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Chemistry > Module 6 > Flashcards

Module 6 Flashcards

1
Q

What are the strong acids?

A 
HCl
HBr 
HI
HNO3
H2so4
HClO3 (chloric acid)
HClO4 (perchloric acid)
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2
Q

What are the strong bases?

A 
LiOH
NaOH
KOH
RbOH
CsOH
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
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3
Q

Acid + Base

A

HX + MOH -> MX + H2O

Acid + Base -> salt + water

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4
Q

Acid + Carbonate

A

HX + MCO3 -> MX + H2O + CO2

Acid + Carbonate -> salt + water + carbon dioxide

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5
Q

Acid + Metal

A

HX + M -> H2 + MX

Acid + Metal -> hydrogen gas + salt

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6
Q

Compare the Arrhenius and Bronsted-Lowry definition of an acid and base, outlining its limitations.

A

Both definitions state that an acid contains H+. However, the Arrhenius definition states that a base must contain OH-. The Arrhenius definition has many limitations including that it only accounts for aqueous acids and bases, it does not explain why Ammonia is a base, and it does not explain why some neutralisation reactions are not neutral. The Bronsted-Lowry definition improves from this as it defines an acid as a proton donor and a base as a proton acceptor.

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7
Q

What does a large Ka value indicate?

A

A large Ka value indicates that the acid is strong as it completely dissociates. A large Ka value will give a low pKa value.

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8
Q

What does a large Kb value indicate?

A

A large Kb value indicates that the base is strong as it completely dissociates. A large Kb value will give a low pKb value.

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9
Q

What is Kw?

A

The Ionic product of water, the value is constant.

Kw = [H3O+][OH-] = 10↑-14

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10
Q

What is the relationship between Ka and pKa?

A

pKa = -log↓10(Ka)

Because of the nature of log scales, a large Ka means a small pKa and a small Ka means the pKa will be large.

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11
Q

What is the relationship between Kb and pKb?

A

pKb = -log↓10(Kb)

Because of the nature of log scales, a large Kb means a small pKb and a small Kb means the pKb will be large.

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12
Q

pKa + pKb =

A

= pKw = 14

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13
Q

Ka x Kb =

A

= Kw

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14
Q

pH + pOH =

A

= 14

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15
Q

What is a buffer and how does it work?

A

A buffer is an aqueous solution that resists rapid change in pH when small amounts of acid or base are added. It is made up of small and equal amounts of a weak acid and its conjugate base, or a weak base and its conjugate acid. The equilibrium between these allows the buffer to work to resist changes in pH.

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16
Q

How to prepare a natural indicator

A
  1. Crush red cabbage using mortar and pestle
  2. Dissolve in ethanol, strain and filter to give you your indicator.
  3. Prepare solutions of known pH (0-14) to test indicator (this provides you with a colour change chart)

Red cabbage contains Anthocyanin pigment which allows it to be an indicator.

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17
Q

How does an indicator work?

A

Indicators are weak acid dyes that exist in equillibrium.
For example, Phenolphthalein works by the following equilibrium:
(Changes colour pH - 10)
X + OH- ⇌ XOH
Colourless Magenta/Pink

If you add OH- ⇌ shifts right (pink)
If you add H+ it reacts with OH- –> ↓[OH-] so ⇌ shifts left (colourless)

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18
Q

Why was the enthalpy of neutralisation less than expected?

A

This was due to the significant heat loss in the experiment. Since a lot of heat energy is being lost to the environment, it is not absorbed by the water, thus the end temperature is lower than it should be, and a lower energy released in the experiment is recorded.

*remember when calculating enthalpy of neutralisation remember it is exothermic.

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19
Q

What is an example of a neutralisation reaction in everyday life? Explain how it works.

A

A neutralisation reaction in everyday life includes the treatment of an insect bite. Most insect bites contain formic acid, and can therefore be treated by sodium hydrogen carbonate (or bicarb soda) which will neutralise the acid from the sting, and is not corrosive, so will not cause damage to the skin.

20
Q

What is an example of a neutralisation reaction in industry? Explain how it works.

A

Limestone, a basic salt is used to neutralise acidic waste. It works by the following equation:

2H+ (aq) + CaCO2(s) –> Ca2+(aq) + H20(l) +CO2(g)

21
Q

Describe a model you could use to distinguish between a strong and weak acid. What are its advantages and disadvantages?

