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Flashcards in module 3 Deck (94):
1

what theories were there before the periodic table?

ancient philosophers such as Aristotle believed the world was made up of four elements- earth, water, air and fire (similar to the states of matter- solid, liquid and gas)

2

what did the chemist Antoine-Laurent de Lavoisier produce?

the first modern chemical textbook in which he compiled the first extensive list of elements, which he described as 'substances that could no be broken down further' .

3

in the early 1800s what were the two ways of categorising elements?

by their physical and chemical properties, and by their relative atomic mass. (protons and electrons were not yet discovered)

4

who was responsible for introducing letter based symbols fo elements?

Jons Jacob Berzelius

5

what did Johann Dobereiner discover?

he noticed that certain groups of three elements, which he called triads, ordered by atomic weight would have a middle element with a weight and properties that were roughly an average of the other two elemants.
eg- chlorine, bromine and iodine.

6

who first attempted to make the table of elements and how were the elements arranged?

John Newlands arranged the elements in order of their relative atomic masses. he suggested that, rather than being in triads, every eighth element showed similar properties, which he named the law of octaves.

7

what were the issues with Newlands' version of the periodic table?

- newlands left no gaps for undiscovered elements, and when they were found this seemed to undermine his suggestions.

- even though he listed the elements in rows of seven so that similar elements lined up in columns, the pattern broke on the third row with transitions metals not following the trend.

8

how did Dmitri Mendeleev arrange the elements in his periodic table and why was his version accepted?

he arranged the known elements by atomic mass. elements with similar properties were arranged in vertical columns and gaps were left where no elements fitted the repeating pattern, which meant he could predict the properties of the undiscovered elements.

- when elements were later discovered, their properties matched Mendeleev's predictions.

9

which elements did Mendeleev leave spaces for in his periodic table?

gallium, scandium and germanium

10

how did Henry Moseley contribute towards the periodic table and what 3 pairs of elements were put in the correct order.?

he arranged the elements by increasing atomic number instead of relative atomic mass.

- his modified periodic law put elements such as argon and potassium, tellurium and iodine, and cobalt and nickel in the correct order.

11

what did Glenn Seaborg discover?

the transuranic elements from Plutonium to Nobelium.
he also remodelled the periodic table by placing the actinide series below the lanthanide series.

12

how is the modern periodic table arranged?

arranged by increasing atomic (proton) numbers, into periods and groups.

13

periodicity

the repeating trends in the physical and chemical properties of the elements across each period.

14

what do all elements within a period or a group have?

- all the elements within a period have the same number of electron shells.

- all the elements within a group have the same number of electrons in their outer shell, and therefore the same chemical properties.
(caused by the repeating pattern of electron configuration).

15

what are semi metals/ metalloids?

non metal elements that have properties between those of a metal and those of a non metal.

16

explain how electron configurations change across a period and down a group

across each group, for each successive element, there is one more electron in the outer shell.

each element in a group has the same outer shell configuration, but with each successive element in a group an extra shell is used.

17

within a shell, in what order do the sub shell energy levels increase?

s < p < d < f

18

why does the 4s orbital fill before the 3d orbital and what effect does this have during ionisation?

because the 4s energy level is lower than the 3d energy level.
the 4s orbital would be emptied before the 3d orbital during ionisation.

19

first ionisation energy

the energy needed to remove one mole of electrons from one mole of gaseous atoms.

20

is ionisation an endothermic or exothermic reaction?

endothermic (energy is put in to ionise an atom/molecule)

21

what has to be done to form a positive ion?

energy must be supplied to an electron in order to overcome the attraction between negative electrons and the positive nucleus

22

why are electrons in the outer shell removed first?

because they experience the smallest nuclear attraction. the outer shell electrons are the furthest away from the nucleus and require the least ionisation energy.

23

what factors does the nuclear attraction experienced by an electron depend on

atomic radius
nuclear charge
electron shielding/ screening

24

nuclear charge

the more protons there are in the nucleus, the more positively charged the nucleus is and the stronger the attraction for the electrons

25

atomic charge

the larger the atomic radius, the smaller the nuclear attraction experienced by the outer electrons, because the positive charge of the nucleus is further away from the outermost electrons.

