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Flashcards in Oceans (O) Deck (50):
1

Dissolving ionic substances in polar substances:

Most ionic substances only dissolve in polar solvents

Ionic bonds are broken when an ionic substance dissolves

Ion-dipole bonds form between solvent molecules and free ions

The strength of the bonds formed is similar to the strength of the bonds broken

Therefore, the energy released by bond formation is sufficient to compensate for the energy required to break the bond between

2

Dissolving ionic substances in non-polar substances:

Non-polar solvents don't interact strongly enough with ions to pull them away from an ionic lattice

The electrostatic forces between the ions are stronger than any bonds that could form between the ions and non-polar solvent molecules

Ion-dipole bonds formed between the solvent molecules and ions are much weaker than the bonds that are broken

So the energy released by bond formation - not sufficient to compensate for the energy required to break the bond between ions

3

Dissolving covalent substances in non-polar solvents

Most covalent substances only dissolve in non-polar solvents

Intermolecular bonds between covalent molecules = weak

They can be broken by non-polar solvent molecules

4

Dissolving covalent substances in polar solvents

Covalent substances don't tend to dissolve in polar substances as

Instantaneous dipole- induced dipole bonds between its molecules and water are weaker than the hydrogen bonds in water

5

Dissolving ionic compounds in water:

1) H-O bonds in water = polar

2) ion-dipole bonds can be formed between water molecules and the dissolved ions

3) Positive H ions are attracted to negative solute ion/ negative O ions are attracted to positive solute ion

4) Ions separate from ionic lattice + become surrounded by water molecules- hydration (if the solvent isn't water = solvation) Ionic dipole bonds pull lattice apart

6

Enthalpy change of hydration definition

Enthalpy change when 1 mole of gaseous ion is hydrated by forming bonds to water molecules

(Gas------------ Aq)

7

Hydrated ion definition

An ion bonded to a water molecule

8

When will the enthalpy change of hydration be more negative?

- Ionic charge greater

- Ionic radius is smaller

- Greater charge density of ion

Ions with greater charge density = attracts more water molecules + forms stronger ion-dipole bonds. Energy released when forming bonds is greater

9

Lattice enthalpy definition

Enthalpy change when 1 mole of an ionic solid is formed from gaseous ions

(Gas-------Solid)

10

When will the lattice enthalpy be more negative?

- Ionic charge greater

- Ionic radius is smaller

- Greater charge density of ions in lattice

Ions with greater charge density = attract each other more strongly- greater electrostatic attraction

11

Why is lattice enthalpy always negative?

Bonds are broken and energy is released (exothermic)

The more negative the lattice enthalpy, the stronger the bonding

12

Enthalpy change of solution definition

The enthalpy change when 1 mole of solute dissolves to form a solution

(Solid------ Aq)

13

Enthalpy change of solution equation

Enthalpy solution = (En hydration (cation) + En hydration (anion)) - lattice enthalpy

Enthalpy change of solution is the overall effect on the enthalpy when something dissolves

It's the net effect of the lattice enthalpy and enthalpy change of hydration

14

For a compound to dissolve, the enthalpy change of solution should be....

exothermic (negative) or very slightly endothermic

15

Techniques and procedures for measuring the enthalpy change of solution

1) Known mass of solid dissolved in a known mass of solvent + temp change measured

2) Energy exchanged with water calculated ( E= mc deltaT )

3) Scale up to find for 1 mole of solute

16

Overall, the greater the charge density of ions......

- The greater the electrostatic attraction + more exothermic the lattice enthalpy

- The greater the attraction of water molecules + more exothermic the hydration enthalpy

17

Greenhouse effect:

1) Solar energy reaches Earth mainly as visible and UV

2) Earth absorbs some of this energy, heating up and radiating IR

3) Greenhouse gases (eg carbon dioxide + methane) in troposphere absorb some of this IR in the 'IR window'

4) Absorption of IR by greenhouse gas molecules increases the vibrational energy of their bonds. The energy is transferred to other molecules by collisions, increasing their kinetic energy and raising the temperature

5) Greenhouse gas molecules also re-emit some of the absorbed IR in all directions, some of which heats up the Earth

Increased concentrations of greenhouse gases------ enhanced greenhouse effect

18

Greenhouse effect definition

Process in which infrared radiation emitted by the Earth is absorbed by gases in the atmosphere, causing a temperature increase

19

Greenhouse gasses definition

Gasses that absorb infrared radiation emitted by the Earth

20

IR window definition

The wavelengths of IR that greenhouse gasses do not absorb

21

Bronsted-Lowery theory

Acid = H+ donor

Base = H+ acceptor

22

Strong acid definition

Dissociates completely in water- the acid is fully ionised into H+ and A- ions (eg HCl)

23

Weak acid definition

Dissociates only very slightly in water- the acid undergoes incomplete ionisation into H+ and A- ions (eg ethanoic acid)

*equations showing dissociation of weak acids must contain the equilibrium arrow as the dissociation is incomplete

24

Strong base definition

Ionise completely in water (eg NaOH)

25

Buffer definition

A solution that resists pH changes when small amounts of acids or alkalis are added to it

26

What do buffer solutions contain?

either:

- A weak acid and one of its salts (eg ethanoic acid and sodium ethanoate)

- A weak base and one of its salts (eg ammonia + ammoium chloride)

All contain large amounts of proton donor + proton acceptor

27

How does a buffer maintain the pH of a solution?

