp for a compound with low relative molecular mass Has a high specific heat capacity ( takes a lot of energy to raise temperature) large masses of water can carry a lot of energy in the oceans and distribute it across the worlds oceans. Has a high enthalpy change of vaporization Density of ice is lower than liquid water, result of bond angles - ice has a very open structure with large spaces. ALL DUE TO H BONDING, WATER CAN FORM 2 HBONDS PER MOLECULE.