26

electron shielding

the inner shell electrons repel the outer shell electrons, so a smaller nuclear attraction is experienced by the outer electrons.

27

what does a high ionisation energy mean?

theres a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron.

28

successive ionisation energy

a measure of the amount of energy required to remove each electron in turn.

29

how many ionisation energies does an element have?

it has as many ionisation energies as it has electrons.

30

give the successive ionisation energies of lithium

li(g) → li+ (g) + e-
li+(g) → li2+ (g) + e-
li2+(g) → li3+ (g) + e-

31

why do successive ionisation energies increase within each shell?

electrons are being removed from an increasingly positive ion and so outweigh the negative charge.
theres also less repulsion amongst the remaining electrons and each shell is drawn in slightly closer to the nucleus. as the distance of each electron from the nucleus decreases, the nuclear attraction increases and more energy energy is needed to remove each successive electron.

32

what causes the sudden increases in ionisation energy?

when an electron is removed from a different shell there is a big increase in the energy required (as that shell is closer to the nucleus)

33

why do noble gases have high ionisation energies?

they have a full outer shell of electrons and a high positive attraction from the nucleus.

34

explain the general trends in ionisation energy across a period. (nuclear charge/ distance/ shielding)

Ionisation energy shows a general increase across each period.
Across each period, the number of protons increases, so there is more nuclear attraction acting on the electrons. Electrons are added to the same shell, so the nuclear attraction draws the outer shell inwards slightly.
There is the same number of inner shells and so electron shielding will hardly change.
The increased nuclear charge is the most significant factor

35

explain the exceptions to the trend of increasing ionisation energy across a period,

There are small decreases in first ionisation energy seen between groups 2 and 3 and again between groups 5 and
6. Between groups 2 and 3 this is due to the outermost electron being in a p-orbital, which is at a slightly higher energy level than the s-orbital used by group 2 outer electrons, and so easier to remove. Between groups 5 and 6, this is due to the outermost electron being spin-paired and experiencing some repulsion, which makes it easier to remove.

36

explain the trends in ionisation energy moving down a group period. (nuclear charge/ distance/ shielding)

Ionisation energy shows a general decrease down each group.
The number of shells increases and the distance of the outer electrons from the nucleus increases, increasing the atomic radius so there is a weaker force of attraction on the outer electrons.
There are more inner shells, so the electrons are more effectively shielded from the nuclear charge, resulting in less attraction. The number of protons in the nucleus also increases, but the resulting increased attraction is outweighed by the increase in distance between the nucleus and the outer electrons (i.e. the resulting atomic radius) and shielding.

37

what causes a sharp decrease in first ionisation energy between the end of one period and the start of the next?

the addition of a new shell, further from the nucleus. this leads to an increase in the distance of the outermost shell from the nucleus and in electron shielding by inner electrons.

38

what are atoms in a solid metal held together by?

metallic bonding

39

metallic bonding

the strong electrostatic attraction between positive ions and negatively charged delocalised electrons. the positive ions attract all the electrons as a whole.

40

giant metallic lattice structure

the electrons in the outermost shell of a metal atom are delocalised, leaving a positively charged metal cation.
there is strong electrostatic attraction between positive ions and negatively charged delocalised electrons.

41

explain why metals can conduct electricity

the delocalised electrons can move freely within the metallic structure, allowing the metal to conduct electricity.

42

explain why metals have high melting snd boiling points

all the positive ions present in the lattice attract the negative delocalised electrons, so a large amount of energy is needed to overcome this strong attraction.

43

why are metals malleable and ductile?

there are no bonds holding specific ions together, so the metal ions can slide past each other, resulting in metals being malleable (can be hammered into sheets) and ductile (can be drawn into a wire).

44

explain how metallic bonding is different from covalent bonding

In metallic bonding, the electrons are delocalised and its impossible to see which electrons belong to which positive ions. They are able to move freely within the lattice.

In a covalent bond, the electrons are paired and shared between two atoms and the bonded pair of electrons is localised – they cannot move freely.

Both giant covalent and giant metallic lattices have high melting and boiling points, but only metals (with metallic bonding) can conduct electricity.

