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The periodic table Flashcards

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1
Q

the number of protons in an atom is the same as the…

A

number of electrons and is called the atomic number

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2
Q

what is the name for the configuration and distance that the electrons are away from the nucleus

A

electronic configuration

and shells

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3
Q

are elements on the periodic table listed in order of increasing or decreasing atomic numbers

A

increasing

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4
Q

what are the horizontal rows called on the periodic table

A

periods

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5
Q

what are the vertical rows called on the periodic table

A

groups

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6
Q

what is the group number the same as

A

the number of electrons in the outermost shell

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7
Q

elements in the same group have similar what

A

chemical properties, because they have the same number of electrons on their outer shell

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8
Q

what is group 1 called

A

alkali metals

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9
Q

why is group 1 called alkali metals

A

because they all react with water to form alkaline solutions

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10
Q

what is group 7 called

A

halogens

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11
Q

why is group 7 called halogens

A

it means ‘salt maker’, halogens form salts when they react with metals

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12
Q

why is group 7 called halogens

A

it means ‘salt maker’, halogens form salts when they react with metals

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13
Q

what is group 8/0 called

A

noble gases

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14
Q

why is group 8/0 called noble gases

A

because they are unreactive (full outer shell)

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15
Q

how does the periodic table change from left to right

A

gradually from metal to non metal

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16
Q

what is another word for a semi-metal

A

a metalloid (it has some metal properties but isn’t actually a metal)

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17
Q

6 metals of group 1

A 
lithium
sodium
potassium
rubidium
caesium
francium (radioactive and little know about its properties)
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18
Q

lithium sodium and potassium physical properties

A
  • good conductors of electricity and of heat
  • soft and can easily be cut with a knife
  • low melting and boiling points compared with more typical metals
  • low densities (all float on water)
  • stored in oil at schools
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19
Q

lithium sodium and potassium chemical properties

A
  • shiny surfaces when freshly cut, but quickly tarnishes as it reacts with oxygen in the air
  • all burn in air or oxygen to form white and solid oxides:
    2M + O2 –> M2O
    M= one of the three metals
  • they react vigorously with water to give an alkaline solution of the metal hydroxide as well as hydrogen gas :
    2M + 2H2O –> 2MOH + H2
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20
Q

observations when lithium is added to water

A
  • moves around the surface of the water
  • hissing sound
  • bubble of gas
  • gets smaller and smaller –> eventually disappears
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21
Q

observations when lithium is added to water

A
  • moves around the surface of the water
  • hissing sound
  • bubble of gas
  • gets smaller and smaller –> eventually disappears
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22
Q

observations when sodium is added to water

A
  • moves around the surface of the water
  • hissing sound
  • bubble of gas
  • gets smaller and smaller –> eventually disappears
  • melts into a shiny ball
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23
Q

observations when potassium is added to water

A
  • moves around the surface of the water
  • hissing sound
  • bubble of gas
  • gets smaller and smaller –> eventually disappears
  • melts into a shiny ball
  • burns with a lilac coloured flame
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24
Q

the further you go down group 1 what should we observe

A

more reactive metals

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25
Q

the further down group 1 less energy is needed to overcome the electrostatic forces of attraction with the ext a outside electron therefore

A
  • francium is the most reactive out of group 1
  • the size of the atom increases and therefore the outer electron gets further away from the nucleus
  • the outer electron is therefore less strongly attracted to the nucleus
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26
Q

the further down group 1 less energy is needed to overcome the electrostatic forces of attraction with the ext a outside electron therefore

A
  • francium is the most reactive out of group 1
  • the size of the atom increases and therefore the outer electron gets further away from the nucleus
  • the outer electron is therefore less strongly attracted to the nucleus
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27
Q

3 most common metals from the halogens

A

chlorine, bromine and iodine (all non metals and poisonous)

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28
Q

state at room temp and colour: chlorine

A

gas

pale green

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29
Q

state at room temp and colour: bromine

A

liquid

red-brown (readily evaporates to form a brown gas)

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30
Q

state at room temp and colour: iodine

A

solid

black (but sublimes when heated to form a purple gas)

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31
Q

how do the three main elements from the halogens react with iron similarly:

A

hot iron wool glows brightly when chlorine passes over it. brown smoke forms and a brown solid is left behind.

