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chemistry > S3.1 - the periodic table: classification of elements > Flashcards

S3.1 - the periodic table: classification of elements Flashcards

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1
Q

What are periods?

A

Horizontal columns on the periodic table.

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2
Q

What are groups?

A

Vertical rows on the periodic table.

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3
Q

Where are the blocks located?

A

S - left hand side
P - right hand side
D - middle
F - bottom

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4
Q

What are metalloids?

A

Metalloids - elements with physical properties of metals, but chemical properties of non-metals.

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5
Q

What are examples of metalloids?

A

Silicon, Germanium, Arsenic, Antimony, Tellerium and Polonium

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6
Q

What does the period number show?

A

Number of energy levels/orbitals

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7
Q

What does group number show?

A

Number of valence electrons

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8
Q

What are valence electrons and what do they show?

A

Group number - number of electrons on the outer shell of an atom.

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9
Q

What elements are located in the d-block?

A

Transition metals

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10
Q

What are group 1 elements?

A

Alkali metals

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11
Q

What are group 7 elements?

A

Halogens

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12
Q

What are group 0 elements?

A

Nobel gases

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13
Q

What is periodicity?

A

Periodicity - physical and chemical properties repeat periodically in the periodic table.

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14
Q

What is effective nuclear charge?

A

ENC - net positive charge experienced in a multi-electron atom
ENC = proton number - electron number in previous noble gas
- Also the group number of an atom

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15
Q

What is shielding?

A

Reduction of attractive forces between the nucleus and outer electron by inner electron shells.

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16
Q

What happens to ENC across the period?

A

Across the period, ENC increases as there is an increasing number of protons in the nucleus with the same amount of shells and shielding, so there is a stronger attraction between the outer electron and nucleus.

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17
Q

What happens to ENC down the group?

A

Down the group, ENC decreases as there is more shells and shielding, so a weaker attraction between the outer electron and nucleus.

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18
Q

What is atomic radius?

A

Atomic radius - half the distance between neighboring nuclei in a covalent bond.

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19
Q

What is ionic radius?

A

Ionic radius - half the distance from the nucleus to the outer electron.

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20
Q

What happens to atomic radius down the group and across the period?

A

Down the group - increases
Across the period - decreases

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21
Q

What happens to ionic radius down the group and across the period?

A

Down the group - increases
Across the period - decreases

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22
Q

Why does ionic radii decrease from group 1 to 4?

A

Ionic radii decreases for positive ions due to the increase in ENC with atomic number across the period, increasing the attraction between the nucleus and outer electron and pulling the outer energy level closer in.

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23
Q

Why does ionic radii decrease from group 4 to 7?

A

Ionic radii decreases from group 4 to 7 for negative ions due to the increase in ENC.

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24
Q

Why is there a big difference in the ionic radii of Si4+ and Si4-?

A

Negative ions are larger than positive ions as they have more energy levels.
- Also explains the discontinuity for transition metals.

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25
Q

Why does the ionic radii decrease down a group?

A

Number of electron energy levels increases.

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26
Q

What are the properties of positive ions?

A
  • Smaller than parent atoms.
  • Formation involves loss of outer energy level.
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27
Q

What are the properties of negative ions?

A
  • Larger than parent atoms.
  • Formation involves addition of electrons to outer energy level.
  • Increased electron-electron repulsion between electrons in the outer energy level causes electrons to move further apart, increasing radius of the outer energy level.
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28
Q

What is first ionization energy?

A

Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms in their ground state.

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29
Q

What is second ionization energy?

A

Energy required to remove 1 mole of electrons from 1 mole of gaseous ions.

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30
Q

What happens to ionization energy down the group?

A

Down the group, ionization energy decreases as the outer electron is further away from the nucleus, meaning there is more shielding and a weaker attraction, requiring less energy to remove an electron.

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31
Q

What happens to ionization energy across the period?

A

Across the period, ionization energy increases as ENC increases, as there is more protons with the same shielding, meaning a strong attraction between nucleus and outer shell electron, which requires more energy to remove.

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32
Q

What is electron affinity?

A

Energy change that occurs when 1 mole of electrons is added to 1 mole of gaseous atoms.

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33
Q

What is second electron affinity?

A

Energy change that occurs when 1 mole of electrons is added to 1 mole of gaseous ions.

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34
Q

What type of reaction is first electron affinity?

A

Exothermic and gives out energy.
- Added electron is attracted to the + charged nucleus.

