O LVL Chemistry > Practical > Flashcards

Practical Flashcards

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1
Q

Give an observation made when zinc reacts with aqueous copper (II) sulfate and explain the observation [2018]

A

observation- the blue colour of the copper (II) sulfate solution turns colourless and a red-brown solid forms
explanation- zinc is more reactive than copper, hence zinc will displace copper from its salt solution to form colourless zinc sulfate and red-brown copper solid

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2
Q

what is the end point of titration

A

when reaction is complete, and a colour change is seen

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3
Q

state the two types of titration

A

acid base titration, redox titration

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4
Q

state the function of a pipette

A

used to measure a very exact volume of liquid (exactly 25.0cm³)

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5
Q

state the function of a burette

A

graduated in 0.10cm³ and measures up to 50.00cm³ of solution. solution can be added 1 drop (approximately about 0.05cm³) at a time by controlling the tap of burette

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6
Q

how do you clean a burette

A

rinse the whole burette (including the tip area with distilled water followed by the solution before filling it up

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7
Q

how do you add the solution into the burette

A

use a filter funnel to fill up the burette and remove it after use

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8
Q

where do you ensure there is no air bubble when using burette

A

at the tip of the burette

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9
Q

how do you read a burette

A

at the bottom of the meniscus at eye level

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10
Q

record burette readings in what decimal place

A

2dp

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11
Q

what is the function of the shape of the conical flask

A

the shape of the conical flask allows the contents to be mixed by swirling without spillage

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12
Q

how do you clean the conical flask in titration

A

only rinse with distilled water before use. you do not need to dry it before using

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13
Q

what is the purpose of the white tile that is placed below the flask under the burette

A

to clearly observe the colour change of the solution

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14
Q

while adding the solution from the burette into the conical flask, what do you do with the conical flask?

A

SWIRL IT

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15
Q

how is the burette clamped

A

vertically upright

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16
Q

what is the consistent readings range

A

within ±0.20cm³

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17
Q

why does the conical flask need to be swirled constantly

A

to mix the reactants thoroughly to ensure reaction

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18
Q

state the observations for the following test:

Put one spatula of calcium carbonate powder in a clean test-tube. Add about 2 cm³ of dilute hydrochloric acid into the test-tube. Test the gas given off using limewater.

A

-effervescence was observed
-white precipitate formed in limewater
-gas is carbon dioxide

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19
Q

state the observations for the following test:

Put one piece of magnesium ribbon in a clean test-tube. Add about 2cm³ of dilute hydrochloric acid into the test-tube. Test the gas given off by placing a lighted splint at the mouth of the test-tube.

A

-effervescence was observed
-lighted splint extinguished with a “pop” sound
-gas is hydrogen gas

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20
Q

state the observations for the following test:

Put one spatula of copper(Il) oxide powder in a clean test-tube. Add about 2 cm³ of dilute hydrochloric acid into the test-tube and warm the mixture gently.

A

colourless solution turns blue

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21
Q

state the observations for the following test:

Put one spatula of ammonium chloride in a clean test-tube. Add about 2 cm³ of aqueous sodium hydroxide into the test-tube and warm gently.
Test the gas given off by holding moist red litmus paper near the mouth of the test tube.

A

-Effervescence was observed
-Most red litmus paper turns blue
-The gas is ammonia gas

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22
Q

define the term acid

A

an acid is a substance that dissociates/ionises in water to produce hydrogen ions

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23
Q

write a balanced chemical equation for the reaction in the following test:

Put one spatula of calcium carbonate powder in a clean test-tube. Add about 2 cm³ of dilute hydrochloric acid into the test-tube. Test the gas given off using limewater.

A

CaCO3 + HCl -> CaCl2 + CO2 + H2O

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24
Q

write a balanced chemical equation for the reaction in the following test:

Put one piece of magnesium ribbon in a clean test-tube. Add about 2cm³ of dilute hydrochloric acid into the test-tube. Test the gas given off by placing a lighted splint at the mouth of the test-tube.

