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1
Q

What determines the solubility of a solute in solvents?

A

The strength of interactions between solute and solvent molecules with respect to intermolecular forces.

Stronger interactions lead to greater solubility.

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2
Q

What type of solute is soluble in polar solvents?

A

Polar molecules

Polar molecules exhibit dipole-dipole interactions.

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3
Q

What type of solute is soluble in nonpolar solvents?

A

Nonpolar molecules

Nonpolar molecules exhibit id-id forces.

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4
Q

Do polar and nonpolar molecules mix?

A

Generally, they do not mix

This is due to the differing intermolecular forces.

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5
Q

What is the solubility behavior of ionic compounds in polar solvents?

A

Ionic compounds are more soluble in polar solvents like water

Ions can form strong favorable interactions with polar solvents as ions are charged.

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6
Q

What does lattice enthalpy measure?

A

The strength of ionic bonding in a giant ionic lattice

It is specific to ionic substances.

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7
Q

Define lattice dissociation enthalpy.

A

The enthalpy change when one mole of a solid ionic compound is completely dissociated into its gaseous constituent ions under standard conditions.

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8
Q

What type of process is lattice dissociation enthalpy?

A

Endothermic process

This indicates that energy is absorbed during the formation of a solid ionic compound from its gaseous ions.

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9
Q

How is lattice formation enthalpy defined?

A

The enthalpy change when one mole of a solid ionic compound is formed from its gaseous constituent ions under standard conditions.

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10
Q

What type of process is lattice formation enthalpy?

A

Exothermic process

This indicates that energy is released when the lattice structure is formed.

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11
Q

Can lattice enthalpy be measured directly?

A

No, it is calculated using experimental values for other enthalpy changes.

This calculation often involves the Born-Haber cycle.

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12
Q

What are the two types of enthalpies discussed that can measure lattice enthalpies indirectly?

A

Enthalpies of Solution (AsolH) and Hydration (AhydH)

These enthalpies provide alternative methods to assess lattice enthalpies by using principles similar to Hess’s Law.

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13
Q

What is the enthalpy of solution?

A

The enthalpy change when one mole of ionic solid is dissolved in water to infinite dilution so that the ions no longer interact under standard conditions.

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14
Q

What is the enthalpy of hydration?

A

The enthalpy change when one mole of gaseous ions is dissolved in water to form one mole of aqueous ions under standard conditions.

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15
Q

Is the enthalpy of hydration typically positive or negative?

A

Negative

This is because water molecules have + regions that naturally attract negative ions.

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16
Q

What is a hydrated ion?

A

A solvated ion surrounded by solvent molecules that are aligned in a spherical shell.

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17
Q

How are the enthalpy of solution and the enthalpy of hydration related?

A

They are combined in an energy cycle.

This relationship is essential for calculating changes in enthalpy during dissolution processes.

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18
Q

What is the enthalpy of hydration?

A

A measure of the attraction between ions and water molecules in solution

Positive ions are attracted to the o- oxygen atoms, and negative ions are attracted to the o+ hydrogen atoms.

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19
Q

How does the size and charge of ions affect hydration enthalpy?

A

Attractions are stronger with smaller ions and ions of greater charge

This explains why hydration enthalpies decrease as you move down a group and why Na+ ions have a lower enthalpy of hydration than Mg2+ ions.

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20
Q

What happens to lattice dissociation enthalpy with decreasing ionic radius and increasing charge?

A

It increases in magnitude

Stronger attractions require greater energy for dissociation.

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21
Q

What does entropy (ΔS) measure?

A

Disorder and the dispersal of energy in a system

Greater disorder equates to greater entropy.

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22
Q

What is the relationship between temperature and entropy?

A

Entropy increases as temperature increases

Particles gain energy and move faster and further apart, becoming less ordered.

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23
Q

Which state of matter has the greatest entropy?

A

Gases

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24
Q

What happens to entropy when a substance melts or evaporates?

A

There is a sudden increase in entropy

This increase occurs because the transition from solid to liquid (fusion) or liquid to gas (vaporization) results in greater disorder.

