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A Level Chemistry OCR B > Oceans > Flashcards

Oceans Flashcards

1
Q

Define Lattice Enthalpy.

A

The enthalpy change when one mole of solid is formed by the coming together of separate ions in the gaseous state. Always exothermic.

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2
Q

Describe the trend in lattice enthalpy.

A

Lattice enthalpy depends on the size and charge of the ions.
It becomes more negative when:
· The ionic charges increase (higher charges attract more strongly)
· The ionic radii decrease (the smaller the radius the closer the ions can get to each other causing stronger attraction).
These two factors make the lattice enthalpy more negative by increasing the charge density.

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3
Q

Define enthalpy change of hydration.

A

Enthalpy change for the formation of a solution of ions from one mole of gaseous ions. Always exothermic.

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4
Q

Describe the trend in enthalpy change of hydration.

A

Enthalpy change of hydration depends on the size and charge of the ions.
It becomes more negative when:
· The ionic charges increase (higher charges attract more strongly)
· The ionic radii decrease (the smaller the radius the closer the ions can get to each other causing stronger attraction)
These two factors make the enthalpy change of hydration more negative by increasing the charge density.

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5
Q

Define enthalpy of solvation.

A

When a solvent other than water is used for ionic compounds.

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6
Q

Define ionic bonding.

A

The electrostatic attraction between oppositely charged ions.

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7
Q

Draw a diagram to represent ionic bonding, include labels.

A
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8
Q

Define hydrogen bonding.

A

A strong dipole–dipole attraction between an electron-deficient hydrogen atom (O–Hδ+, N–H δ+ or F–H δ+) on one molecule and a lone pair of electrons on a highly electronegative atom (H-O:δ–, H-N: δ–, H-F: δ–) on a different molecule.

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9
Q

Draw a diagram of hydrogen bonding between two water molecules.

A
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10
Q

Define enthalpy change of solution.

A

The enthalpy change that takes place when one mole of a compound is completely dissolved in water under standard conditions to form a dilute solution – can be exothermic or endothermic (-/+).

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11
Q

State the equation to calculate enthalpy change of solution.

A
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12
Q

Draw an enthalpy cycle.

A
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13
Q

Draw an enthalpy profile/level diagram for a soluble salt, include labels.

A

Here the ΔHsolution is negative so energy is released. This process is energetically favourable so the solid will normally dissolve.

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14
Q

Draw an enthalpy profile/level diagram for an insoluble salt, include labels.

A

Here the ΔHsolution is positive so energy is required. This process is energetically unfavourable so the solid will not dissolve.

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15
Q

Draw the enthalpy profile for a solute with a slightly positive ΔHsolution.

A

Here the Here the ΔHsolution is only positive so energy is required. This process appears energetically unfavourable but the solid will still dissolve.

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16
Q

Why do non polar solvents not dissolve ionic solids.

A

They do not have dipoles so do not interact strongly with ions. ΔHsolution is a large positive value so dissolving is unlikely.

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17
Q

Define acid.

A

Proton (H+) donor

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18
Q

Define base.

A

Proton (H+) acceptor

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19
Q

Define basic oxide.

A

Reacts with acids to neutralise them.

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20
Q

Define strong acid.

A

An acid that completely dissociates in solution.

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21
Q

Define weak acid.

A

An acid that partially dissociates in solution.

22
Q

Define strong base.

A

A base that completely dissociates in solution.

23
Q

Define Buffer.

A

A system that minimises pH changes on addition of small amounts of an acid or a base.

24
Q

Define solubility product Ksp.

A

This represents the conditions for equilibrium between a sparingly soluble product and its saturated solution.

