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Flashcards in Chemistry Module 3.1 Deck (98):
1

In the early 1800s, what were the only ways to categorise elements?

In the early 1800s, the only two ways to categorise elements were by their physical and chemical properties and their relative atomic mass. This was due tot he fact at this time, they only knew how to measure the relative atomic mass.

2

In 1817, who attempted to group similar elements?

johann Dobereiner attempted to group similar elements - these groups were called Dobereiner's triads. He saw that chlorine, bromine and iodine had similar characteristics. He also realised that the other properties of bromine (e.g. atomic weight) fell halfway between those of chlorine and iodine. He found other groups of three elements (e.g. lithium, sodium and potassium), and called the triads. It was a start.

3

In 1863, who made another table of elements?

An English chemist called John Newlands had the first good attempt at making a table of elements in 1863. He noticed that if he arranged the elements in order of mass , similar elements appeared at regular intervals - every eighth element was similar. He called this the law of octaves, and he listed some known elements in rows of seven so that the similar element s lined up as columns.

4

What was the issue with John Newlands table of elements?

The law of octaves broke down in the third row, with many transition metals like Fe, Cu and Zn disrupting the pattern.

5

Who created the first accepted version of the periodic table?

Russian chemist, Dmitri Mendeleev produced a better table, which isn't far off from the one we have today. He arranged all the known elements by atomic mass, but left gaps in the table where the next element didn't seem to fit. That way he could keep elements with similar chemical properties in the same group. He also predicted the properties of undiscovered elements that would go in the gaps correctly.

6

How is the periodic table arranged?

The periodic table is arranged into periods (rows) and groups (columns).

7

What do all the elements within a period have the same number of?

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

8

Define periodicty.

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

9

What do all the elements within a group have the same number of?

All the elements within a group have the same number of electrons in their outer shell. This means they have similar chemical properties.

10

Split the periodic table into an s block, d block and p block.

{Correctly split periodic table}

11

Define what is means for an atom to be ionised.

When electrons from an atom have been removed it has been ionised.

12

What is meant by the first ionisation energy?

The first ionisation energy is the energy needed to remove 1 mole of electrons from 1 mole of gaseous atoms.

13

Why is ionisation an endothermic reaction?

Ionisation is an endothermic reaction because you have to put energy in to a molecule to ionise it.

14

What is the equation for the 1st ionisation energy of an oxygen atom?

O (g) → O⁺ (g) + e⁻

15

What are the three most important points about ionisation energies?

1. You must always use state symbols, (g), because ionisation energies are measured for gaseous atoms.2. Always refer to 1 mole of atoms, as stated in the definition, rather than to a single atom.3. The lower the ionisation energy, the easier it is to form an ion.

16

How does Nuclear charge affect ionisation energy?

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

17

How does Atomic radius affect ionisation energy?

Attraction falls off very rapidly with distance. An electron close to the nucleus will be much more strongly attracted than one further away.

18

How does shielding affect ionisation energy?

As the number of electrons between the outer electrons and the nucleus increases, the outer electrons feel less attraction towards the nuclear charge. This lessening of the pull of the nucleus by inner shells of electrons is called shielding.

19

What does a high ionisation energy mean in an atom?

A high ionisation energy means there's a strong attraction between the electron and the nucleus, so more energy is needed to overcome the attraction and remove the electron.

20

What happens to ionisation energies as you go down the group?

As you go down a group in the periodic table, ionisation energies generally fall, i.e. it gets easier to remove outer electrons.

21

Why do the ionisation energies fall as you go down the group?

feeIonisation energies go down a group because: Elements further down a group have extra electron shells compared to ones above. The extra shells mean that the atomic radius is larger, so the outer electrons are further away from the nucleus, which greatly reduces their attraction to the nucleus.The extra inner shells shield the outer electrons from the attraction of the nucleus.This provides evidence that electron shells really exist - a decrease in ionisation energy going down a group supports the Bohr model of the atom.