A

Mg + 2HCl -> MgCl + H2

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25
write a balanced chemical equation for the reaction in the following test: Put one spatula of copper(Il) oxide powder in a clean test-tube. Add about 2 cm³ of dilute hydrochloric acid into the test-tube and warm the mixture gently.
CuO2 + 2HCl -> CuCl2 + H2O
26
write a balanced chemical equation for the reaction in the following test: Put one spatula of ammonium chloride in a clean test-tube. Add about 2 cm³ of aqueous sodium hydroxide into the test-tube and warm gently. Test the gas given off by holding moist red litmus paper near the mouth of the test tube.
NH4Cl + NaOH -> NH3 + NaCl + H2O
27
Another student repeated the teacher's demonstration using dilute ethanoic acid of the same concentration instead of dilute hydrochloric acid. Predict the effect of using dilute ethanoic acid on the student's observations. Explain your answer.
-The light bulb would light up dimmer than before or not even light up at all. -Dilute ethanoic acid is a weak acid. It does not completely ionise into water like hydrochloric acid. -Less mobile electrons, less electricity conducted
28
when making copper (II) sulfate crystals, suggest a reason why copper (II) oxide needed to be added till it no longer reacts with the acid
To ensure all sulfuric acid had reacted
29
state how you can tell that copper (II) oxide no longer reacts with the acid when making copper (II) sulfate crystals
When black solid copper (II) oxide can no longer dissolve in the acid even with stirring
30
state why filtration is done before heating the filtrate to obtain a saturated solution when making copper (II) sulfate crystals
to remove excess copper (II) oxide
31
when making copper (II) sulfate crystals, Instead of heating till a saturated solution is obtained, a student proposed to heat the filtrate to dryness. Is this a suitable method? Explain why.
No Water of crystallisation will be completely evaporated forming white anhydrous copper (II) sulfate powder, hence unable to obtain copper (II) sulfate crystals
32
when making copper (II) sulfate crystals, crystals start to form as the saturated solution is cooled. explain why
as temperature decreases, solubility of the crystals decreases, hence crystals start to form as the solution is cooled
33
write the ionic equation for the reaction: Na2SO4 (aq) + Ba(NO3)2 (aq) -> BaSO4 (s) + 2NaNO3 (aq)
SO4²⁻ (aq) + Ba²⁺ -> BaSO4 (s)
34
results table for titration after: 27.35, 38.45 before: 0.05, 1.20
35
state the 3 indicators name
-methyl orange -screened methyl orange -thymolphthalein
36
state the pH range of methyl orange
3-5
37
state the pH range of screened methyl orange
3-5
38
state the pH range of thymolphthalein
9.3-10.5
39
what is the colour of methyl orange in acidic solution
red
40
what is the colour of screened methyl orange in acidic solution
pink/violet
41
what is the colour of thymolphthalein in acidic solution
colourless
42
what is the colour at end point of methyl orange
orange
43
what is the colour at end point of screened methyl orange
grey
44
what is the colour at end point of thymolphthalein
first appearance/disappearance of blue
45
what is the colour in alkaline solution for methyl orange
yellow
46
what is the colour in alkaline solution for screened methyl orange
green
47
what is the colour in alkaline solution for thymolphthalein
blue
48
they provided a pH curve as some titrations do not require indicators (eg redox/iodometric titrations). what indicators do u use
choose an indicator whose range lies on the vertical section of the pH curve
49
define independent variable and where it is on a graph
a variable with values that are selected or changed by the investigator x-axis
50
define dependent variable and where it is on a graph
a variable that is being measured for each change in the independent variable/quantity that depends on the independent variable y-axis
51
define constant variable
a variable which could affect the independent variable and has to be kept constant to ensure a fair test / a variable that do not have direct relationship with what is investigated in the experiment but needs to be kept constant to ensure reliability of the experiment.
52
