1.6 - the periodic table Flashcards
(96 cards)
1
Q
Who discovered the periodic table?
A
Mendeelev 1869
left gaps for undiscovered elements
2
Q
what do the blocks indicate?
A
the presence of the valence electron
3
Q
what increases across the period?
A
ionisation increases across the period
nuclear charge increases across the period
4
Q
what doesn’t increase across the period?
A
shielding
5
Q
As Ionisation increases, atomic radius?
A
decreases
6
Q
Between Group 2 and 3?
A
there’s a decrease in Ionisation ebergy because Group 3’s valance electron is in a new subshell with a higher energy level shielded by the S electrons
7
Q
Between Group 5 and Group 6?
A
there’s a decrease in Ionisation energy
there’s a change between N and O
N = singly occupied
O = pair
7
Q
Difference between singularly occupied orbita anda paired orbital?
A
easier to remove it
8
Q
What decreases down the group?
A
Ionisation decreases
Increase in shielding outweighs the increase in nuclear charge
9
Q
Electronegativity def?
A
measure of tendency of an atom to attract a bonding pair of electrons
10
Q
What increases across a period?
A
Electronegativity
There is an increase in nuclear charge but the bonding electron is always shielded by the same inner electrons
so there is a greater attraction betweenthe nucleus and the bonding pair
11
Q
When does electronegativity decrease?
A
down the group
bonding electrons have increased shielding from nucleus, so attraction between the nucleus and the bonding electrons decrease
12
Q
Where are the more electronegative elements?
A
at the top of the RHS
13
Q
Where are the least electronegative elements?
A
bottom of LHS
14
Q
What is the pattern?
A
there is a general increase from first to 4th electron
a large decrease to the 5th element then a small general decrease to the 8th element
15
Q
What changes across a period?
A
the structure of the elements from metallic to giant covalent then to simple molecular
16
Q
What bonding does Sodium, magnesum, aluminium have?
A
metallic bonding
increase because metallic bonding gets stronger
metal ions have a greater charge + there is an increased number of delocalised electrons
17
Q
Silicon?
A
Giant covalent structure
each atom = bonded covalently to 4 other atoms
a large amount of energy is needed to break all these bonds
18
Q
Phosphorous, Sulfur+ chlorine?
A
simple molecular
althought covalent bonds between toms = strong IMF holding these molecules together = weak + dont need much more energy to break
weak vdw
19
Q
Ar?
A
lowest melting + boiling temp because it exists as seperate atoms that are held together by induced dipole induced dipole
monoatomic
20
Q
Trends in melting + boiling temps?
A
similar in period 2 but boron in Group 3 has a giant covalent non metallic structure
21
Q
Reduction + Oxidation?
A
Redox takes place together
22
Q
Mg + CuO —— MgO + Cu?
A
Mg has gained oxygen so = oxidised
(increase in oxidation state)
Copper oxide has lost oxygen so = reduced
(decrease in oxidation state)
23
Q
Shown by?
A
Mg — Mg2+ + e-
Cu2+ + 2e- —– Cu
24
Oxidation def?
loss of electrons
25
Reduction def?
gain of electrons
26
Mg?
reducing agent and is oxidised in the process
27
What is another way to tell if a reaction = redox?
to work out the oxidation numbers
28
E.g Ba+Cl2 ----- Bacl2
Oxidation number of Barium decreases from 0 to + 2
therefore have been oxidised
Oxidation number of Chlorine has decreased from 0 to -1 therefore have been reduced
Oxidation number of an atom does not always exchange when it reacts.
It can be helpful to where the oxidation numbers of each atom underneath symbol
2HNO3 + 6 Hi ------ 2NO + 3 I2+H20
Nitrogen = reduced, oxidation number decrease from + 5 to + 2
Iodine is oxidised from -1 to 0
29
Group 1 metals with water ?
Group 1metals react vigorously with cold water to form hydroxide + hydrogen
30
for example?
2Na(g)+2H2O(l) ------ 2 NaOH(aq)+H2(g)
2Li(g)+2H2O(l)------ 2 LIOH(aq) + H2(g)
31
Reactions increase in vigour as you go down the group
Li floats on water, gently fizzing
Na melts into a ball that dashes around the surface
K melts into a ball + catches on fire
Cs explodes + shatters glass container
32
Group 2 metals?
reach much less vigorously in fact
Mg reacts very slowly. Again the OH + H = formed
Ca+2H2O----Ca(OH)2+H2
33
Reactivity increase as you go down the group e.g?
Ca produces a steady stream of bubbles + liquid goes cloudy as a white ppt of CaOH = Formed
34
Ba?
produces greater effervesence + solution is clearer since barium Hydroxide = more soluble is colourless
35
Hydroxides?
more soluble in Group 2
36
Does magnesium react with water?
cold water no only steam
mg+H2O(g) ---MgO+H2
37
Hydrogen test?
