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AE 1 Flashcards

1
Q

Outline the steps to prepare and test a natural indicator

A
  1. Place shredded cabbage into a pestle.
  2. Add a large pinch of sand and 10 mL of ethanol into the pestle.
  3. Grind the mixture using a mortar until the ethanol has become intensely coloured.
  4. Decant the solution into a small beaker.
  5. Prepare 10-fold serial dilutions of 1 molL-1 HCl and 1 molL-1 NaOH in separate test tubes.
  6. Add three drops of the indicator solution into each test tube.
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2
Q

What are the variables of the practical:preparing and testing a natural indicator

A

Independent variable: pH of the solution in each test tube.
Dependent variable: Colour of the solution in each test tube after the indicator was put in.
Controlled variable: The amount of indicator used and the amount of HCl and NaOH in each test tube.

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3
Q

What are the results of the practical:preparing and testing a natural indicator

A

acidic ———————–>basic

red >purple> blue> green >yellow

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4
Q

State the range (number of colour) of methyl orange

A

Low ph to high
red to yellow
3.1 to 4.4

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5
Q

State the range (number of colour) of methyl red

A

Low ph to high
red to yellow
4.8-6.0

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6
Q

State the range (number of colour) of bromothymol blue

A

Low ph to high
yellow to blue
6.0-7.6

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7
Q

State the range (number of colour) of phenol red

A

Low ph to high
yellow to red
6.8-8.4

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8
Q

State the range (number of colour) of phenophthalein

A

Low ph to high
colourless to pink
8.2-10

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9
Q

State the range (number of colour) of litmus

A

Low ph to high
red to blue
5-8

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10
Q

Describe the everyday uses of indicators in sol testing

A

o A small sample of soil is mixed with universal indicator solution.
o BaSO4 is sprinkled onto the surface
o BaSO4 provides a neutral white background which contrasts the colour of the indicator solution seeping through from the soil.
o The colour is matched to a pH value on a pH chart.

Soil testing is important as pH of soil affects nutrient availability.
o In agriculture, this affects the productivity and yield.
o pH is determined to affect which plants can be grown as each plant is only suitable in certain conditions.

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11
Q

Describe the everyday uses of indicators in swimming pools

A

o Collect sample of pool water.
o Place drops of indicator in sample (as recommended).
o Use white background to match the colour chart.
o pH can be lowered or raised using substances such as HCl or Na2CO3.

o It is important to maintain pH of swimming pools close to 7.4 to avoid irritation of the eyes and mucous membranes.
o It is kept at 7.4 as it is more basic which prevents algae and bacteria from growing.

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12
Q

Describe the everyday uses of indicators in aquarium waters

A

o pH may be affected by respiration, decay of biological matter (produces acids) and carbonates leaching from decorations such as coral and gravel (raised pH).
o Certain species of fish may favour or can only prosper in certain pH ranges.
o Indicators can be used to ensure aquarium water lies in a suitable range.

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13
Q

What is the trend of basic oxides and acidic oxides in the periodic table

A

o Non-metals on the RHS of the periodic table are acidic and covalent.
o Metals on the LHS tend to form basic oxides.
 Generally the more EN the element, the more acidic its oxide.

o Oxides of elements near the semi-metal region of the periodic table can react with both acids and bases. These oxides are known as amphoteric oxides. Examples include: Al2O3, ZnO and Cr2O3.
o However, these substances do not contain ionisable protons (H+) and hence cannot be amphiprotic (substances which can donate or accept protons).

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14
Q

Define Lechats

A

When a chemical system at equilibrium is disturbed, the system will shift to a new equilibrium by undergoing a reaction which minimises the effect of the disturbance.

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15
Q

What is a dynamic equilbrium

A

o Dynamic Equilibrium: Microscopic changes which balance out to produce no overall net change at a macroscopic level.

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16
Q

Carbon dioxide dissolves in water and reacts to form carbonic acid, H2CO3.