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25
Which has a greater entropy change, vaporization or fusion?
Vaporization ## Footnote The entropy change of vaporization is much greater than that of fusion because a gas is much more disordered than a liquid or solid.
26
In which scenario does a sudden increase in entropy occur during a reaction?
When there are more moles of gaseous product than reactants ## Footnote This increase in the number of gaseous products leads to greater disorder.
27
What occurs to entropy when a lattice is dissolved in solution?
There is an increase in entropy ## Footnote The ions from the lattice become dissociated and can move freely in the solution, increasing disorder.
28
How is the overall entropy change for a reaction measured?
In J K⁻¹ mol⁻¹
29
What is the formula to calculate the entropy change for a reaction?
ΔS_total = ΣS_products - ΣS_reactants
30
What does a positive entropy change indicate about the reactants and products?
The products are more disordered than the reactants ## Footnote A positive change in entropy reflects an increase in disorder.
31
What does a negative entropy change for a reaction indicate?
The products are less disordered than the reactants.
32
What is the natural direction of change in entropy?
Towards a more disorderly state overall, which is a positive change in entropy.
33
Provide an example of a natural and spontaneous process that reflects entropy increase.
The natural and spontaneous spreading of gas throughout a room.
34
What does the solubility product constant, Ksp, indicate?
How much a compound dissociates in water.
35
To what type of compounds does the solubility product, Ksp, apply?
Saturated compounds.
36
What does a higher Ksp value indicate about a compound's solubility?
The more soluble the compound.
37
What happens when an ionic solid dissolves in water?
An equilibrium is established between the ions and the compound.
38
Fill in the blank: BaSO4(s) = Ba2+(aq) + SO4^2-(aq) is an example of _______.
Ionic solid dissolving in water.
39
What is indicated by the formula AS_total = AS_system + AS_surroundings?
The total change in entropy is the sum of the changes in the system and the surroundings.
40
What is the solubility product constant represented as?
Ksp = [Ba²⁺][SO₄²⁻] ## Footnote This formula indicates the concentrations of the ions in a saturated solution of the compound.
41
Why is BaSO₄(s) not included in the Ksp equation?
It is a solid in a heterogeneous equilibrium. ## Footnote Only the concentrations of aqueous ions are included in the equilibrium constant expression.
42
How are the concentrations of substances used in the Ksp equation?
In the same way as in other equilibrium constant equations. ## Footnote For example, if there were 2Ca²⁺ in the equilibrium, Ksp would contain [Ca²⁺]².
43
What are the units for the solubility product of BaSO₄?
mol² dm⁻⁶ ## Footnote Calculated as (mol dm⁻³) x (mol dm⁻³).
44
What are the units for the solubility product of Ca₃(PO₄)₂?
mol⁶ dm⁻¹⁵ ## Footnote Calculated as (mol dm⁻³)³ x (mol dm⁻³)².
45
Under what condition will the solubility product for a compound always be the same?
Under the same conditions if the solution is saturated. ## Footnote This means that the concentrations of the ions remain constant.
46
What happens when two compounds are mixed together regarding precipitation?
A precipitate will form if the ionic concentrations give a value greater than the solubility product. ## Footnote If the product is less than or equal to the Ksp, the solution remains unsaturated.
47
What do acid-base equilibria involve?
The transfer of protons between substances
48
How are substances classified as acids or bases?
Depending on their interaction with protons
49
Define a Bronsted-Lowry acid
A proton donor
50
Give an example of a Bronsted-Lowry acid
Ammonium ions (NH4+)
51
Define a Bronsted-Lowry base
A proton acceptor
52
Give an example of a Bronsted-Lowry base
Hydroxide ions (OH-)
53
What is a conjugate acid?
The species formed when a base accepts a proton
54
What is a conjugate base?
The species formed when an acid donates a proton
55
What do conjugate acids and conjugate bases form?
Conjugate acid-base pairs
56
Provide an example of a conjugate acid-base pair.
HCl (acid) and Cl- (conjugate base)
57
How is a strong acid defined?
An acid that completely dissociates into its ions when in solution.
58
What is the definition of a weak acid?
An acid that does not completely dissociate into its ions when in solution.
59
What is the pH range for strong acids?
0-1
60
What is the pH range for weak acids?
3-7
61