45

as you move across periods 2 and 3, how do the elements change? (state/ structure/ forces/ bonding)

they change from metal elements to nonmetal elements and from solids to gases.

structure: giant metallic-- giant covalent-- simple molecular.

forces: strong forces between positive ions and negative delocalised electrons-- strong forces between atoms-- weak intermolecular forces between molecules.

bonding: metallic-- covalent-- covalent bonding within molecules/ intermolecular bonding between molecules.

46

what is silicon classified as and why?

si is classified as a metalloid; it has the shiny appearance of a metal, but is brittle and it conducts electricity, but very poorly.

47

explain the changes of state from solid to gas and the melting points moving across periods 2 and 3.

From group 1 to group 14, atoms of each element are held together in giant structures by strong bonds, either metallic or covalent, so the melting points increase.

From group 15, the elements have simple molecular structures and decreasing melting points. Individual molecules are held together by weak intermolecular forces, resulting in them being gaseous at room temperature.

48

what is the effect on metallic bonding from extra electrons and greater nuclear charge, from Na to Mg to Al? (across period 2)

the ionic charge increases
the ionic size decreases
the number of outer shell electrons increases and attraction increases, so melting and boiling points also increase.

49

explain why the melting point of aluminium is higher than that of magnesium

The metallic structure of aluminium is made from Al3+ ions and delocalised electrons and the metallic structure of magnesium is made from Mg2+ ions and delocalised electrons. Al3+ ions are smaller than Mg2+ ions. In aluminium, the attraction between the smaller, more positively charged ions and the greater number of delocalised electrons means that a larger amount of energy is required to break these strong metallic bonds than those in magnesium, resulting in a higher melting point.

50

explain why the melting point of carbon is much higher than that of nitrogen

Carbon has a giant covalent structure with strong covalent bonds between the atoms in this structure. So, a large amount of energy is required to break these strong forces, resulting in a high melting point. Nitrogen has a simple molecular structure with weak van der Waals’ forces between the N2 molecules. So, a small amount of energy is required to break these weak forces, resulting in a low melting point.

51

allotropes

different forms of the same element in the same state

52

give examples of giant covalent lattices

allotropes of carbon: diamond, graphite, graphene

silicon

53

describe and explain the structure of diamond

diamond has a tetrahedral structure where each carbon is bonded to four other carbon atoms.
this makes it the extremely hard with a very high melting point.
- its a good thermal conductor because vibrations travel easily through the stiff lattice.

54

uses of diamond

diamond-tipped drills and saws

55

describe and explain the structure/properties of graphite
(uses)

graphite has a layered structure where each carbon atom is bonded to three other carbon atoms and the fourth outer electron of each carbon atom is delocalised.

the weak forces between the layers in graphite are easily broken so the layers can slide over each other, making graphite soft and slippery. (used as a dry lubricant and in pencil led)

56

which is less dense- graphite or diamond and why?

graphite is less dense than diamond because the layers are far apart compared to the length of the covalent bonds, so its used to make strong, lightweight sports equipment.

57

describe the structure of silicon

silicon forms a crystal lattice (tetrahedral) structure with similar properties to carbon. each silicon atom forms four strong covalent bonds.

58

describe the structure and properties of graphene

graphene forms interlocking hexagonal rings that make up a lattice only one atom thick.

- the delocalised electrons are free to move along the sheet and can move quickly without layers, making graphene the best known electrical conductor.

the delocalised electrons strengthen the covalent bonds between the carbon atoms, making graphene extremely strong, and its single layer makes it transparent and very light.

59

uses of graphene

- applications in high-speed electronics and aircraft technology due to its high strength, low mass and good electrical conductivity.

- touchscreens on electronic devices due to its flexibility and transparency.

60

properties of diamond, graphite and graphene (in common)

they have high melting and boiling points and are insoluble in any solvent due to their strong covalent bonds.

61

why are metals good thermal conductors?

because the delocalised electrons can pass kinetic energy to each other.

62

why are metals good electrical conductors?

because the delocalised electrons can move and carry a current

63

compare and explain the electrical conductivities of diamond, graphite and graphene in terms of their structure and bonding.

diamond has all of its outer electrons in localised covalent bonds, so its a poor electrical conductor.
graphite has delocalised electrons betweens the sheets which can flow, so its a good electrical conductor.
graphene also has delocalised electrons which can flow, but it is a better conductor than graphite because it has no layers to slow the electrons down.

