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HSC Chemistry > 9.3 Acidic Environment > Flashcards

9.3 Acidic Environment Flashcards

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1
Q

Identify 2 acidic, basic and neutral substances

A

Acidic: vinger, orange juice, tomato juice
Basic: soap, floor/drain cleaner, bleach
Neutral: salt, sugar, shampoo, milk

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2
Q

Define: Indicator

A

A substance than in solution changes colour, depending on the acidity or basicity or the solution. Each indicator changes into different colours at different pH values.

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3
Q

Identify the range in which phenolphthalein changes colour

A

Between neutral and basic, from clear to pink

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4
Q

Identify the range in which litmus changes colour

A

Between acidic and basic, from red to blue

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5
Q

Identify the range in which bromothymol blue changes colour

A

Between acidic and basic, from yellow to blue

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6
Q

Identify the range in which methyl organise changes colour

A

Between acidic and neutral, from red to yellow

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7
Q

Identify and describe three situations in which testing pH is necessary

A

Soil pH - different plant varieties require different soil pH
Pool water pH - needs to be neutral but cleaning solutions change the pH
Sewerage pH - effluents from factories must be a certain pH

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8
Q

Describe an experiment used to prepare a natural indicator

A

Red cabbage was used as a natural indicator by shredding and crushing the cabbage until a purple liquid was produced. In an acidic solution, this substance turned from purple to pink/red. In a basic solution, the substance turned from blue/green to yellow.

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9
Q

Describe how litmus was discovered

A

Litmus was discovered in the 18th century as it was observed that the colour of litmus changed in different solutions. It was later determined that this was due to the acidity/basicity of different substances. It is a dye that is extracted form lichens

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10
Q

Define acidic oxide and identify their properties

A

Acidic oxides are formed between oxygen and a non-metal. Acidic oxides react with water to form an acid to react with bases to form salts. Carbon dioxide is an acidic oxide

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11
Q

Define basic oxide and identify their properties

A

Basic oxides are formed between oxygen and a metal. Basic oxides react with acids to form salts and do not react with alkalis (soluble bases). Calcium oxide is a basic oxide

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12
Q

Define amphoteric oxides and identify their properties

A

Amphoteric oxides react with acids and bases to form salts. Amphoteric oxides used Zn, Al, Sn and Pb.

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13
Q

Define neutral oxides and identify their propeties

A

Neutral oxides do not react with acids or bases. Neutral oxides are CO and NO

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14
Q

Explain the relationship between the position of elements on the Periodic Table and the acidity/basicity of oxides formed

A

Acidic oxides are formed from non-metals. They are covalent molecules are are thus found on the right of the Periodic Table.
Basic oxides are formed from metals. they are ionic compounds and are thus found on the left of the Periodic Table.

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15
Q

Quote Le Chateleir’s principle

A

“If a system at equilibrium is disturbed, then the system adjusts itself so as to minimise the disturbance.”

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16
Q

Define: equilibrium reaction

A

An equilibrium reaction is a reaction which has the capacity to ‘go both ways’. They are able to go in the forwards direction and produce produce, or go in the reverse and produce the reactants.
A equilibrium, the forwards and backwards reaction occur at the same rate so that there is equal concentration of all substances.

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17
Q

Identify three factors which can disturb equilibrium

A
  1. temperature
  2. pressure
  3. change in concentration
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18
Q

Identify and describe how three factors affect the solubility of carbon dioxide

A

The conversion of carbon dioxide and water produces carbonic acid (and heat).

  • Pressure of CO2 is increased: solubility will be increased as system absorbs reactant
  • Total pressure increased: solubility increases
  • Temperature increases: solubility decreases as system absorbs heat.
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19
Q

Identify the natural and industrial sources of sulphur dioxide

A

Natural: geothermal hot springs and volcanoes
Industrial: burning of fossil fuels and extracting metals from sulfide ores

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20
Q

Identify the natural and industrial sources of oxides of nitrogen

A

Nitric oxide: lightning, combustion
Nitrogen dioxide: lightning, combustion
Nitrous oxide: bacteria in nitrogenous soil, nitrogenous fertiliser

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21
Q

Recite a reaction for the production of sulphur dioxide

A

Zinc sulfide + oxygen -> Zinc oxide and sulfur dioxide

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22
Q

Recite the reactions for the production of nitric oxide and nitrogen dioxide

A

Oxygen + Nitrogen -> Nitric oxide

Nitric oxide + Oxygen -> Nitrogen dioxide

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23
Q

Explain how the concentration of sulphur dioxide and oxides of nitrogen have increased

A

Following the Industrial Revolution, the production of sulphur dioxide increased as a waste product.
By the twentieth century, oxides of nitrogen began to be produced as they required great energy to be formed, primarily from the increased use of motor vehicles and electricity.
The currently annual emissions in 0.01ppm (10x greater than clean air), however mining regions often exceed this.

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24
Q

Recite Avogadro’s law in relation to the volume of gases

A

When measured at the same temperature and pressure, equal volumes of gases contain the same number of molecules.
therefore equal numbers of moles of different gases occupy the same volume.