A

Model of a Chocolate bar vs a lolly

HX + H2O –> H3O+ + X- (Chocolate doesn’t rewrap, therefor represents strong acid)

HX + H2O ⇌ H3O+ + X- (lolly rewraps, therefore represents weak acid)

Advantages of the model:

  • Visualises concept/ makes it tangible
  • Simplifies a complex idea for better understanding
  • Represents key components/processes
  • Allows for predictions
  • The lolly and chocolate have equivalent parts, the sweet is the anion and the wrapper is the proton
  • The parts are separate and can be separated (to represent ionisation in water)

Disadvantages:
- Can be oversimplified/lack detail

22
Q

How do you prepare and test a buffer?

A

Control - Water
Dependent - Weak acid buffer and Weak base buffer

  1. Add 10 mL of a weak acid/base and 10 mL of its conjugate in equal amounts to a beaker. Use a measuring cylinder to measure. This will create an equilibrium as shown below:

CH3COOH + H2O ⇌ CH3COO- + H3O+
CH3COO- + H2O ⇌ CH3COOH + OH-

  1. Test the buffer by adding drop by drop a strong acid/base and recording the change in pH from the original pH.
23
Q

Identify a buffer system that occurs in nature and explain its importance.

A

A buffer system that occurs in nature is human blood. The human body needs to maintain a certain pH range for optimal performance of its enzymes. Enzymes help catalyse all the biochemical reactions in the human body, so when pH varies extremely, the enzymes will cease to work. Without this catalyst, the reaction rate of the body’s metabolism will slow down eventually leading to death.
This buffer system works according to the following equations:
H2O + CO2 ⇌ H2CO3(aq) ⇌ H+(aq) + HCO3-(aq) ⇌ H+(aq) + CO3 2-(aq)

24
Q

Methyl Orange

A

Use for titrations between a strong acid and a weak base (acidic salts) pH range: 3.1 - 4.4 Red > Yellow