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32
Q

how do the three main elements from the halogens react with iron similarly:

A

Chlorine: hot iron wool glows brightly when chlorine passes over it. brown smoke forms and a brown solid is left behind.
Bromine: hot iron wool glows less brightly when bromine vapour passes over it. brown smoke forms and a brown solid is formed.
Iodine: hot iron wool glows even less brightly when iodine vapour passes over it. brown smoke forms and a brown solid is formed.

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33
Q

how do the three main elements from the halogens react with iron similarly:

A

Chlorine: hot iron wool glows brightly when chlorine passes over it. brown smoke forms and a brown solid is left behind.
Bromine: hot iron wool glows less brightly when bromine vapour passes over it. brown smoke forms and a brown solid is formed.
Iodine: hot iron wool glows even less brightly when iodine vapour passes over it. brown smoke forms and a brown solid is formed.

34
Q

what is the reactivity trend down the halogens

A

decreases down the group because of the increasing size of the atoms (attracting an extra atom less attraction of the other atom as there are more shells further away from the nucleus)

35
Q

chlorine reaction with hydrogen

A

a mixture of hydrogen and chlorine explodes when exposed to UV radiation
H2 + Cl2 –> 2HCl

36
Q

bromine reaction with hydrogen

A

a mixture of hydrogen and bromine vapour will react when heated
H2 + Br2 –> 2HBr

37
Q

iodine reaction with hydrogen

A

a mixture of hydrogen and iodine vapour will react when heated, but the reaction does not go completion
H2 + I2 –> 2HI

38
Q

an example of a displacement reaction where chlorine is more reactive than the aqueous solution

A

chlorine + potassium bromide solution –> potassium chloride solution + bromine

39
Q

in general a halogen will displace a … reactive hydrogen for an aqueous solution solution of its halide

A

less

40
Q

in general a halogen will displace a … reactive hydrogen for an aqueous solution solution of its halide

A

less

41
Q

aqueous solutions of the halogen has the following colours: chlorine

A

very pale green, but usually appears colourless since it is often dilute

42
Q

aqueous solutions of the halogen has the following colours: bromine

A

orange, turns yellow when diluted

43
Q

aqueous solutions of the halogen has the following colours: iodine

A

brown

44
Q

in displacement reaction, the halogen molecules is gaining electrons and is therefore being reduced. the ionic half equation for this reaction:

A

X2 + 2e- –> 2X-

45
Q

the halide ions are losing electrons and are therefore being oxidised. redox reaction as both reduction and oxidation are taking place. the ionic half equation for this reaction:

A

2Y- –> Y2 + 2e-

46
Q

HCl dissolved in methylbenzene and in water: effect on blue litmus

A

no change –> not acidic

turns red –> acidic

47
Q

how to show that air contains approx. one fifth oxygen: using iron

A
  • place wet iron filings in the end of a burette and set up the apparatus as shown in the diagram
  • over several days the water will rise up the burette and reach a constant level
  • this is because the iron reacts with the oxygen in the air
  • take the initial and final reading of the water level in the burette
  • the volume of air at the start = (50 - initial burette reading)
  • the percentage of oxygen in air can be calculated using the equation:
    percentage of oxygen = volume of oxygen used/volume of air at start x 100
48
Q

how to show that air contains approx. one fifth oxygen: using phosphorus

A
  • a similar experiment can be done like the one using iron, but with a piece of white phosphorus
  • levels of water inside and outside the tube are equal
  • levels equal; phosphorus touched by hot metal rod
  • levels equal; phosphorus starts burning
  • levels inside higher: phosphorus stops burning
    4P(s) + 5O2(g) –> 2P2O5(s)
49
Q

how to prepare oxygen in a lab

A
  • hydrogen peroxide, H2O2, decomposes slowly to form water and oxygen
  • The speed of the decomposition is increased by adding solid manganese dioxide MnO2, which acts as a catalyst for the reaction
  • the oxygen can also be collected over water
  • since oxygen is not very soluble in water, very little is lost
  • an aqueous solution of hydrogen peroxide is used in the lab prep
    2H2O(aq) –> 2H2O(l) + O2(g)
50
Q

HCl dissolved in methylbenzene and in water: effect of electricity

A

does not conduct –> not contain ions
hydrogen evolved at the negative electrode
chlorine evolved at the positive electrode –> contains ions
HCl(aq) –> H+ (aq) + Cl- (aq)