electronic balance smallest division, uncertainty, examples of recording
53
measuring cylinder smallest division, uncertainty, examples of recording
54
stopwatch smallest division, uncertainty, examples of recording
55
thermometer smallest division, uncertainty, examples of recording
56
ruler smallest division, uncertainty, examples of recording
57
the experiment need individual to record heat gained or lost from or to the surroundings state the source of error, how it affects results, how to rectify them
-Heat loss to surroundings (exothermic) / heat gain from surroundings (endothermic) -Causes temperature recorded to be lower than expected / Causes temperature recorded to be higher than expected -Provide lagging for the experimental setup
58
the experiment need individual to record amount of time via stopwatch state the source of error, how it affects results, how to rectify them
-Human reaction time error when using a stopwatch -Causes time recorded to be greater or less than expected -Use a data logger to measure time
59
the experiment need individual to record amount of solution via measuring cylinder state the source of error, how it affects results, how to rectify them
-use of measuring cylinder -Causes the volume of solution measured to be less accurate (greater or lower than expected) -Use a burette/ pipette instead
60
the experiment need individual to collect and find volume of gas via displacement of water state the source of error, how it affects results, how to rectify them
-gas collected is soluble in water -Causes volume of gas collected to be less than expected -Use a gas syringe (instead of displacement of water) -Repeat measurements and take average -Take more frequent readings
61
state the general rule for qualitative analysis
use minimum quantity of reagents as using too much reagent may lead to inaccurate observations
62
how do you test for solubility of precipitate
pour away at least half of the precipitate and carry out the solubility test on the remaining portion of precipitate by adding excess
63
when do u test for gas
when effervescence is observed or when heating is involved
64
what are somethings (3 points) to note during litmus test
-colour -odour of gas -colour change
65
what does it mean when the test tube feels cold
the reaction is endothermic
66
what does it mean when the test tube feels hot
reaction is exothermic
67
when heating, what do we need to take note (3 points)
-Test the gas (if any) produced during heating as it often indicates the presence of a certain Ions (e.g. NH4⁺ or NO3⁻). -Note the colour of residue when it is hot and when it is cold. (E.g. Zinc oxide, ZnO, is yellow when hot and white when cold.) -Always dry the test tube / boiling tube if only solid reagents are involved. Heat the unknown gently at first then strongly until no further changes takes place
68
4 points for clear recordings
-Identify the subject that causes the change clearly. For example, white precipitate is formed -Record in complete sentences. -Describe colour change by recording both initial and final colour. For example, "Acidified aqueous KMnO4 turned from purple to colourless", -Use only acceptable symbols like state symbols or chemical formula.
69
using NaOH solution test for what
cations
70
using aqueous NH3 test for what
cations
71
why are some precipitates soluble in excess NaOH?
Zn(OH)2 (and Al(OH)3 are amphoteric hydroxides. hence white precipitates formed are soluble in excess NaOH (aq) to form colourless solution
72
Write the ionic equation when copper compound reactions with aqueous NH3
Aqueous NH3 provides the OH⁻ ions as it is a weak alkali
73
Name some precipitates soluble in excess NH3?
Zn(OH)2 (white precipitate) and Cu(OH)2 (light blue precipitate) are soluble in excess NH3 (aq) to form their respective colourless or dark blue solutions
74
used hydrochloric acid. action of it test for what
anions
75
when used hydrochloric acid and CO2 is evolved, what does it mean
carbonate ions are present
76
used BaCl2 (aq) followed by dilute HCl or Ba(NO3)2 (aq) followed by dilute HNO3. action of it test for what
both solutions are used to detect anions that form insoluble barium salts
77
used BaCl2 (aq) followed by dilute HCl or Ba(NO3)2 (aq) followed by dilute HNO3. what is the purpose of adding acid