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25
Q

Identify the effects of acid rain

A

Increase acidity of lakes
Damage to pine forests in Europe and North America
Erosion to marble and limestone
Damage to vegetation

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26
Q

Identify the industrial origins of sulphur dioxide

A

Smelting of sulfide ores, burning of coal (primarily contains sulfide) and refinement of crude oil

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27
Q

Identify the industrial origins of nitrogen dioxide and nitric oxide

A

Combustion in power stations

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28
Q

Identify the health concerns of Sulphur dioxide and the oxides of nitrogen

A

Sulfur dioxide can irritate the respiratory system and causes breathing difficulties. Magnified when small particles are present
Nitrogen dioxide and nitric oxide irritate the respiratory tract and cause breathing difficulties, at high concentration can cause tissue damage

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29
Q

Identify the environmental concerns of sulphur dioxide and oxides of nitrogen

A

When in the atmosphere can lead to the production of acid rain. Nitrogen dioxide also can produce photochemical smog

30
Q

Define the Bronsted-Lowry description of acids and bases

A

Acids are proton donors (lose hydrogen ions)

Bases are proton acceptors (gain hydrogen ions)

31
Q

List the four main acids studied as define them as strong or weak

A

Citric (2-hydroxypropane-1,2,3-tricarboxylic) - weak
Acetic (ethanoic) - weak
Hydrochloric - strong
Nitric - strong

32
Q

Define: Strong acid and give examples

A

when all the acid present in a solution completely ionise to produce hydrogen ions. Not neutral acid molecules are left. HCl, sulphuric acid, nitric acid

33
Q

Define: Weak acid

A

when only some of the acid present in solution has ionised to produce hydrogen ion. Is written as an equilibrium reaction.

34
Q

Define: concentrated/dilute acids

A

The amount/quantity of acid in a solution

35
Q

Describe qualitatively what happens when the pH of a solution increases

A

The concentration of hydrogen ions decreases by a factor of 10

36
Q

Compare the relative strengths of equal concentration of citric, acetic and hydrochloric acids in terms of the degree of ionisation

A

When the three acids are in equal concentration and compared, they will have different pH values. As pH refers to the concentration of hydrogen ions, the value reflects the degree of ionisation. Hydrochloric acids has the lowest pH value as it is a strong acid, meaning that all of the HCl has disassociated. Citric acid is a stronger weak acid than acetic as more of the hydrogen ions has disassociated. The degree of ionisation refers to the fraction of molecules that ionise in solution

37
Q

Identify whether strong acids create equilibrium reactions and explain why

A

Strong acids do not form equilibrium reactions as all the acid molecules completely ionise in solution. This means that the ions will never go in the reverse to recreate the acid molecule

38
Q

Identify whether weak acids create equilibrium reactions and explain why

A

Weak acids do form equilibrium reactions as not all of the acid molecules completely ionise. This means that it does not go to completion.

39
Q

Identify the reasons why acids are frequently added to foods

A

Improve the taste (tartness)

Preservative (bacteria can’t survive in acidic environments)

40
Q

Identify three acids or bases as examples of naturally occurring substances

A

Ammonia: produced during deamination in the body or anaerobic decay of organisms matter, used as a solvent
Amines: compound bound with another group (alkyl) and is formed during anaerobic decomposition of organic matter
Carbonates: found in limestone as calcium carbonate

41
Q

Describe how to find the hydrogen ion concentration of a substance when reacting with a known acid

A

Monoprotic: equal to conc of acid
Diprotic: twice conc of acid
Triprotic: trice conc of acid

42
Q

Outline Lavoisier’s, Davy’s, Arrhenius’ and Bronsted-Lowry’s definition of acids and Bases

A

Lavoisier: acids are substances that contain oxygen
Davy: acid are substances than contain replaceable hydrogen
Arrhenius: acids are substances the produce hydrogen ions in solution. Bases are substances that produce hydroxide ions in solution. Strong and weak acids
Bronsted-Lowry: acids are proton donors, bases are proton acceptors

43
Q

What is the major difference between Arrhenius’ and Bronsted-Lowry’s definition of acids and bases?

A

Arrhenius’ defintion relies on acids and bases interacting in solution
Bronsted-Lowry’s definition does not require the reaction to take place in solution

44
Q

Explain why some salts form solutions that have a pH other than 7

A

When salts enter solution, their anions or cation can act as acids or bases with water

45
Q

Define: Amphiprotic and give an example

A

A substance that can act as either an acid or a base (proton donor or proton acceptor). Hydrogen carbonate ions

46
Q

Describe neutralisation reactions

A

Neutralisation reactions occur between an acid and base and involve proton transfer. They are exothermic as new bonds are formed

47
Q

Explain the purpose of titration experiments

A

To find the concentration of a substance, using another substance of known concentration.

48
Q

Define: equivalence point (end point)

A

The point at which the amounts of the two reactants are just sufficient to cause complete consumption of both reactants.