CO2(aq) + H2O(l) ⇌ H2CO3(aq)

What happens when the bottle is opened for the first time

A

(decrease in pressure), therefore Increase in CO2 will be seen as many gas bubbles will rise to the surface (effervescence)
shifts left

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17
Q

Carbon dioxide dissolves in water and reacts to form carbonic acid, H2CO3.

CO2(aq) + H2O(l) ⇌ H2CO3(aq)

What happens when an open bottle eventually goes flat and is no longer fizzy.

A

Decrease in [CO2(g)]
CO2(g) escapes and hence bubble form and escapes
shifts left

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18
Q

Carbon dioxide dissolves in water and reacts to form carbonic acid, H2CO3.

CO2(aq) + H2O(l) ⇌ H2CO3(aq)

What happens when a cold bottle is placed on a desk and allowed to warm up.

A

increase in temperature shifts left, little bubbles produced

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19
Q

Carbon dioxide dissolves in water and reacts to form carbonic acid, H2CO3.

CO2(aq) + H2O(l) ⇌ H2CO3(aq)

What happens when the volume of CO2 is increased

A

shifts left
An increase in volume means there is a decrease in pressure. Thus the equilibrium will shift to the side with more gas moles to minimise the disturbance, thus favouring the reverse reaction.

20
Q

What are the natural sources of sulfur dioxide

A

Geothermal hot springs and volcanoes

Natural decay of vegetation land, wetlands and oceans.

21
Q

What are the industrial sources of sulfur dioxide

A

Burning of fossil fuels
S(s) + O2(g) → SO2(g)
Smelting of Metals
CuS(s) + O2(g) → Cu(s) + SO2(g)

22
Q

What are the natural sources of NOx

A

Lightning
O2 + N2(g) → 2NO(g) (needs condition >1000 OC)
2NO(g) + O2(g) → 2NO2(g) (reacts with air)

micro-organisms feed on nitrogenous material is soils to produce N2O

23
Q

What are the industrial sources of NOx

A

Combustion of fossil fuels in motor vehicle

Combustion in oil and coal electrical power plants

24
Q

What are two environmental effects of acid rain

A

Damage to forest (trees die from nutrient deficiency rather than direct)- bad for environment in general and can affect agricultural purposes.

increase in lake acidity (kills fish and hence disrupts food chains)

25
Q

What are two societal effects of acid rain

A

Erosion of irreplaceable marble and limestone buildings.
H2SO4(aq) + CaCO3(s) → CaSO4(s) + H2O(l) + CO2(g)

Photochemical smog- hazy, irritates eyes and hard to breathe

26
Q

What are the problems associated with SO2 and NOx

A

o Sulfur oxides are irritating gases which present a major health risk, especially for people with respiratory disorders.
o Nitrogen oxides are important components in formation of photochemical smog and acid rain.
o Sulfur dioxide, sulfur trioxide and nitrogen dioxide are acidic oxides which combine with water to produce acid rain.

27
Q

How can sulfur dioxide emissions be reduced

A

o Reducing sulfur content of fuels (OZ coal)
o Using a scrubber to remove sulfur dioxide
 Scrubber passes gaseous emissions through a slurry of lime (CaO), which neutralises the sulfur dioxide to form calcium sulfite

SO2(g) + CaO(s)  CaSO3(s)

28
Q

How can nitric oxide emissions be reduced

A

o Catalytic converters catalyse decomposition of NO back to N2¬ and O2, and also decompose carbon monoxide into carbon dioxide
o Rhodium or platinum catalyst

2NO(g) + 2CO(g)  N2(g) + 2CO2(g)

29
Q

State the equations in how acid rain is formed

A

2NO2(g) + H2O(l) → HNO3(aq) + HNO2(aq)

SO2(g) + H2O(l) → H2SO3(aq)

30
Q

what are the evidence that show an increase in atmospheric concentration of sulfur and nitrogen
and what are the problems with the evidence shown

A

o Direct: chromatographs show an increase since 1970s.
o Indirect: samples extracted from ice cores from 1700s
o Knowledge from NOx and SO2 from things such as a cars. An increase in cars results to an increase in sulfur and nitrogen. (industrial revolution)
o Measurements of effects of increased areas of acid rain and photochemical smog.