What is the pH range for strong bases?
12-14
62
What is the pH range for weak bases?
7-11
63
What are spectator ions?
Ions that do not change in the reaction and are left out of the ionic equation.
64
What is important to balance in ionic equations?
Elements and charge.
65
In the reactions of acids with carbonates, metal oxides, and alkalis, which ion is the reacting ion?
H+
66
What is the product of the reaction between 2H+ and CO3^2-?
H2O + CO2
67
Fill in the blank: The formula for the Acid Dissociation Constant is Ka = [H+][A-] / _______.
[HA]
68
What does Ka represent in the context of weak acids and bases?
The Acid Dissociation Constant
69
In the dissociation of weak acids, what type of mixture is formed?
An equilibrium mixture
70
What is the relationship between pKa and Ka?
pKa = -log(Ka) ## Footnote This equation allows the calculation of the acid dissociation constant (Ka) from the pKa value.
71
How can Ka be calculated from pKa?
Ka = 10^(-pKa) ## Footnote This formula indicates that Ka is derived from the negative logarithm of the pKa.
72
What does a low value of pKa indicate?
A strong acid ## Footnote A low pKa value corresponds to a high Ka, signifying a greater ability to donate protons.
73
Define pH.
A measure of acidity and alkalinity on a logarithmic scale from 0 to 14 ## Footnote pH indicates the concentration of H+ ions in a solution.
74
What does a pH of 0 represent?
An acidic solution with a high concentration of H+ ions ## Footnote A pH of 0 is the most acidic point on the pH scale.
75
What does a pH of 14 represent?
A basic solution with a low concentration of H+ ions ## Footnote A pH of 14 is the most alkaline point on the pH scale.
76
How is pH calculated from the concentration of hydrogen ions?
pH = -log[H+] ## Footnote This formula shows how to derive pH from the concentration of hydrogen ions in a solution.
77
What is the formula to find the concentration of hydrogen ions from pH?
[H+] = 10^(-pH) ## Footnote This equation allows you to calculate the hydrogen ion concentration based on the pH value.
78
What is the relationship between the concentration of H+ ions and strong acids?
The concentration of H+ ions is equivalent to the concentration of a strong acid as it completely dissociates to ions in solution.
79
What does the ionic product of water (Kw) represent?
Kw represents the equilibrium constant for the dissociation of water into hydroxide and hydrogen ions.
80
What is the equation for the ionic product of water?
Kw = [H+][OH-]
81
What is the value of Kw at 25°C?
At 25°C, Kw has a constant value of 1.0 x 10^-14.
82
How does temperature affect the value of Kw?
As temperature changes, the value of Kw changes.
83
What type of reaction is the forward reaction in the equilibrium of water?
The forward reaction is endothermic.
84
What happens to H+ ions as temperature increases?
As temperature increases, more H+ ions are produced, making the water more acidic.
85
What is the dissociation equation of water?
H₂O = H+ + OH-
86
How can pKa be calculated?
pKa can be calculated from Ka.
87
What is the formula for calculating pKw?
pKw = -log10kW
88
What is the relationship between Kw and concentration of OH for a strong base?
For a strong base, the concentration of OH will be the same as the concentration of the base.
89
What is the formula for Kw?
Kw = 10-pKw
90
How can [H*] be found from Kw?
[H*] = Kw / [OH].
91
What value is used for Kw in calculations?
Kw = 10^-14.
92
What relationships can be used to find the pH of weak acids and bases?
The relationships of Ka, pKa and [H*].
93
What method should be used when HA is in excess?
Use [HA] and [A] along with Ka to find [H*], then pH.
94
What method should be used when A is in excess?
Use Kw to find [H*], then pH.
95
What is the relationship between pKa and pH in this case?
pKa is equal to pH
96
What are the steps to calculate the Ka of weak acids?
1. Find initial concentrations, change in concentrations, and equilibrium concentrations 2. Find concentration of H+ ions using pH 3. Calculate actual equilibrium concentrations 4. Substitute values into the expression for Ka
97
What is a buffer solution?
A system that minimises pH changes on addition of small amounts of an acid or a base ## Footnote A buffer is formed from a weak acid and its salt or an excess of a weak acid and a strong alkali.
98
What is defined as a solution which is able to resist changes in pH when small volumes of acid or base are added?
Buffer solution ## Footnote Buffer solutions maintain a nearly constant pH despite additions of acids or bases.
99
What components typically form a buffer solution?