Place a lit splint in an inverted test tube of gas
If it makes a squeaky pop, its H
38
Why does reactivity increase as you go down the group?
when the S-block metals react, they lose electrons to form + ions
since IE decrease down the group, the energy needed to form ions decrease
39
What does it lead to?
lower activation energies + faster reactions
Group 1 metals = more reactive than Group 2 metals because Group 1 metals lose only 1 electron while Group 2 metals lose 2 electrons
40
Reaction with acids?
all group 2 metals react vigorously with HCL to produce a colourless solution of the metal chloride + bubbles of Hydrogen
41
Equations?
2 Li(s) + 2HCL(aq) ----- 2Licl(aq) +H2(g)
Sr+2HCl----SrCl2+H2
Ba(s)+2HCl-----Bacl2+H2
42
What is BaCl test for?
Sulfate ions
43
Reactivity of Group 2 metals?
increase as you go down the group
Mg reacts with H2SO4 as the other memers have insoluble sulfates
44
Group 1 metals?
too reactive to be added directly to acids
2Li(s) + Cl(aq) ----- 2 LiCl(aq)+H2(g)
45
Reacts with Oxygen?
apart from Mg, all Group 2 metals + end to burn with a characteristic glame
All group 2 metals burn to form solide white oxides
2 Mg+O2---- 2MgO
46
Group 1 metals also burn?
white solids and burn with a characteristic flame
4Li + O2 ---- 2 Li2O
47
What do Group 1 metals also form?
white solids + burn with a characteristic flame
4Li+O2--2Li2O
They also form peroxides + superoxides
48
Oxides + Hydroxides?
Metal oxides = basic
non metal oxides = acidic
49
All S block metal oxides =?
strong bases and they neutralise acids to form a salt and water
MgO + 2 HCl ---- MgCl2 + H2O
50
Group 1 oxides + barium oxides ?
react with water to form a soluble Hydroxide
e.g
Na2O+H2O ---- 2NaOH
51
Since Hyroxides?
soluble - alkalis
52
Other Group 2 hydroxides?
not very soluble so saturated solutions of these hydroxides are only weakly basic because concentration of Hydroxide ions = very low
53
Test for cations?
All S-block elements apart from Mg may be identified by a flame test
a clean metal wire is moistoned with HCl, dipped in compound + held in anion - luminous Bunsen flame
54
Li?
red
55
Na+
Orange - yellow
56
K+?
lilac
57
Mg2+?
no colour
58
Ca2+?
brick red
59
Sr2+?
crimson
60
Ba2+?
apple green
61
Solubility in water?
all Group 1 compounds = soluble
However, many Group 2 compounds are not
62
Group 2 trends?
All nitrates = soluble
All carbonates are insoluble
Hydroxides become more soluble as you go down the group. Therefore Magnesium Hydroxide = insoluble whilst Barium Hydroxides are soluble
Mg2+(aq)+2OH-(aq) ----- Mg(OH)2(s)
sulfates become less soluble as you go down the group.
Therefore Magnesium sulfate = soluble but barium sulfate = insoluble
63
Equation?
Ba2+(aq) + SO42-(aq) ------ BaSo4(s)
trends can be used to distinguish between unkown solutions containing group 2 cations
64
Thermal of solubilities of Hydroxides + carbonates
All Group 2 hydroxides decomposee on heating to the oxide and steam
Ca(OH)2(s) ----- Ca(O)(s) + H2O(g)
Thermal stabilities increase as you go down the group i.e the hydroxides have to be heated more strongly before they decompose
All group 2 carbonates decompose on heating to the oxide and CO2
e.g MgCO3(s)---- MgO(s)+CO2(g)
65
Thermal stabilty?
Increase as you go down the group
Shown in the lab by heating the carbonate and seeing how long the CO2 formed to turn limewater cloudy
66
Chemistry of G7 halogens + halides?
Elements known as halogens because they all form salts called halides
Halogen = salt producer
Halogens exist as 2 molecules containing a single covalent bond e.g Cl2
67
At room temp?
Chlorine = a green gas
Bromine = red Brown liquid
Iodine = grey solid
68
As the number of electrons increase with atomic number?
increased in the induced dipole - induced dipole intermolecular forces holding the diatomic molecule
therefore,the melting + boiling temp increase as you go down the group
69
Volatile def?
substances that form vapours easily
70
Volatility?
substances with a low boiling temp has a high volatility
decreases down the group
71
Trends in reacitivity ?
halogens react by gaining elecrons to form negative halide ions. Since they gain electrons during reactions, halogens = reduced + they oxidise the other substance
72
As you go down the group?
the outer electrons = shielded more + are further from the nucleus. so it gets harder to attract electrons + both reactivity + oxidising power down the group
73
How are halides form?
halogens react directly with most metals to form the halide
for example
Sodium burns in a jar of Cl2 gas whilst forming white Nacl
2 Nacl+cl2 ---- 2 Nacl
74
Iron wool?
burns directly in Cl2 or Br2 vapour to gie the iron III halide
however, when burns in iodine vapour, it produces Iron (II) iodide since Iodine is less reactive + is a weaker oxidisng agent
75
Equations?
