9.3 Acidic Environment Flashcards Preview

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Flashcards in 9.3 Acidic Environment Deck (65):
1

Identify 2 acidic, basic and neutral substances

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

2

Define: Indicator

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.

3

Identify the range in which phenolphthalein changes colour

Between neutral and basic, from clear to pink

4

Identify the range in which litmus changes colour

Between acidic and basic, from red to blue

5

Identify the range in which bromothymol blue changes colour

Between acidic and basic, from yellow to blue

6

Identify the range in which methyl organise changes colour

Between acidic and neutral, from red to yellow

7

Identify and describe three situations in which testing pH is necessary

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

8

Describe an experiment used to prepare a natural indicator

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.

9

Describe how litmus was discovered

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

10

Define acidic oxide and identify their properties

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

11

Define basic oxide and identify their properties

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

12

Define amphoteric oxides and identify their properties

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

13

Define neutral oxides and identify their propeties

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

14

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

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.

15

Quote Le Chateleir's principle

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

16

Define: equilibrium reaction

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.

17

Identify three factors which can disturb equilibrium

1. temperature
2. pressure
3. change in concentration

18

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

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.

19

Identify the natural and industrial sources of sulphur dioxide

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

20

Identify the natural and industrial sources of oxides of nitrogen

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

21

Recite a reaction for the production of sulphur dioxide

Zinc sulfide + oxygen -> Zinc oxide and sulfur dioxide

22

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

Oxygen + Nitrogen -> Nitric oxide
Nitric oxide + Oxygen -> Nitrogen dioxide

23

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

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.

24

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

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.

25

Identify the effects of acid rain

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

26

Identify the industrial origins of sulphur dioxide

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

27

Identify the industrial origins of nitrogen dioxide and nitric oxide

Combustion in power stations

28

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

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

29

Identify the environmental concerns of sulphur dioxide and oxides of nitrogen

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

30

Define the Bronsted-Lowry description of acids and bases

Acids are proton donors (lose hydrogen ions)
Bases are proton acceptors (gain hydrogen ions)

31

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

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

32

Define: Strong acid and give examples

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

Define: Weak acid

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

34

Define: concentrated/dilute acids

The amount/quantity of acid in a solution

35

Describe qualitatively what happens when the pH of a solution increases

The concentration of hydrogen ions decreases by a factor of 10

36

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

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

Identify whether strong acids create equilibrium reactions and explain why

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

Identify whether weak acids create equilibrium reactions and explain why

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

Identify the reasons why acids are frequently added to foods

Improve the taste (tartness)
Preservative (bacteria can't survive in acidic environments)

40

Identify three acids or bases as examples of naturally occurring substances

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

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

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

42

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

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

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

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

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

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

45

Define: Amphiprotic and give an example

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

46

Describe neutralisation reactions

Neutralisation reactions involve proton transfer. They are exothermic as new bonds are formed

47

Explain the purpose of titration experiments

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

48

Define: equivalence point (end point)

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

49

Describe the process of conducting a titration experiment

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

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

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

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

- 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

Describe the titration between a strong acid and a strong base

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

53

Describe the titration between a weak acid and strong base

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

54

Describe the titration between a strong acid and weak base

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

55

Define: Buffer

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

Describe the limitation of using a buffer

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

57

What principle do buffers rely on?

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

58

Identify and describe a naturally occurring buffer

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

Describe the need for neutralisation reactions and provide an example

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

60

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

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

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

- cost
- speed of action
- effect if excess is used
- safety in storing
- amphiprotic property

62

Identify the pH range of Bromothymol blue

6.2-7.6 (strong acid-strong base)

63

Identify the pH range of Methyl orange

3.1-4.4 (strong acid-weak base)

64

Identify the pH range of litmus

6-8 (strong acid-strong base)

65

Identify the pH range of phenolphthalein

8.3-10.0 (weak acid-strong base)