inorganic chemistry Flashcards
(98 cards)
1
Q
group 1 metals
A
alkali metals (form alkaline solutions when they react with water)
2
Q
properties of alkali metals
A
-soft metals, can be cut with a knife
-low densities and melting points
-very reactive (lose on electron to become stable)
-shiny when cut
-stable metals are grey except for caesium
3
Q
group 1 metal reaction with water
A
group 1 metal +water -> metal hydroxide + hydrogen
4
Q
Colour of metal hydroxide
A
Colourless (aq)
5
Q
How does the reactions of alkali metals with water change
A
React more vigorously as you descend the group
6
Q
Observations of Li + H2O
A
-(slow reaction)
-fizzing
-lithium moves on the surface of the water, slowly becomes smaller
7
Q
Na + H2O reaction observations
A
-fizzing
-moves rapidly on the surface
-dissolves quickly
8
Q
Potassium and water reaction observations
A
-Burns with a lilac flame
-moves rapidly on the surface
-dissolves quickly
-fizzing?
9
Q
What does alkali metal + oxygen form and what are the observations
A
-forms metal oxides
-alkali metals tarnish when exposed to air
-metal oxide is a dull coating that covers the surface of the metal
-metal tarnishes more rapidly as you go down the group
10
Q
How do properties of alkali metals change as you go down the group
A
-
11
Q
Properties of rubidium based on trends
A
-soft grey solid
-appears shiny when freshly cut
-more dense that potassium
-lower melting point than potassium
12
Q
Why do alkali metals get more reactive as you go down
A
-More shells of electrons
-outermost electron is further away from the nucleus, weaker forces of attraction between outermost (valence) electron and nucleus
-less energy is required to overcome force of attraction, so outer electron is lost more easily
13
Q
features of halogens
A
-poisonous
-diatomic
14
Q
state and colour of fluorine (solution and by itself)
A
15
Q
state and colour of chlorine (solution and by itself)
A
-pale yellow-green gas
-pale green in solution
16
Q
state and colour of bromine (solution and by itself)
A
-red-brown liquid
-orange in solution
17
Q
state and colour of iodine (solution and by itself)
A
-grey solid
-dark brown in solution
18
Q
how does the boiling point of halogens change as you go down the group
A
increases
-increasing IMF as atoms become larger
19
Q
characteristics of iodine
A
shiny, crystalline solid that sublimes to form a purple vapour
20
Q
metal halides
A
halogens react with some metals to form ionic compounds
21
Q
non metal halides
A
halogens react with non metals to form simple molecular covalent structures
22
Q
halogen displacement reactions
A
a more reactive displaces a less reactive halogen from an aqueous solution of its halide
23
Q
why does the reactivity of halogens decrease as you go down the group
A
-More shells of electrons
-increased distance from outer shells to the nucleus decreases the fore of attraction between the nucleus and the outermost shell
-harder for the atoms to gain electrons as you descend the group
24
Q
composition of air
A
78% nitrogen
21% oxygen
1% includes - CO2, water vapour, argon (argon 0.9%, CO2 0.04%
25
how to find the percentage of oxygen in air
-burn phosphorus in a bell jar that is sitting in a trough of water
-as phosphorus burns it uses up oxygen inside the bell jar and water level rises P2O5 (phosphorus pentoxide)
-use measurements of water levels before and after the experiment to determine % of oxygen in the air
-phosphorus is suitable as it burns readily until all oxygen is used up
-phosphorus is toxic
26
combustion
-burning
-oxygen reacts with elements to produce oxides
-gives out heat, exothermic
27
combustion reaction with magnesium observation and chemical equation
-intense white flame, white powder produced
2Mg (s) + O2 (g) → 2MgO (s)
28
combustion reaction with hydrogen observation and chemical equation
-exothermic, water is produced (flame almost colourless), (squeaky pop)
2H2 (g) + O2 (g) → 2H2O (g)
29
combustion reaction with sulfur observation and chemical equation
-blue flame, colourless, poisonous gas produced
S (s) + O2 (g) → 2SO (g)
30
thermal decomposition
reactions where a substance breaks down due to the action of heat
31
metal carbonate decomposition
metal carbonate -> metal oxide + carbon dioxide
32
observations of the thermal decomposition of copper II carbonate
copper carbonate is a green powder and darkens as black copper II oxide is produced, CO2 also produced
copper II carbonate -> copper II oxide + CO2
33
3 greenhouse gases
CO2, methane, water vapour
34
greenhouse effect
-short wavelength (UV) radiation is emitted from the sun
-some is absorbed when it hits the earth's surface, some is re-emitted as long wavelength radiation (infrared)
35
sources of CO2
-combustion of wood and fossil fuels
-respiration of plants and animals
-thermal decomposition of carbonate rocks and effect of acids on carbonates
36
reactivity series
-
37
what metals reacts with dilute acids
-metals more reactive than hydrogen
38
how do metal displacement reactions work
-a more reactive metal will displace a less reactive metal from its compound