Add ACID---- H+ conc increases - extra H+ ions combine with A- to form HA. Shifts equilibrium left, reducing the H+ concentration close to original value

Add BASE---- OH- concentration increases. Most extra OH- ions react with H+ ions from solution. Causes more HA to dissociate to form H+ ions, shifting equilibrium right. H+ concentration increases until it's close to original value

28

Equation for working out pH from H+ concentration

pH = -log (H+)

29

Conjugate pairs

HA-------- H+ + A-

A- = conjugate base

H+ = conjugate acid

30

What assumptions are made when working out the acidity constant (Ka) for a weak acid?

- Equilibrium concentration (HA) for a weak acid is the same as the initial concentration of the acid (as acid only slightly dissociates)

- The equilibrium concentration (A-) is equal to the equilibrium concentration (H+). A few protons will be provided by the water but these are insignificant compared to those provided by the acid

31

How to work out the pH of a weak acid

Ka = (H+)(A-) / (HA)

also can be written as:

Ka = (H+)^2 /(HA) (due to assumption 2)

32

What is pKa and how to calculate it?

Shows pH

Values for Ka can be very small for weak acids. A logarithmic scale can be used

pKa = -logKa

Ka = 10^-pKa

33

Techniques and procedures to measure the pH of a solution

TITRATION

- Add a standard solution of base to a measured quantity of acid (or vice versa)

- Known concentration goes in burette

34

pH definition

A measure of concentration of H+ ions in a solution

logarithmic scale

pH = -log(H+)

35

How to calculate the pH of a strong acid

Strong acids fully dissociate

so (H+) = concentration of acid (only is monoprotic- eg HCl, if H2SO4- x2!)

pH = -log (conc acid)

Reverse: (H+) = 10^-pH

36

How to calculate the pH of a strong base

Assume that (OH-) is equal to the concentration of the solution of the base, as it fully dissociates in aqueous solution

Use ionisation product of water (Kw)

Kw = (H+)(OH-)

Units of Kw are always mol^2dm^-6

37

Everyday applications of buffers

- Found in Shampoo, food, drink

- In blood- protect us from changes in pH due to formation of CO2 and H+ in metabolic processes- these changes would otherwise affect the action of enzymes

38

Calculating the pH of a buffer

Need to know the Ka of weak acid and the concentrations of the weak acid and its salt

Ka = (H+) x (salt)/(acid)

39

What does the value of H+ and therefore the pH of a buffer solution depend on?

- The value of Ka- provides the rough pH of buffer solution

- The ratio of (salt) : (acid)- provides the more precise pH of a buffer solution. Changing this ratio alters the pH value

40

Solubility product (Ksp) definition

An equilibrium constant for the dissolving of a sparingly soluble salt.

Represents the conditions for equilibrium between sparingly soluble solid and its saturated solution

(Sparingly soluble- ionic salt only dissolves a little bit)

41

What is the solubility product (Ksp) used to predict?

Whether a precipitate will form or not

42

How to calculate the solubility product (Ksp) for an ionic compound

Ksp = (anion) x (cation)

*But check molar ratios

43

Techniques and procedures for determining solubility product

- Need to measure concentrations of ions in solution

- Only necessary to determine concentration of one of the ions as the concentration of the other ion will be proportional to it

1) Make saturated solution of salt in distilled water. Warm distilled water in conical flask + add salt until no more salt dissolves. Allow mixture to cool, filter mixture + discard residue

2) Take temperature of solution as Ksp is temperature dependent

3) Decide appropriate method for determining conc of the ion (titration, colourimetry)

44

Entropy (S) definition

A measure of the number of ways that molecules and their associated energy quanta can be arranged

45

What is entropy used to predict?

Whether a reaction will occur spontaneously or not

46

Factors affecting entropy:

PHYSICAL STATE- gases have most disordered arrangements + highest entropy, as particles are free to move around and take up many different positions

AMOUNT OF ENERGY SUBSTANCE HAS- the more energy quanta a substance has, the more ways they can be arranged and the greater the entropy

NUMBER OF PARTICLES- more particles = higher entropy. The more particles, the more ways they and their energy can be arranged

47

Calculating the entropy change of a system:

Entropy change of products - Entropy change of reactants

48

Calculating the entropy change of the surroundings

-enthalpy change / Temperature

(enthalpy change is in J/mol and temperature in K)

49

Calculating the total entropy change:

The sum of the entropy changes of the system and the surroundings

Total entropy change = entropy change system + entropy change surrounding

also written as

Total entropy change = entropy change system + ( - enthalpy change/T)

50

What must the entropy be for a chemical change to be feasible (reaction occur spontaneously)

Total entropy change must be positive or zero