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35
Q

What type of reaction is second and third electron affinity?

A

Endothermic, requiring energy.
- Added electrons are repelled by the - charged ion.

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36
Q

What happens to electron affinity down the group?

A

Down the group, electron affinity decreases as the energy levels increases, meaning there is a weaker electron between the nucleus and outer electron.

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37
Q

What happens to electron affinity across the period?

A

Across the period, electron affinity increases as ENC increases, meaning there is a greater attraction between the nucleus and outer electron.

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38
Q

What is electronegativity?

A

Ability of an atom to attract a pair of electrons in a covalent bond.
- Measures of the attraction between the nucleus and outer electron.

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39
Q

What happens to electronegativity across the period?

A

Across the period, electronegativity increases due to increased ENC, meaning there is an increased attraction between the nucleus and outer shell electrons.

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40
Q

What happens to electronegativity down the group?

A

Down the group, electronegativity decreases as there are more shells and shielding between bonding electrons and nucleus, so there is a lower attraction.

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41
Q

Why is N and O an exception?

A

Oxygen has a lower IE as a pair of electrons is easier to remove than a single electron due to the repulsion.

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42
Q

Why is B and Be an exception?

A

B has a lower IE as electrons are more easily lost from the 2p1 orbital.

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43
Q

What is the Pauling scale?

A

Scale which assigns electronegativity values to each element.

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44
Q

What determines chemical properties of an element?

A

Number of valence electrons in the outer shell of an atom.

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45
Q

What are physical properties of group 1 metals?

A
  • Good conductors of heat and electricity
  • Low density
  • Shiny grey when freshly cut
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46
Q

What are chemical properties of group 1?

A
  • Very reactive metals (increasing in metallic character across the period)
  • Form ionic compounds with non-metals
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47
Q

What are the reactions of group 1 with water?

A

Metal + Water –> Salt + Hydrogen
- Releases a gas
- Dissolves
- Melts into a sphere
- Produces a flame

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48
Q

What happens when lithium reacts with water?

A

Lithium floats and reacts slowly. It releases H, but keeps its shape.
2Li (s) + 2H2O (l) –> 2LiOH (aq) + H2 (g)

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49
Q

What happens when sodium reacts with water?

A

Sodium reacts with a vigorous release of hydrogen. Heat produced is sufficient to melt the unreacted metal, forming a small ball which moves around on the water surface.
2Na (s) + 2H2O (l) –> 2NaOH (aq) + H2 (g)

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50
Q

What happens when potassium reacts with water?

A

Potassium reacts even more vigorously to produce sufficient heat to ignite the hydrogen produced. It produces a purple flame and moves excitedly on the water surface.
2K (s) + 2H2O (l) –> 2KOH (aq) + H2 (g)

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51
Q

What bonding is present with group 1?

A

Metallic bonding - attraction between a cation lattice and a sea of delocalised electrons.

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52
Q

What are the half-reactions of group 1?

A

Li (s) –> Li+ (s) + e-

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53
Q

What happens to reactivity down group 1?

A

Reactivity increases down 1 as there is more shielding, lower ionization energy and ENC, a larger radii and a weaker attraction between nucleus and electrons, meaning it requires less energy to overcome.

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54
Q

What happens to boiling point down group 1?

A

Boiling point decreases down group 1 as atomic radius increases, so there is a weaker attraction between outer electrons and nucleus, which requires less energy to overcome.

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55
Q

Why are group 1 conductors?

A

Group 1 are conductors as they contain delocalised electrons which can move around the structure and carry charge.

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56
Q

Why is rubidium more reactive than Na?

A

Rb is more reactive than Na, as it is further down group 1, meaning a larger atomic radius, so there is a weaker attraction between outer electron and nucleus, meaning the electron is lost more easily, using less energy to overcome.

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57
Q

What are the properties of noble gases?

A
  • Colourless, monatomic gases
  • Stable electronic configuration
  • No electron affinity
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58
Q

Why are noble gases so unreactive?

A

Don’t form + ions - high IE and ENC, so don’t lose electrons easily.
Don’t form - ions - electron would need to be added to a new shell with lots of shielding, which would get lost.

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59
Q

What are physical properties of group 7?

A
  • They are a range of colours and states.
    Fluorine - pale yellow gas
    Chlorine - yellow-green gas
    Bromine - red-brown liquid
    Iodine - purple gas, brown liquid, grey solid crystal
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60
Q

What are chemical properties of group 7?