22

What is the general trend for the ionisation energies as you move across a period and why?

As you move across a period, the general trend is for the ionisation energies to increase - more energy is required to remove the outer electrons.This is because the no. of protons is increasing. As the positive charge of the nucleus increases, the electrons are pulled closer to the nucleus, making the atomic radius smaller.The extra electrons that the elements gain across a period are added to the outer energy level so they don't really provide any extra shielding effect (shielding works with inner shells mainly).

23

What are the two exceptions to the general trend of ionisation energies increasing across a period?

There are two exceptions to the trend:The first ionisation energy decreases between groups 2 and 3, and between groups 5 and 6.

24

Why is there a drop in ionisation energy between groups 2 and 3?

1. The outer electron in group 3 elements is in a p orbital rather than an s orbital.2. A p orbital has a slightly higher energy than an s orbital in the same shell, so the electron is, on average, to be found further from the nucleus.3. The p orbital also has additional shielding provided by the s electrons.4. These factors override the effect of the increased nuclear charge, resulting in the ionisation energy dropping slightly.

25

Why is there a drop in ionisation energy between groups 5 and 6?

1. In the group 5 elements, the electron is being removed from a singly-occupied orbital when ionised.2. In the group 6 elements, the electron is being removed from an orbital containing two electrons.3. The repulsion between two electrons in an orbital means that electrons are easier to remove from shared orbitals.

26

What is the equation for the second ionisation energy of oxygen?

O⁺ (g) → O²⁺ (g) + e⁻

27

What happens to successive ionisation energies within each shell?

Within each shell, successive ionisation energies increase. This is because electrons are being removed from an increasingly positive ion, and there's also less repulsion amongst the remaining electrons. So the electrons are held more strongly by the nucleus.

28

What causes 'big jumps' in successive ionisation energy?

The 'big jumps' in ionisation energy happen when a new shell is broken into - an electron is being removed from a shell closer to the nucleus.

29

How can one tell what group an element belongs to by looking at a successive ionisation energy graph?

One can find out what group an element belongs to from a successive ionisation energy graph by counting the electrons removed before the first big jump to find the group number.

30

How can one predict the electron configuration of an element by looking at a successive ionisation energy graph?

The graphs can also be used to predict the electronic configuration of an element. Working from right to left, count how many points there are between each big jump to find out how many electrons are in each shell, starting with the first.

31

What is a Giant covalent lattice?

A giant covalent lattice is a huge network of covalently bonded atoms. (They're sometimes called macromolecular structures too.)

32

Give an example of an atom which can form a giant covalent lattice.

Carbon atoms can form this type of structure because they can each form four strong, covalent bonds.

33

What are different forms of the same element in the same state called?

Different forms of the same element in the same state are called allotropes. Carbon has several allotropes like diamond, graphite and and graphene.

34

How many carbon atoms does carbon bond to in diamond?

In diamond, each carbon atom is covalently bonded to four other carbon atoms. The atoms rearrange themselves in a tetrahedral shape - its crystal lattice structure.

35

What are the effects of carbon having lots of strong covalent bonds?

Because it has a lot if strong covalent bonds:1. Diamond has a very high melting point - it actually sublimes at over 3800 K.2. Diamond is extremely hard - it's used in diamond tipped drills and saws.3. Vibrations can travel easily through the stiff lattice, so it's a good thermal conductor.4. It can't conduct electricity - all the outer electrons are held in localised bonds.5. it won't dissolve in any solvent.

36

What other element also forms a crystal lattice structure with similar properties to carbon?

Silicon (which is in the same periodic group as carbon) also forms a crystal lattice structure with similar properties to carbon. Each silicon atom is able to form four strong, covalent bonds.

37

Give the feature of graphite's structure that means it can be used as a dry lubricant and in pencils.

The weak forces between the layers in graphite are easily broken, so the sheets can slide over each other - graphite feels slippery and is used as a dry lubricant and in pencils.

