Flashcards in Module 3: Section 1 Deck (97)
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1
What is a giant covalent lattice?
a huge network of covalently bonded atoms.
2
What is another name for a giant covalent lattice?
macromolecular structure.
3
What is an allotrope?
a different form of the same element in the same state.
4
Name the 3 allotropes of carbon.
diamond, graphite, graphene.
5
In diamond, how many covalent bonds does each carbon atom form?
each carbon atom is covalently bonded to 4 other carbon atoms.
6
In diamond, what shape are the atoms arranged in?
tetrahedral shape. (crystal lattice structure)
7
What does 'sublimes' mean?
when a substance changes straight from a solid to a gas, skipping out the liquid stage.
8
What properties are caused by lots of strong covalent bonds in diamond?
- very high melting point (it sublimes at over 3800K)
-extremely hard (used in diamond-tipped drills and saws)
-good thermal conductor (vibrations travel easily through the stiff lattice)
-insoluble in any solvent
-can't conduct electricity (no delocalised electrons)
9
Why are diamonds 'cut' to form gemstones?
it's structure makes it refract light a lot, so it sparkles.
10
In silicon, how many covalent bonds does each atom form?
each silicon atom forms 4 strong covalent bonds
11
What structure does silicon form?
silicon (same periodic group as carbon) forms a crystal lattice structure with similar properties to carbon.
12
In graphite, how many covalent bonds does each carbon atom form?
each carbon atom is covalently bonded to 3 other carbon atoms.
13
What is the structure of graphite?
-carbon atoms are arranged in sheets of flat hexagons covalently bonded with 3 bonds each
-the fourth outer electron of each carbon atom is delocalised between the sheets of hexagons
-the sheets are bonded together by weak induced dipole-dipole forces
14
What properties of graphite are caused by its structure?
-very high melting point (it sublimes at over 3900K)
-conducts electricity (delocalised electrons aren't attached to any particular carbon and are free to move along the sheets, so an electric current can flow)
-layers can slide over each other (weak intermolecular forces between layers are easily broken)
-insoluble in any solvent (covalent bonds in sheets are too strong to break)
15
Why is graphite less dense than diamond?
graphite is less dense than diamond as the layers are quite far apart compared to the length of the covalent bonds.
16
What are the uses of graphite?
-due to sheets being able to slide over each other, graphite feels slippery and is used as a dry lubricant and in pencils.
-used to make strong, lightweight sports equipment
17
What is graphene?
graphene is one layer of graphite.
18
What is the structure of graphene?
-a sheet of carbon atoms joined together in hexagons
-each carbon atom has 3 covalent bonds and one delocalised electron
-sheet is one atom thick (two-dimensional compound)
19
What are the properties of graphene?
-very high melting and boiling point
-insoluble
-best known electrical conductor (delocalised electrons are free to move and carry charge, without layers they can move quickly above and below the sheet)
-extremely strong (delocalised electrons strengthen covalent bonds between carbon atoms)
-single layer of graphene is transparent and very light
20
Due to its high strength, low mass and good electrical conductivity, what are potential uses of graphene?
potential applications in high-speed electronics and aircraft technology.
21
Due to its flexibility and transparency, what other uses of graphene are there?
touchscreens on smart phones and other electronic devices.
22
What is the structure of metal elements?
metal elements exist as giant metallic lattice structures.
23
How is a giant metallic lattice formed?
the electrons in the outermost shell of a metal atom are delocalised (free to move about the metal), leaving a positively charged metal cation e.g. Na+
24
Describe metallic bonding.
the positively charged metal cations are electrostatically attracted to the delocalised negative electrons. They form a lattice of closely packed cations in a sea of delocalised electrons.
25
What are the properties of metals?
-malleable (can be hammered into sheets) and ductile (can be drawn into a wire), no bonds holding specific ions together so metal ions can slide past each other when the structure is pulled
-good thermal conductors (delocalised electrons can pass kinetic energy to each other)
-good electrical conductors (delocalise electrons are free to move and carry charge)
-insoluble, except in liquid metals (because of strength of metallic bonds)
26
What factors affect the melting point of metals?
