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Chemistry Unit 1 > 1.6 - the periodic table > Flashcards

1.6 - the periodic table Flashcards

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1
Q

Who discovered the periodic table?

A

Mendeelev 1869
left gaps for undiscovered elements

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2
Q

what do the blocks indicate?

A

the presence of the valence electron

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3
Q

what increases across the period?

A

ionisation increases across the period
nuclear charge increases across the period

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4
Q

what doesn’t increase across the period?

A

shielding

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5
Q

As Ionisation increases, atomic radius?

A

decreases

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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

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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

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7
Q

Difference between singularly occupied orbita anda paired orbital?

A

easier to remove it

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8
Q

What decreases down the group?

A

Ionisation decreases
Increase in shielding outweighs the increase in nuclear charge

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9
Q

Electronegativity def?

A

measure of tendency of an atom to attract a bonding pair of electrons

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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

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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

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12
Q

Where are the more electronegative elements?

A

at the top of the RHS

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13
Q

Where are the least electronegative elements?

A

bottom of LHS

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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

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15
Q

What changes across a period?

A

the structure of the elements from metallic to giant covalent then to simple molecular

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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

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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

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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

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19
Q

Ar?

A

lowest melting + boiling temp because it exists as seperate atoms that are held together by induced dipole induced dipole
monoatomic

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20
Q

Trends in melting + boiling temps?

A

similar in period 2 but boron in Group 3 has a giant covalent non metallic structure

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21
Q

Reduction + Oxidation?

A

Redox takes place together

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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)

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23
Q

Shown by?

A

Mg — Mg2+ + e-
Cu2+ + 2e- —– Cu

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24
Q

Oxidation def?

A

loss of electrons

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25
Q

Reduction def?

A

gain of electrons

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26
Q

Mg?

A

reducing agent and is oxidised in the process

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27
Q

What is another way to tell if a reaction = redox?

A

to work out the oxidation numbers

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28
Q

E.g Ba+Cl2 —– Bacl2

A

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

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29
Q

Group 1 metals with water ?

A

Group 1metals react vigorously with cold water to form hydroxide + hydrogen

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30
Q

for example?

A

2Na(g)+2H2O(l) —— 2 NaOH(aq)+H2(g)
2Li(g)+2H2O(l)—— 2 LIOH(aq) + H2(g)

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31
Q

Reactions increase in vigour as you go down the group

A

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

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32
Q

Group 2 metals?

A

reach much less vigorously in fact
Mg reacts very slowly. Again the OH + H = formed
Ca+2H2O—-Ca(OH)2+H2

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33
Q

Reactivity increase as you go down the group e.g?

A

Ca produces a steady stream of bubbles + liquid goes cloudy as a white ppt of CaOH = Formed

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34
Q

Ba?

A

produces greater effervesence + solution is clearer since barium Hydroxide = more soluble is colourless

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35
Q

Hydroxides?

A

more soluble in Group 2

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36
Q

Does magnesium react with water?

A

cold water no only steam
mg+H2O(g) —MgO+H2

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37
Q

Hydrogen test?

A

Place a lit splint in an inverted test tube of gas
If it makes a squeaky pop, its H

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38
Q

Why does reactivity increase as you go down the group?

A

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
Q

What does it lead to?

A

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
Q

Reaction with acids?

A

all group 2 metals react vigorously with HCL to produce a colourless solution of the metal chloride + bubbles of Hydrogen

41
Q

Equations?

A

2 Li(s) + 2HCL(aq) —– 2Licl(aq) +H2(g)
Sr+2HCl—-SrCl2+H2
Ba(s)+2HCl—–Bacl2+H2

42
Q

What is BaCl test for?

A

Sulfate ions

43
Q

Reactivity of Group 2 metals?

A

increase as you go down the group
Mg reacts with H2SO4 as the other memers have insoluble sulfates

44
Q

Group 1 metals?