38

Give the feature of graphite's structure that means it's able to conduct electricity.

the delocalised electrons in graphite aren't attached to any particular carbon atom and are free to move along the sheets, so an electric current can flow.

39

Give the feature of graphite's structure that means it can be used to make lightweight, strong sports equipment.

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

40

Give the feature of graphite's structure that gives it such a high melting point.

Because of the strong covalent bonds in the hexagon sheets, graphite has a very melting point.

41

Give the feature of graphite's structure that makes it insoluble in any solvent.

Like diamond, graphite is insoluble in any solvent. The covalent bonds in the sheets are too strong to break.

42

What is Graphene?

Graphene is a sheet of carbon atoms joined together in hexagons. The sheet is just one atom thick, making it a two-dimensional compound.

43

What makes graphene the best known electrical conductor?

Like in graphite, the delocalised electrons in graphene are free to move along the sheet. Without layers, they can move quickly above and below the sheet, making graphene the best known electrical conductor.

44

What makes graphene extremely strong?

The delocalised electrons also strengthen the covalent bonds between the carbon atoms. This makes graphene extremely strong.

45

Describe a single layer of graphene.

A single layer of graphene is transparent and incredibly light.

46

What properties of graphene make it a useful compound in industry?

Due to its high strength, low mass, and good electrical conductivity, graphene has potential applications in high-speed electronics and aircraft technology. its flexibility and transparency also make it a potentially useful material for touchscreens on smartphones and other electronic devices.

47

What structure do metal elements exist as?

Metal elements exist as giant metallic lattice structures.

48

Describe the giant metallic lattice structure of metal elements.

The electrons in the outermost shell of a metal atom are delocalised - the electrons are free to move about the metal. This leaves a partially charged metal cation.The metal cations are electrostatically attracted to the delocalised negative electrons. They form a lattice of closely packed cations in a sea of delocalised electrons. - This is metallic bonding.

49

How does metallic bonding explain the properties of metals?

1. The number of delocalised electrons per atom affects the melting point. The more there are, the stronger the bonding will be and the higher the melting point. The size of the metal ion and the lattice structure also affect the melting point. a smaller ionic radius will hold the delocalised electrons closer to the nuclei.

2. As there are no bonds holding specific ions together, the metal ions can slide past each other when the structure is pulled, so metals are malleable (can be hammered into sheets) and ductile (can be drawn into a wire).

3. The delocalised electrons can pass kinetic energy to each other, making metals good thermal conductors.

4. Metals are good electrical conductors because the delocalised electrons can move and carry current.

5. Metals are insoluble, except in liquid metals, because of the strength of the metallic bonds.

50

How does the structure of simple molecules explain their properties?

1. Simple molecular structures contain only a few atoms.2. The covalent bonds between the atoms in the molecule are very strong, but the melting and boiling points of simple molecular substances depend on the induced dipole-dipole forces between their molecules. These intermolecular forces are weak and easily overcome, so these elements have low melting and boiling points.3. More atoms in a molecule mean stronger induced dipole-dipole forces. Meaning a higher boiling point.4. The noble gases have very low melting and boiling points because they exist as individual atoms (they're monatomic), resulting in very weak induced dipole-dipole forces.

51

How are the melting points in metals affected across a period?

For the metals, melting and boiling points increase across the period because the metallic bonds get stronger as the ionic radius decreases and the number of delocalised electrons increases.

52

How are the melting points in elements with giant covalent lattice structures (C and Si) affected across a period?

The elements with giant covalent lattices (C and Si) have strong covalent bonds linking all their atoms together. A lot of energy is needed to break these bonds.

53

How are the melting points in elements that form simple molecular structures affected across a period?

The elements that form simple molecular structures have only weak intermolecular forces to overcome between their molecules, so they have low melting and boiling points.

54

How are noble gases' melting and boiling points affected across a period?

The noble gases have the lowest melting and boiling points in their periods because they are held together by the weakest forces.