-number of delocalised electrons per atom (more delocalised electrons, stronger bonding, higher MP) e.g. Mg2+ has a higher MP than Na+
-size of metal ion (smaller ionic radius will hold delocalised electrons closer to the nuclei)
27
What are simple molecular structures?
simple molecular structures contain only a few atoms e.g. O2, Cl2
28
Why do simple molecular substances have low melting and boiling points?
the covalent bonds between the atoms in the molecule are very strong, but the intermolecular forces between the molecules are weak and easily overcome, so these elements have low melting and boiling points.
29
What do the melting and boiling points depend on?
the melting and boiling points depend on the strength of the induced dipole-dipole forces between the molecules. (more atoms in a molecule = stronger induced dipole-dipole forces)
30
E.g. in period 3, why does sulfur have a higher melting and boiling point than phosphorus or chlorine?
in period 3, sulfur is the biggest molecule (S8), so it contains more atoms so has stronger induced dipole-dipole forces and has a higher melting and boiling point than phosphorus or chlorine.
31
Why do the noble gases have very low melting and boiling points?
the noble gases exist as individual atoms (they're monatomic), resulting in very weak induced dipole-dipole forces.
32
As you go across a period, what affects the melting and boiling points of the elements?
as you go across a period, the type of bond formed between the atoms of an element changes (different bond strengths), this affects the melting and boiling point of an element.
33
Describe how the melting and boiling points change across period 3.
-for the metals (Na, Mg, Al), MP and BP increase across the period because the metallic bonding gets stronger as the ionic radius decreases (increasing nuclear charge) and the number of delocalised electrons increases
-MP and BP increase at Si because it is a giant covalent lattice structure with lots of strong covalent bonds which require a lot of energy to break
-MP and BP decrease for elements that form simple molecular structures (P, S, Cl) because the weak intermolecular forces between molecules are easy to overcome
-the noble gases (Ar) have the lowest MP and BP in their period because they are monatomic and are held together by the weakest forces
34
E.g. why is the melting point of Mg higher than that of Na?
Mg has more delocalised electrons per atom and a smaller ionic radius. So the electrostatic attraction between the metal ions and the delocalised electrons is stronger.
35
What ions do group 2 elements form?
they lose their 2 outer electrons to form 2+ ions.
36
What is the trend in reactivity down group 2?
reactivity increases down group 2 due to the increasing atomic radius and shielding effect. When group 2 elements react they lose electrons, forming positive cations. The easier it is lose electrons (the lower the first and second ionisation energies), the more reactive the element, so reactivity increases down the group.
37
When group 2 elements react, are they oxidised or reduced?
group 2 elements are oxidised from a state of 0 to +2
M ---> M2+ + 2e-
(act as reducing agents)
38
Redox Reactions of group 2 elements:
How do group 2 elements react with oxygen?
group 2 elements burn in oxygen to form solid white oxides.
2M(s) + O2(g) --> 2MO(s)
39
How do group 2 elements react with water?
Observations?
group 2 elements react with water to produce a metal hydroxide and hydrogen.
M(s) + 2H2O(l) --> M(OH)2(aq) + H2(g)
Observations:
solid dissolves, fizzing, ph of solution more than 7 (metal hydroxides are alkalis)
half equation for water:
2H2O + 2e- --> 2OH- + H2
40
How do group 2 elements react with dilute acid?
group 2 elements react with dilute acid to produce a salt and hydrogen. (different acids produce different salts)
e.g. group 2 metals react with dilute HCl to produce a metal chloride and hydrogen
M(s) + 2HCl(aq) --> MCl2(aq) + H2(g)
41
What is a base?
a base is a proton acceptor.
42
What is an alkali?
an alkali is a base that's soluble in water and will release OH- ions in an aqueous solution.
43
What are group 2 oxides and hydroxides classed as?
group 2 oxides and hydroxides are bases. Most of them are soluble in water so are also alkalis.
44
How do group 2 oxides and hydroxides react with acids?