49
Q

Describe the process of conducting a titration experiment

A
  1. Fill a burette of known concentration and record the starting volume
  2. Pipette 25mL of a solution of unknown concentration into a conical flask
  3. Place under burette
  4. Place 2-3 drops of indicator into conceal flask
  5. Run the titrant into the conical flask slowly until the indicator changes colour
  6. Repeat until an accurate volume of titrant it determined
50
Q

Define: Primary standard and identify substances that do not qualify as primary standards

A

A primary standard is a substance of high purity and stability that a solution of known concentration can be made. It must be non-hygroscopic and have an accurately known molar mass. HCl, sulphuric, nitric acid, Na and K hydroxide can not be used as a primary standard solution

51
Q

Describe the variables that must be controlled in order to make a standard solution

A
  • primary standard used is as pure as possible
  • substance used is placed in oven to evaporate
  • substance wieget accurately
  • all solute transferred to volumetric flask
  • all solute is completely dissolved
  • distilled water is used for cleaning
  • burette and pipette are rinsed with substance being measured
52
Q

Describe the titration between a strong acid and a strong base

A

Equivalence point is a pH 7 because salt produced is neutral. Possible indicator - Bromothymol Blue

53
Q

Describe the titration between a weak acid and strong base

A

Equivalence point is >7 as the salt produced is basic. Possible indicator: Phenolphthalein

54
Q

Describe the titration between a strong acid and weak base

A

Equivalence point is <7 as the salt produced is acidic. Possible indicator: methyl orange

55
Q

Define: Buffer

A

Solutions that a used to maintain a pH range. They resist rapid change in pH and contain comparable amounts of weak acid/conjugate base or weak base/conjugate acid

56
Q

Describe the limitation of using a buffer

A

Only work to a certain extent as at some point, the ions will become depleted.

57
Q

What principle do buffers rely on?

A

Le Chatelier’s principle. When acidic or basic substances are introduced the system will work to minimise the change in pH

58
Q

Identify and describe a naturally occurring buffer

A

In rivers or lakes, a buffer maintains the water at approximately pH 7. A weak acid, carbonic acid (from the reaction between the water and carbon dioxide) and sodium hydrogen carbonate ions (from dissolving rocks and soil) create a buffer

59
Q

Describe the need for neutralisation reactions and provide an example

A

Neutralisation reactions produce a product of neutral pH. Sewerage monitoring bodies have restrictions on chemical effluents. Neutralisation reactions ensure the standards are met.

60
Q

Give and example of a substance used in neutralisation reactions and identify reasons for tis use

A

Sodium hydrogen carbonate

  • solid to easy and safe to store
  • minimal damage if excess is used
  • amphiprotic, so can be used for acids and bases
61
Q

identify some factors that need to be considered when choosing a neutralisation substance

A
  • cost
  • speed of action
  • effect if excess is used
  • safety in storing
  • amphiprotic property
62
Q

Identify the pH range of Bromothymol blue

A

6.2-7.6 (strong acid-strong base)

63
Q

Identify the pH range of Methyl orange

A

3.1-4.4 (strong acid-weak base)

64
Q

Identify the pH range of litmus

A

6-8 (strong acid-strong base)

65
Q

Identify the pH range of phenolphthalein

A

8.3-10.0 (weak acid-strong base)

66
Q

Describe the differences between alkanols and alkanoic acids

A

Alkanols: contain a hydroxyl group. Have dipole-dipole bonds, hydrogen bonds and dispersion forces. Has a lower MP and BP than relative alkanoic acids due to lower molecule mass = weaker dispersion forces.
Alkanoic acids: contains a carboxylic group. Has dipole-dipole bonds, hydrogen bonds and dispersion forces.

67
Q

Describe the solubility of different esters

A

Smaller esters are highly soluble as they have a larger polar region in comparison to their non-polar region. Larger esters are insoluble as they have greater non-polar regions.

68
Q

Identify the two purposes of a catalyst in esterification and identify the catalyst used.

A

Sulfuric acid is used in esterification as a catalyst and a dehydrating agent. It absorbs the water product in the reaction, forcing the position of equilibrium to move to the right, resulting in an increase in the yield of ester, according to Le Chatelier’s principle.

69
Q

Explain the need for refluxing in esterification

A

Due to the heat applied to esterification, some liquid substances can become gases. As the substances used are volatile, this can cause a significant safety hazard if released as they are easily combustible. A condenser tube is to limit the volume of gaseous substances form escaping the reaction. As gases enter the tube, the cool water which surrounds the tube cools the case, causing it to condense and precipitation back into the reaction. Refluxing also decreases the pressure in the vessel by providing gases with a space to expand.

70
Q

Identify where esters occur

A

Esters occur in nature as flavouring and scents in plants and fruits.

71
Q

Identify where esters are produced

A

Esters are produced by extraction from natural sources or artificial production

72
Q

Identify uses of esters

A

Esters are used in nail polish remover, food flavouring and perfume.

HSC Chemistry (4 decks)

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