64

why do simple molecular structures have low melting and boiling points?

they have induced dipole to dipole forces between their molecules which are very weak and easily overcome.

65

why do noble gases have very low melting and boiling points?

because they exist as individual atoms, resulting in vey weak induced dipole to dipole forces.

66

what are the group 2 elements called?

the alkaline earth metals

67

what are the general properties of group 2 elements?

they have high melting and boiling points.
theyre light metals with low densities, and form colourless (white) compounds

68

why does reactivity increase down group 2?

as you go down group 2, the ionisation energies decrease because each successive element has its outer electrons in a higher energy level, a larger atomic radius and feels more shielding from the positive pull of the nucleus. a lower ionisation energy results in a higher reactivity as the outer electrons are lost more easily).

69

show the ionisations of group 2 elements

M → M+ + e-
M+ → M2+ + e-

70

describe the reaction of group 2 elements with water and give the general equation

- group 2 metals except beryllium react with water to produce a metal hydroxide and hydrogen.

M (S) + 2H20 (l) → M(OH)2 (aq) + H2 (g)

- the oxidation number of the metal goes from 0 to +2 (metal is oxidised)

71

describe the reaction of group 2 elements with oxygen and give the general equation

- group 2 metals burn in oxygen, to produce solid white oxides.

2M (s) + O2 (g) → 2MO (s)

- the oxidation number of the metal goes from 0 to +2 and oxygen from 0 to -2. (metal is oxidised and oxygen is reduced)

72

describe the reaction of group 2 elements with dilute acid and give the general equation

- group 2 elements except Be react with dilute acid to form a salt and hydrogen gas. eg- with HCL acid to produce a metal chloride and hydrogen.

M (S) + 2HCL (aq) → MCl2 (aq) + H2 (g)

- the oxidation number of the metal goes from 0 to +2 (metal is oxidised)

73

explain why the reaction between group 2 metal oxides and water results in an alkaline solution

group 2 oxides react with water to form metal hydroxides

MO (S) + H20 (l) → M(OH)2 (aq)

The metal hydroxide is soluble in water. When it dissolves it releases OH− ions into the water, giving an alkaline solution with a pH of between 10 and 12.

74

describe the trend in PH of the solutions formed moving down group 2

As you move down the group, the hydroxides become more soluble; this results in more OH− ions being present and a higher, more alkaline, pH.

(Be is insoluble in water. Mg(OH)2 is slightly soluble, forming a dilute solution. Ba(OH)2 is more soluble, with a more alkaline solution.

75

what does it mean by the oxides, hydroxides and carbonates of group 2 metals are basic?

they will react with acids to form a salt and water

76

what can you see when the oxides/ hydroxides/ carbonates of group 2 metals react with HCL acid?

you will see a solid oxide or hydroxide dissolve.

77

what are the uses of group 2 compounds?

- (ca(OH)2) calcium hydroxide is used in agriculture to neutralise acidic soils.
- (Mg(OH)2) magnesium hydroxide and (CaCO3) calcium carbonate are used in indigestion tablets as antacids.

Mg(OH)2 + 2HCL→ MgCl2 + 2H20

- Calcium carbonate is used as a building material- present in limestone and marble.

78

give two properties of the group 7 halogens

- they have low melting and boiling points
- they exist as diatomic molecules
- very reactive and highly electronegative- strong oxidising agents

79

explain how and why the physical state of the halogens change moving down the group

The state changes from gas, to liquid, to solid down the group. This is because the amount of electrons increases with each successive group member – causing greater London intermolecular forces. This increases the boiling point of each element.

80

explain how and why the reactivity of the halogens change moving down the group

The reactivity decreases down the group. This is because the atomic radius increases, the electron shielding increases and the ability to gain an electron into the p sub-shell and form 1– ions decreases. These factors result in each successive group member being less able to oxidise other elements.

81

what is a halogen displacement reaction?

a more reactive halogen will oxidise and displace a halide of a less reactive halogen

82

what do colour changes in halogens indicate?

that redox reactions have occurred.