55

How many electrons do group 2 elements have in their outer shell?

Group 2 elements all have two electrons in their outer shell. They lose tow outer electrons to form 2+ ions.

56

What happens to the ionisation energies of group 2 elements as you go down the group?

As you go down the group, the ionisation energies decrease. This is due to the increasing atomic radius and shielding effect.

57

why does reactivity increase going down group 2?

When group 2 elements react they lose electrons, forming cations. The easier it is to lose electrons, the more reactive the element, so reactivity increases down the group.

58

What happens to the oxidation state of group 2 elements when they react?

When group 2 elements react, they are oxidised from a state of 0 to +2, forming M²⁺ ions.M → M²⁺ + 2e⁻0 +2

59

What do group 2 metals react with water to produce?

The group 2 metals react with water to give a metal hydroxide and hydrogen.

60

What do group 2 metals burn in oxygen to form?

When group 2 metals burn in oxygen, you get solid white metal oxides.

61

What do group 2 metals react with dilute acid to produce?

They react with dilute acid to produce a salt and hydrogen.

62

What are the metal oxides and metal hydroxides of group 2 metals examples of?

The metal oxides and metal hydroxides of group 2 metals are bases. Most of them are soluble in water so are alkalis.

63

How do group 2 metal oxides form very alkaline solutions?

The oxides of the group 2 metals react readily with water to form metal hydroxides, which dissolve. The hydroxide ions, OH⁻, make these solutions strongly alkaline.CaO (s) + H₂O (l) + Ca²⁺ (aq) + 2OH⁻ (aq)

64

What makes Magnesium oxide an exception to group 2 metal oxide reactions with water?

Magnesium oxide is an exception as it only reacts slowly and the hydroxide isn't very soluble.

65

Why do the group 2 metal oxides form more strongly alkaline solutions when reacted with water as you go down the group?

The group 2 metal oxides form more strongly alkaline solutions as you go down the group, because the hydroxides get more soluble.

66

What are common examples of group 2 compounds used for neutralising acids?

1. Calcium hydroxide (slaked lime (Ca(OH)₂) is used in agriculture to neutralise acidic soils.2. Magnesium hydroxide (Mg(OH)₂) and calcium carbonate (CaCO₃) are used in some indigestion tablets as antacids.

67

Give the ionic equation for neutralisation.

H⁺ (aq) + OH⁻ (aq) → H₂O (l)

68

What are the main properties of fluorine?

Formula: F₂ Colour: Pale yellow Physical state at (20°C): Gas Electronic structure: 1s² 2s² 2p⁵

69

What are the main properties of chlorine?

Formula: Cl₂ Colour: Green Physical state at (20°C): Gas Electronic structure: 1s² 2s² 2p⁶ 3s² 3p⁵

70

What are the main properties of bromine?

Formula: Br₂ Colour: Red-brown Physical state at (20°C): Liquid Electron structure: 1s² 2s² 2p⁶ 3s² 3p⁶ 4s² 3d¹⁰ 4p⁵

71

What are the main properties of Iodine?

Formula: I₂ Colour: Grey Physical state at (20°C): Solid Electron structure: 1s² - 5p⁵

72

Do halogens exist as diatomic or monatomic molecules?

halogens exist as diatomic molecules (two atoms joined together by a single covalent bond).

73

Why does the boiling/melting point of the halogens increase down the group?

This is due to the increasing strength of the London (induced dipole-dipole) forces as the size and relative mass of the atoms increases. This trend is shown in the changes of physical state from chlorine (gas) to Iodine (solid).

74

What does it mean for a substance to be described as volatile?

A substance is said to be volatile if it has a low boiling point.

75

What happens to the volatility of the halogens as you go down group 7?

Volatility decreases down the group.

76

Why are halogens described as oxidising agents?

Halogens are described as oxidising agents because they react by gaining an electron in their outer shells forming 1- ions. This means they're reduced. As they're reduced, they oxidise another substance (it's a redox reaction) - so they're oxidising agents.