39
reactivity series
Potassium
Sodium
Lithium
Calcium
Magnesium
Aluminium
Carbon
Zinc
Iron
Tin
Lead
Hydrogen
Copper
Silver
Gold
Platinum
Please send Lions, Cats, Monkeys and cute Zebras into the least hot country, Signed: GP
40
conditions for iron to rust
-oxygen and water
41
iron rusting equation
iron +oxygen + water -> hydrated iron(III) oxide
42
how can iron structures be damaged by rusting
-rust is a soft solid that flakes off the iron surface easily, exposing fresh iron below
-over time all of the iron rusts and the structure is weakened
43
barrier methods
-coat iron in paint, oil, grease or electroplating
-prevents iron from coming into contact with water and oxygen
44
sacrificial protection
-a more reactive metal is attached to iron
-more reactive metal will oxidise first and corrode first, protecting the iron
-for continued protection, the sacrificial metal must be replaced after it had corroded (not fully)
45
galvanising
-process where the iron to be protected is coated with zinc
-done by electroplating or by dipping it into molten zinc
-Zinc carbonate is formed after contact with air and protects iron by barrier
-if the coating is damaged, the iron is protected by sacrificial protection
46
oxidation
a reaction in which a substance gains oxygen
47
reduction
a reaction in which a substance loses oxygen
48
redox reaction
a reaction where reduction and oxidation both take place
49
oxidising agent
the substance that supplies the oxygen (loses it)
50
reduction agent
the substance that removes that oxygen (gains the oxygen as a product)
51
oxidation and reduction reactions in terms of electrons
-oxidation is gaining electron
-reduction is losing electron
OILRIG
52
how are useful metal that are chemically combined with other substances to form ores extracted
-electrolysis
-blast furnace
-reacting with a more reactive metal
53
extraction method metals more reactive than carbon
-electrolysis of the molten chloride or oxide
-expensive as lots of electricity is required
54
extraction method metals less reactive than carbon
-heating with a reducing agent such as carbon or CO in a blast furnace
-cheap as carbon is cheap and can be a source of heat
55
extraction of silver and gold
found as pure elements already
56
steps to extract iron
1. coke burns in the hot air to form carbon dioxide, exothermic reaction heats the furnace
carbon (coke) + oxygen -> CO2
2. more coke reacts with carbon dioxide forming carbon monoxide
CO2 + carbon -> carbon monoxide
3. carbon monoxide reduces the iron (III) oxide in the iron ore to form iron. It will melt and collect at the bottom of the furnace, where it is tapped off
Fe2O3 + 3CO -> 2Fe(I) + 3CO2
4. calcium carbonate added to the furnace to reduce impurities in the ore
Calcium carbonate -> calcium oxide + carbon oxide
-calcium Oxide formed reacts with the impurity silicon dioxide to form calcium silicate by neutralisation
calcium oxide + silicon dioxide -> calcium silicate (CaSiO3). This melts and floats on top of the molten iron and is separately tapped off
57
steps to extract aluminium
1. bauxite oxide is purified to produce aluminium oxide
2. aluminium oxide is dissolved in molten cryolite (molten aluminium oxide would make it expensive, molten cryolite is lower melting point)
3. mixture placed in an electrolysis cell
4. at the cathode: aluminium ions gain electrons, molten aluminium forms at the bottom of the cell
5. at the anode: oxide ions lose electrons and oxygen is produced. carbon in the graphite anodes react with oxygen to produce CO2. Anode has to be replaced regularly
58
uses of aluminium
-aircraft bodies (high strength to weight ratio)
-saucepans (good conductor of heat and unreactive)
-overhead electrical cables (good conductor of electricity)
-food cans (non toxic, resistant to corrosion and acidic food)
59
uses of copper
-electrical wiring (good conductor, ductile)
-saucepan (good conductor, unreactive, malleable)
-water pipes (unreactive, non toxic, malleable)
60
uses of iron
-building material (strong, malleable and ductile, cheap)
-catalyst (increases rate of reaction without being used up)
61
uses of steel
mild steel: car body panels and writing (soft and malleable)
high carbon steel: tools (hard)
stainless steel: cutlery, sinks, chemical plants (strong and resistant to corrosion)
62
alloy
mixture of two or more metals or metal with a non metal e.g. carbon
63
why are alloys harder
-alloys contain atoms of different sizes
-this distorts the regular arrangement of atoms
-it is more difficult for the layers of atoms to slide over each other
63
properties of alloys
-stronger and harder
-resistant to corrosion or high temperatures
64
red litmus paper colours
a: red
n: red
b: blue
65
blue litmus paper colours
a: red
n: blue
b: blue
66
phenolphthalein colours
a: colourless
n: colourless
b: pink
67
methyl orange colours
a: red
n: yellow
b: yellow
68
universal indicator colours
strong acid: red
medium acid: orange
weak acid: yellow
neutral: green
weak alkaline: dark green: light blue
medium alkaline: dark blue