A
  • Very reactive non-metals
  • Form ionic compounds with group 1
  • Form covalent compounds with non-metals.
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61
Q

What bonding is present in group 7?

A

Covalent bonding

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62
Q

What are displacement reactions?

A

When a more reactive element displaces a less reactive element from its compound.
2Na (s) + Cl2 (g) –> 2NaCl (s) = salt

63
Q

What are the half-equations for group 7?

A

Cl2 (g) + 2e- –> Cl (g)

64
Q

What happens to reactivity down group 7?

A

Down group 7, reactivity decreases as there is a high ENC, very exothermic electron affinity which decreases due to increased shielding and atomic radii, requiring more energy to overcome.

65
Q

What happens to melting points down group 7?

A

Melting points increase down group 7 as there are increasing electrons and LDF, which require more energy to overcome.

66
Q

What is produced when a group 7 reacts with silver nitrate?

A

Precipitation reaction - silver ions react with halide ions to form different coloured precipitates.

67
Q

What colors are produced when silver nitrate reacts with Cl, I and Br?

A

Cl - white
Br - cream
I - yellow

68
Q

What is the trend of period 3 from left to right?

A

Continuum from basic metal oxides (Na2O and MgO) to amphoteric oxides (Al2O3) to acidic non-metal oxides (SiO2, P4O10, SO3, Cl2O7).
- Goes from ionic to covalent bonding.

69
Q

What is the trend of period 3 from left to right in terms of oxidation number?

A

Oxidation number increases, and transition metals have multiple oxidation states.

70
Q

What is the trend of period 3 from left to right in terms of electrical conductivity?

A

Goes from high conductivity to very low and no conductivity.

71
Q

What is the trend of period 3 from left to right in terms of structure?

A

Goes from giant ionic to giant covalent to molecular covalent sructure.

72
Q

What is the trend of period 3 from left to right in terms of acid-base character?

A

Goes from basic to amphoteric to acidic period 3 oxides.

73
Q

What are basic oxides?

A

Metal oxides that are ionic.

74
Q

What are acidic oxides?

A

Non-metal oxides that are covalent.

75
Q

What are amphoteric oxides?

A

Show both acidic and basic properties, such as aluminium oxide.

76
Q

What are alkalis?

A

Bases that are soluble in water to form hydroxide ions in aqueous solutions.

77
Q

How do basic oxides react with water?

A

Basic oxides dissolve in water to form alkaline solutions due to the presence of OH ions.
MgO (s) + H2O (l) –> Mg(OH)2 (aq)

78
Q

How do basic oxides react with acid?

A

Basic oxides react with an acid to form salt and water.
MgO (s) + 2HCl (aq) –> MgCl2 (aq) + H2O (l)

79
Q

How do phosphorus oxides react with water?

A

P4O10 (s) + 6H2O (l) –> 4H3PO4 (aq) - Phosphorus (V)
P4O6 (s) + 6H2O (l) –> 4H3PO3 (aq) - Phosphorus (III)

80
Q

How do acidic oxides react with water?

A

Non-metallic oxides react with water readily to produce acidic solutions.

81
Q

How does carbon dioxide react with water?

A

CO2 (g) + H2O (l) –> H2CO3 (aq)

82
Q

How does silicon dioxide react with alkalis?

A

SiO2 (s) + 2OH- (aq) –> SiO3 2- (aq) + H2O (l) - forms silicates with alkalis

83
Q

How does sulfur trioxide react with water?

A

SO3 (l) + H2O (l) –> H2SO4 (aq) - Sulphuric (VI) acid

84
Q

How does sulfur dioxide react with water?

A

SO2 (l) + H2O (l) –> H2SO3 (aq) - Sulphuric (IV) acid

85
Q

What happens when an amphoteric oxide is added to water?

A

Doesn’t affect the pH as it is insoluble when added to water, instead showing both acidic and alkaline properties.

86
Q

How does an amphoteric oxide react with acids?

A

Behaves as a base when it reacts with an acid.
Al2O3 (aq) + 3H2SO4 (aq) –> Al2(SO4)3 (aq) + 3H2O (l)

87
Q

How does an amphoteric oxide react with alkalis?

A

Behaves as an acid when it reacts with an alkali.
Al2O3 (aq) + 2OH- (aq) –> 2Al(OH)4- (aq)

88
Q

How does Cl2O7 and Cl2O react with water?