51
Q

typical properties of hydrochloric acid

A
  • turns litmus paper red
  • has a low pH and hence turns Universal indicator red
  • reacts with fairly reactive metals i.e. magnesium, zinc, iron to form a metal chloride in solution and liberate hydrogen gas
52
Q

typical properties of hydrochloric acid

A
  • turns litmus paper red
  • has a low pH and hence turns Universal indicator red
  • reacts with fairly reactive metals i.e. magnesium, zinc, iron to form a metal chloride in solution and liberate hydrogen gas
  • reacts with metal carbonates to form a metal chloride in solution, water and carbon dioxide gas
  • reacts with bases (metal oxides and metal hydroxides) to form a metal chloride in solution and water
53
Q

typical properties of hydrochloric acid

A
  • turns litmus paper red
  • has a low pH and hence turns Universal indicator red
  • reacts with fairly reactive metals i.e. magnesium, zinc, iron to form a metal chloride in solution and liberate hydrogen gas
  • reacts with metal carbonates to form a metal chloride in solution, water and carbon dioxide gas
  • reacts with bases (metal oxides and metal hydroxides) to form a metal chloride in solution and water
54
Q

both carbon dioxide and sulphur dioxide are … oxides. they dissolved in water to form acidic solutions, give the equations for these reactions:

A

acidic
CO2(g) + H2O(l) –> H2CO3(aq) = carbonic acid (approx pH 5-6)
SO2(g) + H2O(l) –> H2SO3(aq) = sulphurous acid (approx pH 3-4)

55
Q

how to show that air contains approx. one fifth oxygen: using copper

A
  • set up the apparatus with 100cm(cb) of air in one of the gas syringes
  • heat the copper at one end of the silica tube using a blue Bunsen flame
  • pass the air backwards and forwards over the copper
  • as the volume of gas in the syringes decreases, move the Bunsen flame along the tube so that it is always heating fresh copper
  • stop heating when the volume of gas in the syringes stops decreasing, the copper is reacting with the oxygen in the air to form black copper (II) oxide:
    2Cu(s) + O2(g) –> 2CuO(s)
  • the final volume of air in the syringe will be approx. 79 cm(cb) showing that 21cm(cb) has reacted
56
Q

how to show that air contains approx. one fifth oxygen: using iron

A
  • place wet iron filings in the end of a burette and set up the apparatus as shown in the diagram
  • over several days the water will rise up the burette and reach a constant level
  • this is because the iron reacts with the oxygen in the air
  • take the initial and final reading of the water level in the burette
57
Q

how to prepare carbon dioxide in a lab

A
  • the reaction between any metal carbonate and an acid will produce carbon dioxide
  • calcium carbonate is the most commonly used metal carbonate in the lab prep for CO2
  • the most convenient form of calcium carbonate to use is marble chips
  • CaCO3(s) + 2HCl(aq) –> CaCl2(aq) + H2O(l) + CO2(g)
  • carbon diode is not very soluble in water so it van be collected over water without much being lost, it can also be collected downward delivery in air as it is more dense than air
  • carbon dioxide is also produced when most metal carbonates are heater
58
Q

how to prepare oxygen in a lab

A
  • hydrogen peroxide, H2O2, decomposes slowly to form water and oxygen
  • The speed of the decomposition is increased by adding solid manganese dioxide MnO2, which acts as a catalyst for the reaction
  • the oxygen can also be collected over water
  • since oxygen is not very soluble in water, very little is lost
  • an aqueous solution of hydrogen peroxide is used in the lab prep
    2H2O(aq) –> 2H2O(l) + O2(g)
59
Q

observation and equation: magnesium heated in oxygen

A

burns with a bright white flame to form a white powder

2Mg(s) + O2(g) –> 2MgO(s)

60
Q

observation and equation: carbon heated in oxygen

A

burns with a yellow-orange flame to form a colourless gas

C(s) + O2(g) –> CO2(g)

61
Q

observation and equation: sulphur heated in oxygen

A

burns with a blue flame to form a colourless gas

S(s) + O2(g) –> SO2(g)

62
Q

magnesium oxide MgO, is a … oxide. it is very slightly … in water and a sat rated solution with have a pH of about … it reacts with water to form a solution of … Mg(OH)2: give this equation

A 
basic
soluble
10
magnesium hydroxide 
MgO(s) + H2O(l) --> Mg(OH)2(aq)
63
Q

uses of carbon dioxide

A
  • making fizzy/carbonated drinks –> CO2 not usually soluble in water at normal atmospheres but becomes soluble when put under pressure
  • in fire extinguishers (prevents oxygen from getting to the flame), does not support combustion
64
Q

oxides of metals are basic, if they dissolve in water they form what

A

alkaline solutions

65
Q

oxides of non-metals are often … if they dissolve in water they form … solutions. some oxides of non-metals are neutral e.g. carbon monoxide