strong alkaline: purple
69
what happens when acids are added to water
-proton donors
-they form positively charged hydrogen ions (H+) in water
-the presence of H+ ions is what makes a solution acidic
70
what are bases
-proton acceptors
-they form hydroxide ions in water which can accept protons (OH-)
-presence of OH- ions is what makes the aqueous solution an alkali
71
what happens in a neutralisation reaction
-acid reacts with an alkali, H+ reacts with OH- to produce water
72
use of titration reactions
-analysing concentration of solutions
-prepare salts
73
method for finding the volume needed to neutralise x amount of a solution
1. use the pipette to place 25 cm alkali solution into the conical flask
2. fill burette with acid and place an empty beaker underneath the tap, run some acid through the burette to remove any air bubbles
3. record the starting point on the burette
4. place the conical flask on a white tile
5. add few drops of indicator to the solution
6.add the acid into the flask, swirling continuously
7. continue until end point is reached, sharp colour change and record the volume
8. record the final burette reading and calculate the volume of acid added
9. repeat the titration, adding drops slowly as the end point is reached, to get more accurate results
(use the solution you need to find the volume of in the burette, use methyl orange is base starts in conical flask
74
what salts are always soluble
SAP
sodium
ammonium
potassium
75
nitrates solubility
all soluble
76
chlorides solubility
all except silver and lead II
77
sulfates solubility
all except barium, calcium and lead II
78
carbonates solubility
insoluble except SAP
79
hydroxides solubility
insoluble except SAP and slightly calcium
80
acid + base reactions
salt + water
-salt formed depends on the acid, can be chloride, sulfate or nitrate salt
81
what can act as bases
metal oxides, hydroxides and carbonates
82
acid and metal carbonate reaction
salt, carbon dioxide and water
fizzing observed due to CO2
83
ammonia in water ions
NH4 + OH
84
how to prepare a soluble salt with insoluble base and acid
1. insoluble reactant is added in excess to ensure all of the acid has reacted (unreacted acid would be dangerously concentrated after evaporation and crystallisation)
2. stir mixture to make sure reaction is complete
3. filter the mixture to remove the excess
4. transfer the solution into an evaporating dish
5. heat the filtrate until saturated
6. check saturation point using a glass rod - if crystals form then saturation point is reached
7. allow to cool and crystalise
8. filter to remove crystals
9. dry crystals in a warm oven
85
method to prepare a soluble salt from soluble base
1. add alkali and indicator to a conical flask using a pipette (phenolphthalein)
2.add acid to burette, record the starting volume
3. add acid to alkali slowly, continuously swirling the flask until indicator changes colour, calculate volume of acid added
4. repeat using the same volumes without indicator
5. transfer solution to evaporating dish
6. heat the solution until saturated, check sat. point using glass rod - if crystals form then sat. point has been reached
7. allow to cool and crystalise
8. filter to remove crystals
9. dry crystals in a warm oven
86
prepare an insoluble salt from two soluble salts
1. mix the solutions
2. stir the mixture
3. filter off the precipitate
4. wash with distilled water
5. dry the solid in a warm oven or between paper towel
87
test for ammonia
turn damp red litmus paper blue
88
test for carbon dioxide
-bubble gas through an aqueous solution of limewater
-turns cloudy white/milky if carbon dioxide is present
89
chlorine gas test
damp blue litmus paper will turn red then be bleached white (turn reds as acids are produce when chlorine comes into contact with water
90
test for hydrogen gas
hold a burning splint at the open end of a test tube of gas
-if gas is hydrogen it burns with a loud "squeaky pop" (combustion of hydrogen with oxygen to produce water
-insert the splint right into the tube as the gas needs air o burn
91
test for oxygen
place a glowing splint inside a test tube of gas, if it relights, oxygen is present
92
what is the flame test
used to identify the positive metal ion by the colour of the flame they produce
93
how to carry out a flame test
1. dip the loop of an unreactive metal wire such as nichrome or platinum in dilute acid
2. hold it in the blue flame of a Bunsen burner until there is no colour change
dip the loop into the solid sample/solution and place it in the edge of the blue Bunsen flame
94
flame test colours
Li - Red
Na - Yellow
K - Lilac
Ca - orange red
Cu - blue green
95
test for cations and anions
-
96
chemical test to identify water
anhydrous copper II sulfate turns from white to blue on contact with water
CuSO4 (s) + 5H2O (l) → CuSO4.5H2O (s)
97
Physical test for water
sample of liquid is placed in a boiling tube and gently heated
-using a thermometer, check if the boiling point is exactly 100C
(impurities will raise the boiling point and lower the melting point of pure substance)