Observations?
group 2 oxides and hydroxides are bases so react with acids to form a salt + water.
Observations:
solid dissolves
45
How do group 2 oxides react with water?
What will be the ph of the solution?
group 2 oxides react readily with water to form metal hydroxides, which dissolve.
the OH- ions make the solution strongly alkaline (e.g. 10-13)
e.g. CaO(s) + H2O(l) --> Ca2+(aq) + 2OH-(aq)
Ca(OH)2(s) + H2O(l) --> Ca2+(aq) + 2OH-(aq)
46
Which group 2 metal oxide is an exception?
magnesium oxide is an exception - it only reacts slowly and the hydroxide isn't very soluble.
47
Why do the oxides form more strongly alkaline solutions as you go down group 2?
as you go down group 2, the hydroxides become more soluble so solutions become more alkaline (higher concentration of OH- ions)
48
How do group 2 carbonates react with acids?
Observations?
group 2 carbonates neutralise acids to produce a salt + water + CO2
Observations:
solid dissolves, fizzing/bubbles of gas
49
What happens when group 2 carbonates are heated?
group 2 carbonates break down when heated - thermal decomposition.
e.g. CaCO3(s) --> CaO(s) + CO2(g)
50
What is the trend in thermal decomposition down group 2?
group 2 elements become more difficult to decompose as you go down the group.
51
What are the group 2 elements known as?
the alkaline earth metals.
52
What are uses of group 2 compounds?
group 2 compounds are used to neutralise acidity:
-calcium hydroxide, slaked lime Ca(OH)2, is used in agriculture to neutralise acidic soils
-magnesium hydroxide, Mg(OH)2, and calcium carbonate, CaCO3, are used in some indigestion tablets as antacids
53
What is the ionic equation for neutralisation?
H+(aq) + OH-(aq) --> H2O(l)
54
What group are the halogens?
group 7
55
What are the colours and physical states (at 20°C) of the following halogens:
-Fluorine
-Chlorine
-Bromine
-Iodine
-fluorine, pale yellow, gas
-chlorine, green, gas
-bromine, red-brown, liquid
-iodine, grey, solid
56
In what form do the halogens exist?
halogens exist as diatomic molecules (two atoms joined by a single covalent bond).
57
What is the trend in melting and boiling points down group 7?
melting and boiling points increase down the group due to increasing strength of London forces as the size and relative mass of the atoms increases.
58
What is meant by volatile?
a substance is volatile if it has a low boiling point.
59
What is the trend in volatility down group 7?
volatility decreases down group 7.
60
What does 'halogen' describe?
'halogen' is used to describe the atom (X) or molecule (X2).
61
What does 'halide' describe?
'halide' is used to describe the negative ion (X-)
62
When halogens react, are they oxidised or reduced?
halogen atoms react by gaining an electron into their outer shell to form 1- ions, so they're reduced. (they're oxidising agents)
oxidation number 0 --> -1
63
What is the trend in reactivity down group 7?
reactivity decreases down group 7. As you go down the group, the atomic radii increase so outer electrons are further from the nucleus and shielding increases as there are more inner electrons. This makes it harder for larger atoms to attract the electron needed to form an ion (despite increasing nuclear charge), so larger atoms are less reactive.
64
What is another way of saying that the halogens get less reactive down the group?
the halogens become less oxidising down the group.
65
What is an oxidising agent?
an electron acceptor.
66
What do displacement reactions of halogens with halide ions show?
displacement reactions show the relative oxidising strengths of the halogens.
67
What colour is chlorine water?
Colour in cyclohexane?
colourless
colourless
68
What colour is bromine water?
Colour in cyclohexane?
yellow
orange
69
What colour is iodine solution?
Colour in cyclohexane?
orange/brown
purple
70
Cl2(aq) + 2KBr(aq) --> ?
Cl2(aq) + 2Br-(aq) -->
Colour change:
-in aqueous solution?
-in organic solution (cyclohexane)?