A

Cl2O7 (aq) + H2O (l) –> 2HClO4 (aq) - Chloric (VII) oxide
Cl2O (aq) + H2O (l) –> 2HClO (aq) - chloric (I) oxide

89
Q

What is acid rain?

A

Substance with a pH less than 5.6

90
Q

What is rainwater?

A

Rainwater is naturally acidic due to the dissolved CO2.

91
Q

How does sulfur make acid rain?

A

1) S (s) + O2 (g) –> SO2 (aq)
2) SO2 (aq) + H2O (l) –> H2SO3 (aq)
3) 2SO2 (aq) + O2(g) –> 2SO3 (aq)
4) SO3 (aq) + H2O (l) –> H2SO4 (aq)

92
Q

How is sulfur dioxide produced?

A

Sulfur dioxide is produced from the burning of fossil fuels, such as coal and heavy oil, in power plants to generate electricity.
It is also released in industrial processes such as smelting, where metals are extracted from their ores.

92
Q

How are nitrogen oxides produced?

A

NO is produced from internal combustion engines, where the burning of the fuel releases heat energy that causes N and O in the air to combine.

93
Q

How does ocean acidification occur?

A

Ocean acidification occurs as carbon dioxide dissolves in the oceans.
CO2 (g) + H2O (l) –> H2CO3 (aq)
The carbonic acid then reacts to form hydrogencarbonates or carbonates:
H2CO3 (aq) <–> H+ (aq) + HCO3- (aq)
HCO3- (aq) <–> H+ (aq) + CO3 2- (aq)

94
Q

What are the consequences of ocean acidification?

A
  • Higher ocean acidity near the surface, where CO2 is absorbed.
  • Higher acidity inhibits shell growth in marine animals.
  • Causes reproductive disorders in some fish.
  • Coral bleaching
95
Q

What is an oxidation state?

A

Number assigned to an atom to show the number of electrons that are transferred when forming a bond. It is the charge the atom would have if the compound were composed of ions.

96
Q

What is oxidation?

A
  • Gain of oxygen
  • Loss of hydrogen
  • Loss of electrons
  • Increase in oxidation state
97
Q

What is an increase in oxidation state?

A
  • Atom has lost an electron.
  • Atom has been oxidized.
98
Q

What is a decrease in oxidation state?

A
  • Atom has gained an electron.
  • Atom has been reduced.
99
Q

What are the 10 rules for oxidation states?

A
  1. atoms in the form of elements have an oxidation state of 0
  2. simple ions have the same oxidation state as the charge of the ion
  3. in a compound, the total oxidation state must add up to 0
  4. oxidation states in a poly atomic ion adds up to the charge of the ion
  5. oxidation state is usually the same as the group number or valence electrons
  6. F has an oxidation state of -1 in all compounds as it is the most electronegative
  7. O has an oxidation state of -3 except in H2O2 where it is -1 and it is positive with Fl
  8. Cl has an oxidation state of -1, but is positive when bonded to O or Fl as it is less electronegative
  9. H has the oxidation state of +1 except when bonded to group 1 and 2 elements to form ionic hydrides
  10. oxidation states of transition metals in a complex ion can be worked out from the ligand charge
100
Q

What discontinuities occur across period 4?

A

Discontinues occur in the trend of increasing first ionization energy across a period.

101
Q

What do discontinuities provide evidence for?

A

Discontinuities across period 4 provides evidence for the existence of energy sub-levels.

102
Q

Where are transition metals located?

A

In the middle of the periodic table

103
Q

What happens to atomic radius across period 4 and why?

A

Across period 4, there is a relatively small decrease in atomic radius. There is a small increase in ENC experienced by 4s electrons. This is because when every extra proton is added, another inner 3d electron is also added.

104
Q

Why can transition metals form alloys?

A

This is because atoms of 1 d-block metal can be replaced by atoms of another without too much disruption to the solid structure.

105
Q

Why is there a small range in first ionization energies across the first d-block elements?

A

This is due to the small increase in ENC.

106
Q

Why is zinc not a transition metal?

A

Zinc is a d-block element, but is not a transition metal as it doesn’t form colored compounds and exists only in the +2 oxidation state, which results in a completely filled d shell, which isn’t a property of transition metals.
Zn = 1s2, 2s2, sp6, 3s2, 3p6, 4s2, 3d10
Zn2+ = 1s2, 2s2, sp6, 3s2, 3p6, 3d10

107
Q

Why is Scandium not a transition metal?