Weak acid and its salt or excess weak acid and strong alkali ## Footnote These components provide a reservoir of H+ and OH- ions.
100
What happens to the OH concentration in a buffer solution when a small amount of base is added?
It increases, making the solution more basic ## Footnote The extra OH ions react with NH4+ ions to form original reactants.
101
In the reaction NH3 + H2O = NH4+ + OH-, what occurs when OH- ions are added?
The equilibrium shifts to the left to remove OH- ions ## Footnote This reaction helps to stop significant changes in pH.
102
True or False: Buffer solutions are effective for both weak and strong acids.
False ## Footnote The approximations for buffer solutions are limited to weak acids.
103
What are the two types of buffer calculations?
1. Acid + Base 2. Acid + Salt
104
In acid + base buffer calculations, what is the first step?
Find the number of moles of each species
105
What is the second step in acid + base buffer calculations?
Calculate their concentrations when at equilibrium using the total volume
106
What is the final step in acid + base buffer calculations?
Use Ka to find [H+] and therefore pH
107
In acid + salt buffer calculations, what is the first step?
Find the moles of the salt
108
What is the second step in acid + salt buffer calculations?
Use Ka to find pH
109
What is the concentration of methanoic acid in the example buffer solution?
0.35 mol dm⁻³
110
What is the concentration of sodium methanoate in the example buffer solution?
0.67 mol dm⁻³
111
What is the value of Ka for methanoic acid in the example?
1.6 × 10⁻⁴ mol dm⁻³
112
What assumption is made about sodium methanoate in the buffer solution?
It completely dissociates
113
What assumption is made about methanoic acid in the buffer solution?
It only slightly dissociates
114
What is the formula for the acid dissociation constant (Ka)?
Ka = [H*] × [HCOO] / [HCOOH] ## Footnote This formula represents the equilibrium constant for the dissociation of a weak acid in solution.
115
How do you rearrange the expression for Ka to find [H*]?
[H*] = Ka × [HCOOH] / [HCOO] ## Footnote This rearrangement allows for the calculation of the hydrogen ion concentration from known values.
116
Calculate [H+] if Ka = 1.6 x 10^-4, [HCOOH] = 0.35, and [HCOO] = 0.67.
[H+] = 1.6 x 10^-4 x (0.35 / 0.67) = 8.4 × 10^-5 ## Footnote This calculation demonstrates the process of finding the hydrogen ion concentration in a buffer solution.
117
What is the pH when [H+] = 8.4 × 10^-5?
pH = -log10(8.4 × 10^-5) = 4.08 ## Footnote pH is calculated using the negative logarithm of the hydrogen ion concentration.
118
True or False: The pH of a buffer solution changes significantly when small volumes of acid or base are added.
False ## Footnote The pH of a buffer solution does not change much, typically in the order of 0.1 or 0.01 units.
119
What happens to the pH when small amounts of acid (H*) are added to a buffer solution?
The solution becomes slightly more acidic ## Footnote Adding acid increases the concentration of the acid component in the buffer.
120
What effect does adding small amounts of base (OH) have on the pH of a buffer solution?
The solution becomes slightly more basic ## Footnote Adding base decreases the concentration of the acid component in the buffer.
121
What are buffer solutions used for in nature?
To keep systems regulated ## Footnote This regulation is crucial for maintaining specific pH levels required for enzymatic reactions in living organisms.
122
How does the human circulatory system maintain pH?
Through a buffer between carbonic acid and hydrogencarbonate ions ## Footnote These ions neutralize acidic substances in the bloodstream, converting them to carbonic acid and water.
123
What is the pH range maintained by the buffer in human blood?
Between 7.35 and 7.45
124
What is the greenhouse effect?
The process of radiation from atmosphere usually UV and visible radiation warming the planet above temperatures that it would have reached without atmosphere. The planet then radiates often IR radiation outwards in all directions. Greenhouse gases such as CO2 and methane in the troposphere absorb some of this IR which increases their kinetic energy alongside their vibrational energy in their bonds which can be passed along to other molecules by collisions which increases kinetic energy. This causes a general increase of energy and therefore an increasing temperature. increasing greenhouse gases being admitted into the atmosphere leads to an increase in the greenhouse effect

chemistry (10 decks)

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