to remove interfering ions like CO3²⁻ and prevent precipitation of carbonate
78
acidify by HNO3 (aq) then add AgNO3 (aq). what does this test for
halides
79
if the gas produced is colourless and odourless, what do you do?
do NOT waste too much time with moist litmus paper. start with using a wooden splint
80
state the colourless and odourless gases (3)
carbon dioxide, oxygen, hydrogen
81
state the colourless and pungent gases (2)
ammonia, sulfur dioxide
82
state the gas that is coloured and pungent
chlorine
83
how will a positive test for oxygen look like
it relights a glowing splint
84
how will a positive test for hydrogen look like
extinguishes lighted splint with a “pop” sound
85
how will positive result for carbon dioxide look like
white precipitate CaCO3 is observed in limewater Ca(OH)2
86
how does a positive result for ammonia gas look like
the gas is very soluble so heating is usually required when tested for this gas. it turns moist red litmus paper blue
87
how does a positive result for sulfur dioxide gas look like
an acidic gas that decolourises purple acidified KMnO4
88
how does a positive result for chlorine gas look like
it turns moist blue litmus paper red, then bleaches it
89
what do you use to test for oxidising agents
by using a reducing agent
90
how do you test for an oxidising agent
-use acidified aqueous potassium iodide (KI) -if there is a presence of oxidising agents, it turns from colourless to brown -if there is an absence of oxidising agents, acidified aqueous KI remained colourless
91
how do you test for reducing agents
-use acidified aqueous KMnO4 -in the presence of reducing agents, acidified aqueous KMnO4 turned from purple to colourless -in the absence of reducing agents, acidified aqueous KMnO4 remains purple
92
add dilute acid. effervescence, carbon dioxide is produced what is the anion that is tested
carbonate (CO3²⁻)
93
add dilute nitric acid, then add aqueous silver nitrate. white ppt is formed. what is the anion that is tested
chloride (Cl⁻)
94
add dilute nitric acid, then add aqueous silver nitrate. yellow ppt is formed. what is the anion that is tested
iodide (I⁻)
95
add dilute nitric acid, then add aqueous silver nitrate. what are the possible anions that is tested for
chloride (Cl⁻), iodide (I⁻)
96
add dilute nitric acid. what are the possible anions that is tested for
carbonate (CO3²⁻) chloride (Cl⁻), iodide (I⁻), sulfate (SO4²⁻)
97
add dilute nitric acid, then add aqueous barium nitrate. white ppt is formed. what is the anion that is tested
sulfate (SO4²⁻)
98
add aqueous sodium hydroxide, then add aluminium foil, then warm carefully. ammonia gas is formed. what is the anion that is tested
nitrate (NO3⁻)
99
how do you test for carbonate ions
-add dilute acid -effervescence observed -carbon dioxide produced
100
how do you test for chloride ions
-acidify with dilute nitric acid -add aqueous silver nitrate -white precipitate is observed
101
how do you test for iodide ions
-acidify with dilute nitric acid -add aqueous silver nitrate -yellow precipitate is observed
102
how to test for nitrate ions
-add aqueous sodium hydroxide -add aluminium foil -warm carefully -ammonia gas produced
103
how to test for sulfate ions
-acidify with dilute nitric acid -add aqueous barium nitrate -white precipitate observed
104
how to test for carbon dioxide
gives white precipitate in limewater. precipitate dissolves with excess CO2
105
white precipitate is formed when added aqueous sodium hydroxide. state the possible cations present
-aluminium (Al³⁺) -calcium (Ca²⁺) -zinc (Zn²⁺)
106
white precipitate formed and is soluble in excess in aqueous sodium hydroxide. state the possible anions
aluminium (Al³⁺) zinc (Zn²⁺)
107
add sodium hydroxide into solution and warm. what cation is tested
ammonium (NH4⁺)
108
add excess aqueous sodium hydroxide into Cu²⁺ cation present solution. what are the observations?
light blue precipitate is insoluble in excess
109
green precipitate is formed when aqueous sodium hydroxide is added. what is the cation present
Fe²⁺
110
red-brown precipitate is formed when aqueous sodium hydroxide is added. what is the cation present
Fe³⁺
111
state the observations when NaOH is added to solution with aluminium cation present