25
Define entropy.
Entropy is a measure of the number of ways of arranging molecules and distributing energy.
26
What is Kw?
The ionic product of water. Kw = [H+][OH-] = 1 x 10^-14 mol2dm-6 at room temperature.
27
Give the template for calculating pH of a strong acid.
1. Write out an ionisation expression for the acid 2. Use molar ratios to deduce the [H+] 3. Substitute into pH = -log [H+]
28
Give the template for calculating pH of a weak acid.
1. Write out the Ka expression for the weak acid, Ka = ([H+][A-])/[HA] 2. Assume that [H+] is the same as [A-], Ka = [H+]^2/[HA] 3. Rearrange to make [H+] the subject, √(Ka x [HA]) = [H+] 4. Substitute into pH = -log [H+]
29
Give the template for calculating pH of a strong base.
1. Write out an ionisation expression for the base 2. Use molar ratios to deduce the [OH-] 3. Substitute [OH-] into Kw, Kw = [H+][OH-] 4. Rearrange to make [H+] the subject, 1 x 10-14 / [OH-] = [H+] 5. Substitute into pH = -log [H+]
30
Give the template for calculating the pH of a buffer solution.
1. Write out the Ka expression for the buffer, [A-] = salt, [HA] = weak acid Ka = ([H+][A-])/[HA] 2. Recalculate the concentrations of HA and A- in the new total volume of the buffer. C1 x V1 = C2 x V2 rearranges to give C2 = (C1 x V1 / V2 ) 3. Rearrange to make [H+] the subject, Ka x ([HA]/[A-]) = [H+] 4. Substitute in Ka and concentrations of A- and HA to calculate [H+] 5. Substitute into pH = -log [H+]
31
How to calculate Ksp?
Write Ksp expression e.g. Ksp = [Ag+][Cl-]. Use molar ratios to deduce conc of each ion and calculate Ksp and units using expression.
32
How is calculated value of Ksp used?
The databook/ calculated value for Ksp provides a benchmark to decide whether an ionic salt will dissolve and form a solution or a precipitate.
33
How is Ksp used - 3 conditions?
* If calculated value of Ksp is smaller than benchmark value solution will form. * If calculated value of Ksp is same as benchmark value saturated solution is formed * If calculated value of Ksp is larger than benchmark value a precipitate/ solid will form.
34
How would you calculate Enthalpy change of reaction from experimental data?
Q = mcΔT ΔH = (-Q/1000)/n
35
How to calculate total entropy?
∆total S = ∆sysS + ∆surrS or ∆total S = ∆sysS + (- ∆H/T)
36
How to calculate the entropy of a system.
∆sysS = S of products - S of reactants
37
How to calculate the entropy of surroundings.
∆surrS = - ∆H/T
38
How is the minimum temperature a reaction is feasible at calculated?
∆totalS = ∆sysS + (-∆H/T) 0 = ∆sysS + (-∆H/T) T = ΔH/ΔsysS
39
Summarise the greenhouse effect.
Visible and ultra-violet radiation from Sun reaches earth * The Earth absorbs some of the solar radiation - so Earth warms up. * The warm Earth emits infrared radiation. * Greenhouse gases include carbon dioxide and methane. They are present in the troposphere. * These gases absorb infra-red radiation. This increases the vibrational energy of the bonds. Molecules in the atmosphere collide. * The kinetic energy of some molecules increases as energy is transferred between molecules. * This raises the temperature of the atmosphere. * Greenhouse gases emit some of the absorbed radiation. * Some of this is also absorbed by the Earth. * Increased concentrations of greenhouse gases leads to an enhanced greenhouse effect.
40
Could a weak acid solution act as a buffer? Explain your answer
No. Although it will be able to react with small volumes of OH- ions added, it will not contain a reservoir of A- ions to react with additional H+ ions added to solution. There is no equilibrium present in a solution of weak acid alone that can shift position to minimise pH changes on addition of both small amounts of acid and alkali
41
Describe a practical technique that can be used to measure the enthalpy change of hydration of an ionic salt.
* Take a known volume of water (e.g. 100 cm3) in a polystyrene cup. * Record the temperature of the water. * Add a known mass of salt (e.g. 1.2 g) and stir to dissolve the salt. * Record the highest/lowest maximum temperature. * Calculate the temperature change, temp change = final temperature-initial temperature. * Use Q = mcΔT to calculate the energy transfer to the surroundings in J. * Calculate the number of moles of salt dissolved (n = m/Mr) * Use ΔH = (-Q/1000)/n to calculate the overall enthalpy change of hydration in kJmol-1.
42
What is the IR Window?
Wavelengths of IR radiation not absorbed by greenhouse gases.
43
Define alkali.
A base that dissolves in water to form OH- ions.
44
Explain what is meant by conjugate acid-base pair.
A pair of species which transform into each other by loss or gain of a proton.
45
Give the mathematical definition of pH.
pH = -log10[H+]
46
What does pH measure?
The concentration of H+ ions in a solution. The lower the pH, the higher the concentration of hydrogen ions in the solution.
47
What is the acid dissociation constant Ka?
The extent to which acids dissociate to donate H+. Units of mol dm-3. A large Ka = large extent of dissociation and strong acid. A small Ka = small extent of dissociation and weak acid.
48
What are the conversions associated with calculating pH?
pH = -log[H+] [H+] = 10^-pH pKa = -log(Ka) Ka = 10^-pKa
49
What are the assumptions for Ka?
[H+] = [A-] Amount of HA at equilibrium = Amount of HA put into the solution.
50
What is a buffer made of?
A mixture of a weak acid and its salt or a weak base (HA) and it’s conjugate base (A-).
51
How does a buffer work?
The weak acid, HA, removes most of any added alkali. The conjugate base, A- removes most of any acid.
52
What are the assumptions for a buffer?
All the A- ions come from the salt. Almost all the HA molecules put into the buffer solution remain unchanged.

A Level Chemistry OCR B (15 decks)

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