Problems with Evidence

o It only showed from 1970s and onwards.
o Evidence on Antarctica might not reflect levels worldwide or concentrated areas may only be generally geographically localised to a specific region.

31
Q

Outline the steps in the practical:use of pH meters/probes and indicators to distinguish between acidic, basic and neutral chemicals

A
  1. Rinse the pH probe with distilled water.
  2. Calibrate the pH probe using standard buffer solutions with pH 4, 7 and 10.
  3. Place a small amount of the solution to be tested into a beaker.
  4. Insert the pH probe into the substances and record the pH of the solution.
  5. Rinse the pH probe with distilled water and repeat steps 3, 4 and 5.
32
Q

What is the funky name for citric acid

A

2-hydroxypropane-1,2,3-tricarboxylic acid

33
Q

Identify examples of naturally occurring acids and their chemical composition

A

citric acid: (2-hydrohypropane-1,2,3-tricarbocylic acid)- C6H8O7: found in citric fruits such as lemon and orange.
Acetic acid: CH3COOH - Acetic acid is produced naturally through bacterial fermentation which is oxidised to acetic acid.

34
Q

Identify examples of naturally occurring bases and their chemical composition

A

Calcium carbonate: (CaCO3)- found in limestone.

Ammonia: (NH3) Certain organisms can produce ammonia from atmospheric nitrogen.

35
Q

What is the difference between strength and concentration of an acid

A

o Strength depends on the identity of the substance and its degree of ionisation in water.
o Concentration is the amount of acid in a given volume of a solution.

36
Q

How do you convert [H] to pH

A

-log (H)

37
Q

How do you convert pH to [H]

A

10^(-pH)

38
Q

How do you convert [H] to [OH]

A

[H] x [OH] = 10^-14

39
Q

How do you convert pH to pOH

A

pH + pOH = 14

40
Q

State two reasons in which acids are used as food preservatives.

A

o Used as opposed to bases because they taste better. E.g. lactic acid, acetic acid and citric acid.
o Used as a preservative- destroys microorganisms involved in decomposition; oxidation is limited and it cannot survive in low pH.

41
Q

Outline what Lavoisier (1776) proposed

A

o He concluded an acid must contain oxygen.

o He was wrong when HCl(l) was discovered to have no oxygen and reacted in water to produce acidic solutions.

42
Q

Outline what Davy (1815) proposed

A

o Noted that all known acids contained a hydrogen that could be displaced by a reaction with a metal.

43
Q

outline three things arrhenius (1884) proposed

A

o Defined acids as a molecule or a polyatomic ion with a H in its formula and dissociates in water to produce hydrogen ions, H+(aq).
o A base was defined as a molecule or ion with OH in its formula and dissolves in water to produce hydroxide ions, OH-(aq).
o He noted the strongest acids produced the strongest degree of dissociation.

44
Q

What are the limitations to arrhenius’ definition

A

o His definition however, could not explain the basic nature of substances such as ammonia which did not contain hydroxide ions in their formulae, but behaved like bases.
o He could not explain why some salts were basic or acidic.
o His definition only covered acids and bases in aqueous solution, hence could not explain the role of the solvent.

45
Q

what is BL propose

A

o Acids were proton donors and bases were proton acceptors.

o A proton was defined as a hydrogen ion.

46
Q

Outline how BL’s definition affected Arrhenius’s definition

A

o BL did not contradict Arrhenius’ definition, it only generalises it. I.e. it allowed for more species to be defined as an acid or a base.
o It could explain the basic nature of substances like ammonia. Ammonia dissolves in water and reacts with a water molecule to produce NH4+(aq) and a hydroxide OH-(aq).

47
Q

what is an amphiprotic substance

A

o A molecule that can act as a proton acceptor and a proton donor is amphiprotic

AE (1 decks)

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