83

what are the colours of the halogens- cl2, br2 and I2 in water and in cyclohexane?

halogen-- in water-- in cyclohexane

cl2-- pale green-- pale green
br2-- yellow-- orange
I2-- brown-- violet

84

state what ions cl2, br2 and I2 oxidise and give the equations

chlorine oxidises both br- and I- ions:
Cl2(aq) + 2Br- (aq) → 2Cl- (aq) + Br2(aq)
Cl2(aq) + 2I- (aq) → 2Cl- (aq) + I2(aq)

bromine can only iodise I- ions:
br2(aq) + 2I- (aq) → 2br- (aq) + I2(aq)

iodine does not oxidise cl- or br- ions

85

disproportionation

the oxidation and reduction of the same element in a redox reaction

86

describe the test for halide ions

add dilute nitric acid to the halide to remove ions that might interfere with the test.
then add silver nitrate solution (AgNO3). a precipitate is formed of the silver halide.
- results can be tested by adding ammonia solution (the solubility decreases as the ion becomes larger.)

cl- = white precipitate/ dissolves in dilute NH3
br- cream precipitate/ dissolves in conc. NH3
I- yellow precipitate/ insoluble in conc. NH3

87

explain how the disproportionation of halogens is useful in water treatments and the formation of bleach.

1) if chlorine gas is mixed with cold, dilute aqueous sodium hydroxide, sodium chlorate(I) solution- bleach is formed.
2NaOH + Cl2 = NaClO + NaCL + H2O

2) mixing chlorine with water causes disproportionation to occur, resulting in a mixture of hydrochloric acid and chloric(I) acid.
Cl2 + H2O = HCl + HClO
aqueous chloric acid ionises to make chlorate(I) ions, which kill bacteria. so adding chlorine to water makes it safer to drink or swim in.
HCLO + H2O = ClO- + H3O+

88

what are the benefits, risks and ethical considerations of using chlorine to treat water?

+ it kills disease causing microorganisms.
+ some chlorine remains in the water and prevents reinfection
+ it prevents the growth of algae, eliminating bad tastes and smells, and removes discolouration caused by organic compounds.

- chlorine gas is harmful if breathed in as it irritates the respiratory system.
- liquid chlorine on the skin or eyes can cause severe chemical burns.
- chlorine reacts with organic compounds in water to form chlorinated hydrocarbons, which are carcinogenic (cancer causing).
(however, increased cancer risk is small compared to the risks from untreated water eg- cholera epidemic.)

- we dont get a choice about having our water chlorinated- forced medication.

89

explain the alternatives of using chlorine to treat water

ozone (o3)- a strong oxidising agent so its good at killing microorganisms, but its expensive to produce and its short half-life means the treatment isnt permanent.

ultraviolet light- kills microorganisms by damaging their DNA, but its ineffective in cloudy water and isnt permanent.

90

describe the test for carbonates and include observations

add a dilute acid to the suspected carbonate. test for co2- collect any gas formed and pass it through limewater.
CO3^2- + 2H+ = CO2 + H2O

if carbonates are present: fizzing/ colourless gas produced, which turns limewater cloudy.

91

describe the test for carbonates and include observations

add dilute HCL acid and barium chloride to the suspected sulfate.
if sulfate ions are present- a whit precipitate of barium sulfate is produced.

Ba2+ + SO4^2- = BaSO4

92

describe the test for halide ions and include observations

add dilute nitric acid and then silver nitrate solution (AgNO3) to the suspected halide.
note the colour of any precipitate formed. if the colours are hard to distinguish, add aqueous ammonia and note the solubility of the precipitate.

cl- = white precipitate/ soluble in dilute NH3
br- cream precipitate/ soluble in conc. NH3 only.
I- yellow precipitate/ insoluble in conc. NH3

93

describe the test for ammonium ions (NH4+) and include observations

add sodium hydroxide solution to the suspected ammonium compound and warm gently.
test any gas evolved with red litmus paper.

ammonia gas will turn red litmus paper blue, and has a distinctive smell.
(ammonia gas is hazardous)

NH4+ +OH- = NH3 + H2O

94

what should be done to prevent mix ups when testing unknown substances?

first test for carbonates, then sulfates and then halides.

add dilute acid to the test solutions to get rid of unwanted anions.