Cl2(aq) + 2KBr(aq) --> 2KCl(aq) + Br2(aq)
Cl2(aq) + 2Br-(aq) --> 2Cl-(aq) + Br2(aq)
-aqueous solution = yellow(Br2)
-cyclohexane = orange(Br2)
71
Cl2(aq) + 2KI(aq) --> ?
Cl2(aq) + 2I-(aq) -->
Colour change:
-in aqueous solution?
-in organic solution(cyclohexane)?
Cl2(aq) + 2KI(aq) --> 2KCl(aq) + I2(aq)
Cl2(aq) + 2I-(aq) --> 2Cl-(aq) + I2(aq)
-aqueous solution = orange/brown(I2)
-cyclohexane = purple(I2)
72
Br2(aq) + 2KI(aq) --> ?
Br2(aq) + 2I-(aq) -->
Colour change:
-in aqueous solution?
-in organic solution(cyclohexane)?
Br2(aq) + 2KI(aq) --> 2KBr(aq) + I2(aq)
Br2(aq) + 2I-(aq) --> 2Br-(aq) + I2(aq)
-in aqueous solution = orange/brown(I2)
-cyclohexane = purple(I2)
73
Why do you shake the reaction mixture with an organic solvent like cyclohexane?
the halogen that's present will dissolve readily in the organic solvent, which settles out as a distinct layer above the aqueous solution, making the colour changes easier to see.
74
How can displacement reactions help to identify which halogen (or halide) is present in a solution?
a halogen will displace a halide from solution if the halide is below it in the periodic table. (a halogen will oxidise a halide if the halide is below it in the periodic table).
75
Describe the test for halides.
1. Add dilute nitric acid, to remove ions the might interfere with the test
2. Add silver nitrate solution, AgNO3(aq), a precipitate of the silver halide is formed (colour of ppt identifies the halide)
3. To be certain, add ammonia solution as each silver halide has a different solubility in ammonia - the larger the ion, the more difficult it is to dissolve.
Cl- : white ppt, dissolves in dilute NH3(aq)
Br- : cream ppt, dissolves in conc. NH3(aq)
I- : yellow ppt, insoluble in conc. NH3(aq)
Ag+(aq) + X-(aq) --> AgX(s)
76
What happens when halogens react with cold dilute alkali solutions?
the halogen undergoes disproportionation (it is simultaneously oxidised and reduced)
77
What oxidation states can chlorine exist in?
chloride, Cl- = -1
chlorine, Cl2 = 0
chlorate(I), ClO- = +1
78
Which halogen cannot exist in a wide range of oxidation states?
fluorine
79
What is produced when chlorine gas reacts with cold, dilute aqueous sodium hydroxide?
2NaOH(aq) + Cl2(g) --> ?
if you react chlorine gas with cold, dilute aqueous sodium hydroxide, you get sodium chlorate(I) solution, NaClO(aq), which is household bleach.
2NaOH(aq) + Cl2(g) --> NaClO(aq) + NaCl(aq) + H2O(l)
80
Why is it a disproportionation reaction?
2NaOH(aq) + Cl2(g) --> NaClO(aq) + NaCl(aq) + H2O(l)
2NaOH(aq) + Cl2(g) --> NaClO(aq) + NaCl(aq) + H2O(l)
0 +1 -1
the oxidation number of Cl increases and decreases, so it's a disproportionation reaction.
81
What are the uses of bleach (sodium chlorate(I) solution)?
water treatment, to bleach paper and textiles, cleaning
82
What is a disproportionation reaction?
when a species is both oxidised and reduced in the same reaction.
83
Balance the disproportionation equation:
Cl2 + NaOH --> NaCl + NaClO3 + H2O
hot conc.
3Cl2 + 6NaOH --> 5NaCl + NaClO3 + 3H2O
0 -1 +5
*use oxidation numbers to help balance
84
What happens when you mix chlorine with water?
Equation?
when you mix chlorine with water, it undergoes disproportionation, you produce a mixture of hydrochloric acid and chloric(I) acid (also called hypochlorous acid).