A

too reactive to be added directly to acids
2Li(s) + Cl(aq) —– 2 LiCl(aq)+H2(g)

45
Q

Reacts with Oxygen?

A

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
Q

Group 1 metals also burn?

A

white solids and burn with a characteristic flame
4Li + O2 —- 2 Li2O

47
Q

What do Group 1 metals also form?

A

white solids + burn with a characteristic flame
4Li+O2–2Li2O
They also form peroxides + superoxides

48
Q

Oxides + Hydroxides?

A

Metal oxides = basic
non metal oxides = acidic

49
Q

All S block metal oxides =?

A

strong bases and they neutralise acids to form a salt and water
MgO + 2 HCl —- MgCl2 + H2O

50
Q

Group 1 oxides + barium oxides ?

A

react with water to form a soluble Hydroxide
e.g
Na2O+H2O —- 2NaOH

51
Q

Since Hyroxides?

A

soluble - alkalis

52
Q

Other Group 2 hydroxides?

A

not very soluble so saturated solutions of these hydroxides are only weakly basic because concentration of Hydroxide ions = very low

53
Q

Test for cations?

A

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
Q

Li?

A

red

55
Q

Na+

A

Orange - yellow

56
Q

K+?

A

lilac

57
Q

Mg2+?

A

no colour

58
Q

Ca2+?

A

brick red

59
Q

Sr2+?

A

crimson

60
Q

Ba2+?

A

apple green

61
Q

Solubility in water?

A

all Group 1 compounds = soluble
However, many Group 2 compounds are not

62
Q

Group 2 trends?

A

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
Q

Equation?

A

Ba2+(aq) + SO42-(aq) —— BaSo4(s)
trends can be used to distinguish between unkown solutions containing group 2 cations

64
Q

Thermal of solubilities of Hydroxides + carbonates

A

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
Q

Thermal stabilty?

A

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
Q

Chemistry of G7 halogens + halides?

A

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
Q

At room temp?

A

Chlorine = a green gas
Bromine = red Brown liquid
Iodine = grey solid

68
Q

As the number of electrons increase with atomic number?

A

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
Q

Volatile def?

A

substances that form vapours easily

70
Q

Volatility?

A

substances with a low boiling temp has a high volatility
decreases down the group

71
Q

Trends in reacitivity ?

A

halogens react by gaining elecrons to form negative halide ions. Since they gain electrons during reactions, halogens = reduced + they oxidise the other substance

72
Q

As you go down the group?

A

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
Q

How are halides form?

A

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
Q

Iron wool?

A

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
Q

Equations?

A

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
Q

Displacement reactions

A

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
Q

What happens when these displacement reactions happen?

A

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
Q

Test for Halide ions?

A

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
Q

Silver nitrate gives?

A

Cl - white ppt
Br - cream ppt
I - yellow ppt

80
Q

Precipitate?

A

insoluble in silver halide
Ag +(aq) + cl - (aq)
—— AgCl(s)

81
Q

AgCl?

A

ppt dissolves in dilute NH3

82
Q

AgBr?

A

ppt does not dissolve much in dilute NH3 but does dissolve in concentrated NH3

83
Q

AgI?

A

ppt insoluble in dilute + concentrated NH3

84
Q

Uses of chlorine + fluoride in water + treatment?

A

Chlorine = commonly added to water as the gaseous element + equilibrium = established
Cl2+H2O —— HCl + HOCl

85
Q

ClO-?

A

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
Q

Risks of Chlorine?

A

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
Q

Why do some people object to water chlorination?

A

as they are forced due to mass medication

88
Q

Fluorine?

A

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
Q

Why do many people object to fluoridation?

A

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
Q

Practical activity?

A

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
Q

Steps?

A

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
Q

Insoluble salt formation?

A

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
Q

Steps?

A

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
Q

Gravimetric anylasis?

A

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
Q

for example?

A

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

Chemistry Unit 1 (7 decks)

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