25
Bromothymol Blue
Use for titrations between a strong acid and a strong base (neutral salts) pH range: 6.0 - 7.6 Yellow > Blue
26
Phenolphthalein
Use for titrations between a weak acid and a strong base (basic salts) pH range: 8.2 - 10.0 Colourless > Magenta
27
Describe an investigation that demonstrates the use of pH to indicate the differences between the strength of acids and bases.
- Calculate the theoretical pH of solutions of a known concentrations - Use a pH meter to test the pH of these solutions and compare this to the theoretical pH Account for this difference: - The solution with lower or higher pH than reality has this pH because the solution has not fully ionised.
28
Examples of advantages and disadvantages of a model?
Advantages: - helps understanding as major components are represented - Accurate representation of the process due to choice of material Disadvantages: - Not to scale - Some components not represented/incorrect representation
29
Why must the conical flask be rinsed with distilled water?
This ensures all the molecules come into contact (so none are stuck on the side of the conical flask), and in doing so ensures that the equivalence point is correctly measured.
30
Why must the burette and pipette be rinsed with the solutions they will contain?
The burette must be rinsed with the titre and the pipette with the aliquot, as if they are cleaned with water it results in residual water molecules and other impurities, causing the solutions to become dilute, giving incorrect calculations of the concentration of the unknown solution.
31
What are the features of an appropriate primary standard solution? (there are 5)
- Can be readily obtained in a chemically stable, solid form - It must have a well-known chemical formula - Highly soluble in water - Easy to store, not hygroscopic (absorb water and change the chemical formula of the substance which causes molar mass changes) - High molar mass, to minimise % errors in weighing by reducing the proportion of the error of the solute weighed
32
What improvements could be made to the enthalpy of neutralisation experiment to improve accuracy?
- Constructing the calorimeter with two styrofoam cups, with an airspace in between. This increases insulation - Significant heat loss occurs through the lid, due to the radiation of heat, through the styrofoam and the surrounding air inside the cup, this could be improved through the use of an insulated stopper instead of a cardboard/paper lid - Not touching the cup as this transfers body heat, causing an inaccuracy in measurement, as extra heat from the environment is being added. - Stirring using a stirrer through a hole in the insulating device. This will help make sure all the moles have been reacted and prevent heat loss.
33
Describe how you would perform an experiment to find the molar enthalpy of neutralisation in a neutralisation reaction.
1. Set up your calorimeter with one/two foam cups, without any mixture or lid. 2. Measure HCl and NaCl (for example) according to their stoichiometric ratios. 3. Measure the temperature of each solution, ensuring they are both the same. 4. Pour both chemicals simultaneously into the calorimeter. 5. Place the lid with the thermometer already poked through, over the reaction mixture. Leave on the bench and don't touch. 6. Stir using the stirrer, without touching the cup. 6. Observe the temperature change on the thermometer. Record the highest temperature reached. (Look at eye level)
34
Describe a method you would use to create a standard solution and perform a titration.
1. Calculate the number of grams of solute you would need to create a standard solution of a certain concentration. 2. Accurately weigh this amount. 3. Transfer this to a 500mL volumetric flask (volume depends on conc this is an example) 4. Add distilled water, place the stopper in and swirl the solution around to dissolve all the solid solvent. 5. Add distilled water to the flask until the bottom of the meniscus is level with the 500mL mark. 6. Rinse the pipette with the standard solution you have created. 7. Add 25mL of this solution to a conical flask that has been rinsed with distilled water. 8. Add a few drops of indicator to the flask. 9. Rinse a burette with your solution of unknown concentration. Add your solution of unknown concentration until the meniscus is at the 0 mark of the burette. 10. Titrate the acid against the Na2CO3 solution until the indicator starts to become colourless and then continue drop by drop until it is completely colourless. 11. Record result and repeat titration at least 3 times to improve reliability. 12. Eliminate any outliers and then calculate the average volume. Use this to calculate the unknown concentration.
35
Why is Sodium Hydrogen Carbonate an amphiprotic salt?
When sodium bicarbonate dissociates in water, it becomes a sodium ion (Na+) and a bicarbonate ion (HCO3−). The bicarbonate ion itself is what is amphiprotic, and this is because it can lose the hydrogen it has to become a carbonate ion (CO32−), or it can gain a hydrogen to become carbonic acid (H2CO3). Seen below: HCO3 - + H3O+ --> H2CO3 + H2O HCO3 - + OH- --> CO3 2- + H2O
36
What is an example of an acid/base analysis technique used in industry?
Many industrially produced fertilisers are the products of neutralisation reactions. • These reaction vessels can be monitored using industrial pH probes as a method of knowing when reactions are complete or which reagent needs increasing. • Ammonium nitrate is reacted with nitric acid in a neutraliser. Here the pH is measured to ensure that the correct mixture has been achieved. • Ammonium nitrate has a pH of 5.3 as HNO3 is a strong acid.
37
What is pH
-log[H+]
38
What is pOH
-log[OH-]
39
Write an ionic equation to represent the dissociation of sodium hydrogen carbonate, and its amphiprotic nature.
Amphiprotic salt - NaHCO3(s) --> Na+(aq) + HCO3–(aq) HCO–3(aq) + H2O(l) ⇌ H2CO3(aq) + OH–(aq) Base Acid conjugate acid conjugate base HCO–3(aq) + H2O(l) ⇌ CO3 2–(aq) + H3O+(aq) Acid Base conjugate base conjugate acid
40
Write an ionic equation to represent the dissociation of potassium dihydrogen phosphate.
Amphiprotic salt - KH2PO4(s) ⇌ K+(aq) + H2PO4–(aq) Amphiprotic ion H2PO4–(aq) + H2O(l) ⇌ HPO4 2–(aq) + H3O+(aq) Acid Base Conj base Conj acid H2PO4–(aq) + H2O(l) ⇌ H3PO4(aq) + OH–(aq) Acid Acid Conj base Conj base
41
draw titration curve and conductivity graph of strong acid / strong base
check
42
draw titration curve and conductivity graph of strong acid / weak base
check
43
draw titration curve and conductivity graph of weak acid / strong base
check
44
Describe an example of an acid/base analysis technique used in industry.
* Many industrially produced fertilisers are the products of neutralisation reactions. * These reaction vessels can be monitored using industrial pH probes as a method of knowing when reactions are complete or which reagent needs increasing. * Ammonium nitrate is reacted with nitric acid in a neutraliser. Here the pH is measured to ensure that indeed the correct mixture has been achieved. * Ammonium nitrate has a pH of 5.3 as HNO3 is a strong acid.
45
Describe an example of an acid/base analysis technique used by Aboriginal and Torres Strait Islander Peoples.
• Bracken fern (Pteridium esculentum) has alkaline juices in the leaves when crushed. • Bull ants inject painful formic acid as it stings. • The alkali juices of the bracken fern have been known to neutralise the formic acid, seen in the equation below: HCOOH(aq) + BOH(aq) --> BHCOO(aq) + H2O(l)
46
Describe how you conducted a chemical analysis on a common household substance for its acidity or basicity.
Titrated white wine (burette) with standardised Na2CO3 (conical flask). Issues with the experiment: - There were minimum 3 acids in the wine (tartaric, citric, mallic and lactic) that weren't separated. Only [H+] conc could be determined, so the concentrations of each acid were unknown. - The colour of the wine complicated after the end point, therefore a pH meter could be used. - There was not enough of the wine sample to neutralise the Na2CO3

Chemistry (7 decks)

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