77

Why are halogen atoms less reactive further down the group?

As you go down the group, the atomic radii increase so the outer electrons are further from the nucleus. The outer electrons are also shielded more from the attraction of the positive nucleus, because there are more inner electrons. This makes it harder for the larger atoms to attract the electron needed to form an ion (despite the increased charge on the nucleus), so larger atoms are less reactive.

78

What's another way of saying that halogens get less reactive down the group?

Another way of saying that the halogens get less reactive down the group is to say they become less oxidising.

79

How can one observe the halogens' relative oxidising strengths?

The halogens' relative oxidising strengths can be seen in their displacement reactions with halide ions.

80

How can one observe a displacement reaction?

When a displacement reaction happens, there are a colour changes. You can make these changes easier to see by shaking the reaction mixture with an organic solvent like hexane. the halogen that's present will dissolve readily in the organic solvent, which settles out a distinct layer above the aqueous solution.

81

Table for colour change in displacement reactions of chlorine with other halogens.

  KCl (aq) - colourless KBr (aq) - colourless KI (aq) - colourless   In aqueous solution In organic solution In aqueous solution In organic solution In aqueous solution In organic solution Chlorine water Cl₂ (aq) - colourless No reaction  No reaction Yellow (Br₂) Orange (Br₂) Orange/brown (I₂) Purple (I₂) Bromine water Br₂ (aq) - yellow No reaction  No reaction No reaction No reaction Orange/brown (I₂) Purple (I₂) Iodine solution I₂ (aq) - orange/brown No reaction  No reaction No reaction No reaction No reaction No reaction   KCl (aq) - colourless KBr (aq) - colourless KI (aq) - colourless   In aqueous solution In organic solution In aqueous solution In organic solution In aqueous solution In organic solution Chlorine water Cl₂ (aq) - colourless No reaction  No reaction Yellow (Br₂) Orange (Br₂) Orange/brown (I₂) Purple (I₂) Bromine water Br₂ (aq) - yellow No reaction  No reaction No reaction No reaction Orange/brown (I₂) Purple (I₂) Iodine solution I₂ (aq) - orange/brown No reaction  No reaction No reaction No reaction No reaction No reaction

82

How does one test for halides?

  1. Add the dilute nitric acid to remove ions that may interfere with the test. 
  2. Add the silver nitrate solution. A precipitate is formed (of the silver halide).
  3. The colour of the precipitate identifies the halide.
  4.  To be extra sure, you can test the results by adding ammonia solution. (Each silver halide has a different solubility in ammonia - the larger the ion is, the more difficult it is to dissolve.

  • Chloride Cl⁻: white precipitate, dissolves in dilute ammonia
  • Bromide Br⁻: cream precipitate, dissolves in concentrated ammonia
  • Iodide I⁻: yellow precipitate, insoluble in concentrated ammonia

   

83

Describe what a disproportionation reaction is with halides and alkalis.

The halogens will react with cold dilute alkali solutions. In these reactions, the halogen is simultaneously oxidised and reduced - this is disproportionation.

84

What does cold, dilute aqueous sodium hydroxide react with chlorine gas to produce?

If you mix chlorine gas with cold, dilute aqueous sodiu hydroxide you get sodium chlorate (I) solution, NaClO (aq) which is the common household bleach.

85

List the uses of sodium chlorate (I) solution.

Used in:

  • Water treatment
  • Bleaching paper and textiles

86

What happens when you mix chlorine with water?

When you mix chlorine with water it undergoes disproportionation. You end up with a mixture of hydrochloric acid and chloric (I) acid (also called hypochlorous acid).

87

What does aqeous chloric acid ionise in water to form?

Aqeous chloric acid ionises to make chlorate (I) ions. 

88

What is useful about chlorate ions?

Chlorate ions kill bacteria so adding chlorine (or a compound containing chlorate ions) to water can make it safe to swim in or drink.