white precipitate is observed, soluble in excess giving a colourless solution
112
state the observations when NaOH is added to solution with ammonium cation present
ammonia produced on warming
113
state the observations when NaOH is added to solution with calcium cation present
white precipitate observed, insoluble in excess giving
114
state the observations when NaOH is added to solution with copper (II) cation present
light blue precipitate observed, insoluble in excess giving
115
state the observations when NaOH is added to solution with iron (II) cation present
green precipitate observed, insoluble in excess giving
116
state the observations when NaOH is added to solution with iron (III) cation present
red-brown precipitate formed, insoluble in excess giving
117
state the observations when NaOH is added to solution with zinc cation present
white precipitate, soluble in excess giving a colourless solution
118
after adding aqueous ammonia, what cation causes no precipitate to be observed
calcium (Ca²⁺)
119
white precipitate is formed when added aqueous ammonia. state the possible cations present
aluminium (Al³⁺), zinc (Zn²⁺)
120
which cations are soluble in excess ammonia
zinc (Zn²⁺). white precipitate is soluble in excess giving a colourless solution copper (II) (Cu²⁺). light blue precipitate is soluble in excess giving a dark blue solution
121
state the observations when aqueous ammonia is added to solution with aluminium (Al³⁺)
white precipitate observed, insoluble in excess
122
state the observations when aqueous ammonia is added to solution with calcium (Ca²⁺)
no precipitate formed
123
state the observations when aqueous ammonia is added to solution with copper (II) (Cu²⁺)
light blue precipitate is observed, soluble in excess giving a dark blue solution
124
state the observations when aqueous ammonia is added to solution with iron (II) (Fe²⁺)
green precipitate, insoluble in excess giving
125
state the observations when aqueous ammonia is added to solution with iron (III) (Fe³⁺)
red-brown precipitate, insoluble in excess
126
state the observations when aqueous ammonia is added to solution with zinc (Zn²⁺)
white precipitate, soluble in excess giving a colourless solution
127
colour of copper
pink (fresh), red brown
128
colour of lead (II) oxide
yellow. solid.
129
colour of lead (II) iodide
yellow, solid
130
iron (II) chloride solid colour
pale green
131
iron (II) chloride solution colour
pale green
132
iron (II) nitrate solid colour
pale green
133
iron (II) nitrate solution colour
pale green
134
iron (II) sulfate solid colour
green
135
iron (II) sulfate solution colour
green
136
iron (II) carbonate solid colour
green
137
iron (II) hydroxide solid colour
green
138
iron (II) oxide solid colour
black
139
state colour change when potassium manganate (VII) is added to iron (II) sulfate
green solution turns brown. if left to stand, red-brown precipitate is observed
140
potassium oxide solid colour
pale yellow
141
zinc oxide solid colour
white (cold), yellow (hot)
142
copper (II) sulfate solid colour
blue (hydrated), white (anhydrous)
143
copper (II) carbonate solid colour
green
144
copper (II) oxide solid colour
black
145
copper (II) hydroxide solid colour
blue
146
iron (III) chloride solid colour
yellow
147
iron (III) sulfate solid colour
yellow
148
iron (III) carbonate solid colour
red-brown
149
iron (III) oxide solid colour
red-brown
150
iron (III) hydroxide solid colour
red-brown
151
iron (III) chloride solution colour
yellow
152
iron (III) sulfate solution colour
yellow
153
colour of fluorine gas
pale yellow
154
colour of chlorine gas
yellow-green
155
colour of aqueous chlorine
yellow
156
colour of liquid bromine
red-brown
157
colour of aqueous bromine
red-brown
158
colour of solid iodine
purple-black
159
colour of aqueous iodine (dissolved in iodide solution
brown
160
colour of gas iodine
violet
161
metal reactivity series
Professor (potassium) Snape (sodium) Caught (calcium) Malfoy (magnesium) Carbonating (carbon) Zabini (zinc) Hermione (hydrogen gas) Investigates (Iron) Later (Lead) Hermione (hydrogen) Cried (copper) Softly (silver) Goodbye (gold)
O LVL Chemistry
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