2Fe+3Br2 ----- 2 FeBr3
Fe+I2----FeI2
Bromine oxidises iron to the +3 oxidation state
Iodine oxidises the iron to the + 2 oxidation state
76
Displacement reactions
halogen in a higher position in the group will oxidise a halide ion from lower in the group
oxidising powers decrease down the group
when a halogen is added to an aqueous solution containing a halide ion
chloride displaces bromide + iodide
bromine displaces only iodide
iodine does not displace either chloride or bromide
77
What happens when these displacement reactions happen?
colour change
e.g when chlorine water = mixed with potassium Bromide, solution changes from colourless to change
since chlorine has oxidised bromide ions
when chlorine mixed with potassium iodine, solution changes from colourless to brown, since chlorine has oxidised from iodide to iodine
78
Test for Halide ions?
Silver nitrate test
test has to be done in solution
if it starts from a solid, must first be dissolved in water
few drops of nitric acid = added first to make sure that any other anions is removed as they would also form ppt
79
Silver nitrate gives?
Cl - white ppt
Br - cream ppt
I - yellow ppt
80
Precipitate?
insoluble in silver halide
Ag +(aq) + cl - (aq)
------ AgCl(s)
81
AgCl?
ppt dissolves in dilute NH3
82
AgBr?
ppt does not dissolve much in dilute NH3 but does dissolve in concentrated NH3
83
AgI?
ppt insoluble in dilute + concentrated NH3
84
Uses of chlorine + fluoride in water + treatment?
Chlorine = commonly added to water as the gaseous element + equilibrium = established
Cl2+H2O ------ HCl + HOCl
85
ClO-?
kills bacteria + other microbes , adding chlorine makes it safe to drink
Chlorination = also used to prevent outbreak of serious disease such as typhoid + chlorea
86
Risks of Chlorine?
risks in using Chlorine to treat water
Highly toxic because naturally organic compounds found in the water supply to form chlorine hydrocarbons which cause liver+ kidney cancer
87
Why do some people object to water chlorination?
as they are forced due to mass medication
88
Fluorine?
generally added to water to reduce tooth decay by preventing cavities
water fluoridation reduces cavities in children but its effectiveness in adults is less clear
although fluoridation can cause dental fluorosis which leads to tooth discolouration
no clear evidence of other adverse effects from water fluoridation
only beneficial effects below 1 nanometres
89
Why do many people object to fluoridation?
mass medication
fluoride in toothpaste, mouth rinses to other dental products, many people think adding fluoride to water supplies = detrimental to long term health
90
Practical activity?
Soluble salt formation
Copper(II) sulfate can be formed by neutralising sulfuric acid with the insoluble base copper(II) oxide
H2SO4(aq) + CuO ----- CuSO4(aq) + H2O(l)
91
Steps?
1) some copper(II) oxide is added to dilute H2SO4 more = added until no more dissolves
solution turns blue
2) As the acid has been used up
excess solid = removed by heating
this blue solution of Copper(II) sulphate in water
3) solution is heated to evaporate some of the water
4) It is left to cool. Blue crystals of Copper(II) sulfate starts to form
the water should not be fully evaporated because if this happens, a powder will form = rather than crystals
If copper (II) carbonate is used, method = exactly the same but effervescence formed is seen when the carbonate is added to the acid because CO2 is given off
when no more effervescence is seen acid is used up
92
Insoluble salt formation?
as insoluble salt can be made using a ppt reaction by reacting 2 soluble solutions.
In a ppt reaction, the positive ions + negative ions in the 2 solutions switch to form 2 new compounds
1 insoluble + 1 soluble salt
CaCO3 can be formed from CaNO3 + NaCl
Ca(NO3)2 (aq) + Na2 CO3(aq) ---- CaCO
93
Steps?
1)seperately dissolve sodium carbonate + calcium nitrate in water + mix them together using a stirring rod in a beaker
2)filter to remove ppt from mixture
wash ppt with water to remove traces of solutions
3)leave in an oven to dry
94
Gravimetric anylasis?
technique through which the amount of an anylatic (the ion being anylased can be determined through the measurement of mass
depends on through the measurement of mass
depends on comparing masses of 2 compounds
containing the anylate
mass of an ion in a pure compound can be determined then used to find mass percentage of the same in a known quantity of an impure compound
95
for example?
determination of chlorine in a compound
silver chloride = insoluble ion + can be formed pure + easily filtered, a soluble salt can be used to determination the percentage of a chloride