89

List what makes chlorine so imoortant in water treatment.

  • It kills some disease causing microorganisms.
  • Some chlorine remains in the water and prevents reinfection further down the supply.
  • it prevents the growth of alage, eliminating bad tastes and smells, and removes discolouration caused by organic compounds.

90

What are the risks of using chlorine to treat water?

Chlorine gas is very harmful if it's breathed in - it irritates the respiratory system. Liquid chlorine on the skin or eyes causes severe chemical burns. accidents invloving chlorine could be really serious, or fatal. 

Water contains a variety of organic comounds e.g. from the decomposition of plants. Chlorine reacts with these compounds to form chloronated hydrocarbons and many of them are carcinogenic. However, this increased risk is small compared to the risk from untreated water - a chloera epiudemic could kill thousands of people.

91

Give an ethical consideration about treating water.

We don't actually get a choice about having our water chlorinated - some people object to this as forced 'mass medication'.

92

GIve two alternatives to chlorine for water treatment.

Ozone (O3) - a strong oxidising agent, which makes it great at killing microorganisms. However, it's expensive to produce and its short half life in water means that the water treatment isn't permanent.

Ultraviolet light - it kills microorganisms by dmaaging their DNA, but it's ineffective in cloudy water and, like O3, it won't stop the water being contaminated further down the line. 

93

Decsribe the test for carbonates.

To test for carbonates (CO3)2-, add a dilute acid (like HCl) to the unknown sample. If carbbonates are present then CO2 will be released

CaCO3(s) + HCl(aq)  → CO2(g) + H2O(l) + CaCl2(aq)

You can test for carbon dioxide using limewater. carbon dioxide turns limewater cloudy - just bubble the gas through a test tube of limewater and watch the solution turn cloudy.

94

Describe the test for sulfates.

Most sulfates are soluble in water, but barium sulfate is insoluble. So test for a sulfate ion (SO42-), add dilute HCl, followed by barium chloride solution, BaCl2. 

If you get a white precipitrate it'll be barium sulfate, whic tells you the unknown substance is a sulfate.

95

Describe the test for halides with silver nitrate.

To test for halide ions just add nitric acid, then silver nitrate solution. if chlorine, bromide or iodide is present, a precipitate will form. the colour of the precipitate depends on the halide present.

  • Silver chloride (AgCl) is a white precipitate
  • Silver bromide (AgBr) is a cream precipitate
  • Silver iodide (AgI) is a yellow precipitate.

You can test for the solubility of these precipitates in ammonia to help you tell them apart.

AgCl dissolves in dilute NH3 and concentrated NH3.

AgBr doesn't dissolve in dilute NH3 but it doesn in concentrated NH3.

AgI doesn't dissolve in dilute or concentrated NH3. 

96

Describe the test for ammonium compounds.

  1. Ammonia gas (NH3) is alkaline, so you can check for it using a damp piece of red litmus paper. If there's ammonia present, the paper will turn blue.  (The litmus paper has to be damp so the ammonia gas can dissolve and make the colour change).
  2. You can use this to test to whether a substance contains ammonium ions (NH4+). Add some sodium hydroxide to the unknown substance in a test tube and warm the mixture. If there's ammonia given off this means there are ammonium ions in your unknown substance.

97

What are the possible false positives in the test for ions?

1. As well as barium sulfate, barium carbonate and barium sulfite are also insoluble. So if you're testing for sulfate ions, you want to make sure there are no carbonate ions or sulfite ions around first.

2. Likewise, if you're testing for a halide ion, you want to rule out the presence of sulfate ions first. this is because sulfate ions will also produce a precipitate with silver nitrate.

3. A gooday of getting around this is to first add dilute acid to your test solutions. the acid will get rid of any anions you don't want.

98

To avoid inaccuracies in results, what order should you do the tests for ions?

1. Test for carbonates;

    No CO2 produced?

2. Test for sulfates;

    No precipitate produced?

3. Test for halides.