A

acidic

acidic

66
Q

how to prepare carbon dioxide in a lab

A
  • the reaction between any metal carbonate and an acid will produce carbon dioxide
  • calcium carbonate is the most commonly used metal carbonate in the lab prep for CO2
  • the most convenient form of calcium carbonate to use is marble chips
  • CaCO3(s) + 2HCl(aq) –> CaCl2(aq) + H2O(l) + CO2(g)
67
Q

thermal decomposition (breakdown by heating) observations and equation: copper (II) carbonate

A

green to black

CuCO3(s) –> CuO(s) + CO2(s)

68
Q

thermal decomposition (breakdown by heating) observations and equation: magnesium carbonate

A

stays white

mgCO3(s) –> MgO(s) + CO2(g)

69
Q

thermal decomposition (breakdown by heating) observations and equation: calcium carbonate

A

stays white

CaCO3(s) –> CaO(s) + CO2(g)

70
Q

thermal decomposition (breakdown by heating) observations and equation: zinc carbonate

A

white to yellow when hot and white again when cold

ZnCO3(s) –> ZnO(s) + CO2(g)

71
Q

thermal decomposition (breakdown by heating) observations and equation: sodium carbonate

A

stays white

doesn’t decompose in a Bunsen flame

72
Q

uses of carbon dioxide

A
  • making fizzy/carbonated drinks

- in fire extinguishers (prevents oxygen from getting to the flame)

73
Q

metals … hydrogen in the reactivity series react with both dilute hydrochloric acid and dilute sulphuric ace to form a …

A

above

salt and hydrogen

74
Q

metal + dilute hydrochloric acid –> ? + hydrogen

metal + dilute sulphuric acid –> ? + hydrogen

A

metal chloride

metal sulphate

75
Q

observation of reaction with HCl or H2SO4 and equations: magnesium

A
  • bubble of gas
  • magnesium disappears
  • reaction mixture gets hot
  • colourless solution formed
    Mg (s) + 2HCl(aq) –> MgCl2(aq) + H2(g)
    Mg(s) + H2SO4(aq) –> MgSO4(aq) + H2(g)
76
Q

observation of reaction with HCl or H2SO4 and equations: aluminium

A
  • slow to start reacting when cold, but bubbles form when heated
  • aluminium disappears
  • colourless solution formed
    2Al(s) + 6HCl(aq) –> 2AlCl3(aq) + 3H2(g)
    22Al9s) + 3H2SO49aq) –> Al2(SO4)3(aq) + 3H2(g)
77
Q

observation of reaction with HCl or H2SO4 and equations: zinc

A
  • bubbles of gas
  • zinc disappears
  • colourless solution formed
    Zn(s) + 2HCl(aq) –> ZnCl2(aq) + H2(g)
    Zn(s) + H2SO4(aq) –> ZnSO4(aq) + H2(g)
78
Q

observation of reaction with HCl or H2SO4 and equations: iron

A
  • bubbles of gas
  • iron disappears
  • pale green solution formed
    Fe(s) + 2HCl(aq) –> FeCl2(aq) + H2(g)
    Fe(s) + H2SO4(aq) –> FeSO4(aq) + H2(g)
79
Q

hydrogen burns when heated un air or oxygen to form…

A

water

2H2(g) + O2(g) –> 2H2O(l)

80
Q

how to test for hydrogen

A

a mixture of hydrogen and air/oxygen will explode when ignited by a spark of a flame and water is the product. this provides the characteristic test for hydrogen. mix a test tube full of hydrogen with air and place the lame of a lighted spill at the mouth of the test tube, hear a squeaky pop

81
Q

2 tests for pure water

A
  1. add the liquid from condensing the water vapour to anhydrous copper(II) sulphate. the white powder will turn blue and hydrated copper (II) sulphate is formed:
    CuSO4(s) + 5H2O(l) –> CuSO4.5H2O(s)
    white –> blue
  2. boil the liquid and measure its boiling point (100 degrees Celsius).
    theses two tests together confirm that the liquid is pure water. if the water contains impurities, its boiling point will be above 100 degrees celcius and its freezing point will be less than 0 degrees celcius

Chemistry (11 decks)

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