A

This is because the most stable ion (Sc3+) doesn’t have any electrons in the d-orbitals which are required for transition metals, but the Sc2+ ion, although not common has a single d-electron.
Sc =1s2, 2s2, sp6, 3s2, 3p6, 4s2, 3d1
Sc3+ = 1s2, 2s2, sp6, 3s2, 3p6
Sc2+ = 1s2, 2s2, sp6, 3s2, 3p6, 3d1

108
Q

What are physical properties of transition metals?

A
  • high electrical and thermal conductivity
  • high melting point
  • malleable = they are easily beaten into shape
  • high tensile strength = they can hold large loads without breaking
  • ductile = they can be easily drawn into wires
  • iron, cobalt and nickel are ferromagnetic
109
Q

What are the magnetic properties of transition metals?

A
  • iron, cobalt and nickel are ferromagnetic and show strong magnetic properties.
  • this is due to the presence of unpaired electrons in the d-orbitals which give transition metals magnetic properties.
  • spinning electrons behave like a tiny magnet, when paired they have opposite spins so they cancel out.
  • when there is an unpaired electron, there is a specific polarity and direction (such as in NMR), leading to magnetic properties.
110
Q

What are chemical properties of transition metals?

A
  • forms compounds with more than 1 oxidation state
  • forms a variety of complex ions
  • forms colored compounds
  • Acts as a catalyst when either an element or compound
111
Q

How can transition metals be used as catalysts?

A

1) Fe in the Haber process - important material for production of other chemicals such as fertilizers, plastics, drugs and explosives.
2) Ni in the conversion of alkenes into alkanes - e.g) unsaturated vegetable oil into margerine.
3) Pd, Rh and Pt in catalytic converters - remove harmful primary pollutants from car exhaust engines and turn into less harmful secondary pollutants.

112
Q

What makes something a transition metal?

A

In order to be considered a transition metal, the metal IONS involved must contain a partially filled d-subshell.
- Not all elements in the d-block are transition metals.

113
Q

How can the formation of variable oxidation states be explained?

A

The formation of variable oxidation states of transition metals can be explained by the fact that their successive ionization energies are close in value.

114
Q

What are the key points on variable oxidation states of transition elements?

A
  • All metals have +2 and +3 oxidation states with M3+ being more stable for elements from Sc to Cr and M2+ being more stable for elements Cr to Cu.
  • Maximum oxidation state is Mn (Mn7+).
  • Oxidation numbers increases by +1 across the period till Mn and then decreases by -1 after Mn to Cu.
115
Q

Why does Ca form Ca2+ ions, but not Ca3+ ions?

A

Ca only forms Ca2+ ions as you would only need to remove the outer 4s shell. However to produce Ca3+, a big energy jump to remove the electrons from the 3p sub-shell is required, yet there is no energy available.

116
Q

Why can Titanium and other transition metals form multiple transition states?

A

Titanium and other transition metals can form multiple transition states as the 3d and 4s energy levels from which electrons are removed are very close together in energy, so is easier to remove as it requires less energy.

117
Q

What are characteristics of compounds with high oxidation states?

A

Compounds with high oxidation states tend to be oxidizing agents (which get reduced), such as K2Cr2O7 or MnO4-.

118
Q

What does high oxidation numbers mean in terms of bonding?

A

High oxidation numbers means more covalent characters (+3 and higher oxidation states).

119
Q

What is the relationship between charge density and covalent character?

A

The high charge density of the transition metal pulls the weakly held outer electron of a negative ion. The polarization (electron is pulled closer to the cation from the ion due to stronger attraction, leading to sharing of electrons) of the negative ions increases the covalent character of the compound.

120
Q

What bonding is present in K2Cr2O7 (potassium dichromate)?

A

Despite there being bonding between a metal and non-metal, the Cr2O7 acts as a single entity and has a charge of 2- as it is the only way for the elements to stick together. As a result of the high charge density, there is higher covalent character and covalent bonding occurs.

121
Q

Why is copper considered a transition metal?

A

Copper is considered a transition metal as it forms coloured compounds, has more than 1 oxidation state, has electrons in the d-shell and acts as a catalyst.

122
Q

Why is Scandium not considered a transition metal?

A

Scandium is not considered a transition metal as it doesn’t have any d-shell electrons, doesn’t form coloured compounds, only has 1 oxidation state and doesn’t act as a catalyst.

123
Q

Why are transition elements complexes coloured?