Cl2(g) + H2O(l) HCl(aq) + HClO(aq)
0 -1 +1
hydrochloric acid chloric(I) acid
85
How is chlorine used to kill bacteria in water?
After mixing chlorine with water, the aqueous chloric(I) acid produced ionises to make chlorate(I) ions (also called hypochlorite ions). Chlorate(I) ions kill bacteria.
So, adding chlorine (or a compound containing chlorate(I) ions) to water can make it safe to drink or swim in.
86
Why is chlorine important in water treatment?
-it kills disease-causing microorganisms
-some chlorine remains in the water and prevents reinfection further down the supply
-it prevents the growth of algae, eliminating bad tastes and smells, and removes discolouration caused by organic compounds
87
What are the risks of using chlorine to treat water?
-chlorine gas is very harmful if breathed in (irritates respiratory system)
-liquid chlorine on skin or eyes causes severe chemical burns
-water contains organic compounds (e.g. from decomposition of plants), chlorine reacts with these compounds to form chlorinated hydrocarbons (e.g. chloromethane, CH3Cl) many of which are carcinogenic.
However, increased cancer risk is small compared to risks from untreated water e.g. a cholera epidemic could kill thousands.
88
What are the ethical considerations of chlorinated water?
we don't get a choice about having our water chlorinated - some people object to this as forced 'mass medication'.
89
What are the alternatives to chlorine?
Disadvantages?
Ozone, O3 - a strong oxidising agent so its good at killing microorganisms. But, its expensive to produce and its short half-life in water means that treatment isn't permanent.
Ultraviolet light - kills microorganisms by damaging their DNA. But, its ineffective in cloudy water and, like O3, it won't stop the water being contaminated further down the line.
90
What reaction takes place when you react bromine or iodine with water?
bromine and iodine also undergo disproportionation when mixed with water.
91
What is the test for ammonium ions (NH4+)?
1. add some NaOH(aq) to the unknown substance
2. warm the mixture and hold damp red litmus paper in the mouth of the test tube
If ammonium ions are present, ammonia gas is given off which turns the red litmus paper blue (ammonia gas, NH3, is alkaline)
NH4+(aq) + OH-(aq) --> NH3(g) + H2O(l)
e.g. NH4Cl(aq) + NaOH(aq) --> NH3(g) + H2O(l) + NaCl(aq)
92
Why does the litmus paper need to be damp?
the litmus paper needs to be damp so the ammonia gas can dissolve and make the colour change.
93
What is the test for carbonate ions(CO3 2-)?
1.add some dilute HCl(aq) to the unknown substance
If carbonate ions are present, carbon dioxide is produced (fizzing observed).
CO3 2-(s) + 2H+(aq) --> CO2(g) + H2O(l)
carbonate + acid --> carbon dioxide + water
2. bubble gas through a test tube of limewater, Ca(OH)2, carbon dioxide turns limewater cloudy.
CO2(g) + Ca(OH)2(aq) --> CaCO3(s) + H2O(l)
e.g. CaCO3(s) + HCl(aq) --> CO2(g) + H2O(l) + CaCl2(aq)
94
What is the test for sulfate ions(SO4 2-)?
1. add some dilute HCl, followed by barium chloride solution, BaCl2(aq) (or Ba(NO3)2 )
If sulfate ions are present, a white precipitate of barium sulfate, BaSO4, is produced.
SO4 2-(aq) + Ba2+(aq) --> BaSO4(s)
*barium sulfate is insoluble
e.g. Na2SO4(aq) + BaCl2(aq) --> BaSO4(s) + 2NaCl(aq)
95
Test for halides - what would be the positive result for fluorine (silver fluoride)?
silver fluoride is soluble, so won't give any precipitate.
96
What can be done to avoid false positives and what order should the ion tests be done in to avoid mix-ups?
first add a dilute acid to the test solutions to remove any unwanted anions. (the dilute acid must not interfere with the test e.g. don't use HCl when testing for chloride ions)
test for carbonates ----> no carbon dioxide
test for sulfates ----> no precipitate
test for halides
97