A

Transition element complexes are coloured due to the absorption of light when an electron is promoted between orbitals in the split 3d-sublevels.
- They absorb some of the colours in the visible light spectrum and the colour the solution appears is due to the ions in the solution absorbing light in the complementary region to the colour or solution.

124
Q

What do transition metals form and why??

A

Transition metals form complex ions in solution as the ions have a high charge density.

125
Q

When does a complex ion form?

A

A complex ion forms when a central ion is surrounded by molecules with a lone pair of electrons, called ligands.

126
Q

What is the relationship between the colour absorbed and the colour observed?

A

The colour absorbed is complementary to the colour observed.

127
Q

What is a coordination bond?

A

Uses a lone pair of electrons to form a covalent bond

128
Q

What is a co-ordination number?

A

Co-ordination number is the number of ligands surrounding a molecule.

129
Q

What are ligands?

A

Ligands - species that uses a lone pair of electrons to form a coordination bond with a metal ion.
- Can be neutral (H2O) or charged (Cl-)
- All ligands have at least 1 atom with a lone pair of electrons

130
Q

What are some properties of complex ions?

A
  • Overall ion is charged
  • Can be a combination of different ligands
131
Q

What occurs when a transition metal ion dissolves in water?

A

When a transition metal ion dissolves in water, the d orbital splits due to the electric field produced by the ligand’s lone pair of electrons and when light passes through the solution, a photon of visible light is absorbed by one of the d electrons, which is promoted to a higher energy level.

132
Q

Why does Sc3+ appear colourless?

A

Sc3+ doesn’t have any d electrons, so splitting cannot occur.

133
Q

Why does Zn2+ appear colourless?

A

Zn2+ has a full d sub-shell so there is no place for the excited electron to move.

134
Q

What colours are seen if there are small or big jumps needed between energy levels?

A

Small jump - needs to absorb a low frequency of energy such as red light, meaning that violet is seen.
Big jump - needs to absorb a high frequency of energy such as violet light, meaning that red light is seen.

135
Q

What does absorbance of light depend on and what is the relationship?

A

Absorbance of a compound at a fixed wavelength is directly proportional to its concentration.

136
Q

What are 2 factors that affect the colour of the complex?

A
  • Charge density of the ligand
  • Concentration of the ligand
137
Q

How does charge density of the ligand impact the colour of the complex?

A

The higher the charge density of the ligand, the greater the split in the d-orbital, so shorter wavelengths of light (red) is absorbed, and longer wavelengths (violet) are seen.

138
Q

How does concentration of the ligand impact the colour of the complex?

A

The higher the concentration of the ligand, the higher the amount of light being absorbed.

139
Q

What is the purpose of a calibration curve?

A

It can help identify the concentration of unknown substances from the absorbance of certain frequencies and amount of light.

140
Q

What are the 4 possible shapes of complex ions?

A
  • Linear shape
  • Square planar shape
  • Tetrahedral shape
  • Octahedral shape
141
Q

What are the properties of a linear complex ion?

A

Coordination number = 2
Bond angle = 180

142
Q

What are the properties of a square planar complex ion?

A

Coordination number = 4
Bond angle = between 90 and 180

143
Q

What are the properties of a tetrahedral complex ion?

A

Coordination number = 4
Bond angle = 109

144
Q

What are the properties of an octahedral complex ion?

A

Coordination number = 6
Bond angle = between 90 and 180

145
Q

What are some examples of common ligands?

A

1) Cl- = chloro (charged)
2) CN- = cyano (charged)
3) H2O = aqua (neutral)
4) NH3 = ammine (neutral)

146
Q

How do you work out the charge on a central metal ion?

A

Charge on central metal ion = charge on complex - charge on ligands
- charge on complex is given, if not it is 0 (charge of the whole compound)
- charge on ligand is the charge of the atom bonded to the central metal atom or the oxidation state

147
Q

What are polydente ligands?

A

Ligands that contain more than 1 lone pair, so can form multiple coordination bonds with the central metal ion.

148
Q

What are chelating agents?

A

Compounds that bind to metal ions through multiple bonds, forming stable complexes.

149
Q

What is chelating?

A

EDTA 4- will form 6 coordinate bonds with the central metal ion to occupy all 6 potential octahedral sites.

150
Q

What is chelating used for?

A

Used to remove poisonous transition metals from solutions and inhibits enzyme catalysed oxidation reactions.

151
Q
A
152
Q
A
153
Q
A

chemistry (22 decks)

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