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IB Chemistry HL > 3 periodicity > Flashcards

3 periodicity Flashcards

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1
Q

groups go from

A

1-18

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2
Q

periods go from

A

1-6

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3
Q

what is the order of block/orbitals in the periodic table

A

s, d, p block

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4
Q

exception to the blocks/orbitals in the periodic table

A

helium

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5
Q

what does n mean

A

the outer energy level that is occupied by electrons

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6
Q

what are the d block elements called

A

transition metals

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7
Q

what are the f block elements called

A

lanthanoids and the actinoids

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8
Q

what is the atomic radius

A

half the distance between the nuclei of two bonded atoms of the same element

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9
Q

across a period atomic radius

A

outer electrons in same shell
more protons in nucleus
same amount of shielding
so stronger attraction between nucleus and outer shell electrons
so outer shell electrons pulled closer to nucleus
DECREASE

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10
Q

down a group atomic radius

A

outer electrons are in a new shell
many more protons in nucleus
much more (whole new energy level) shielding
so weaker attraction between nucleus and outer shell electrons
INCREASE

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11
Q

what is shielding

A

inner electrons shield the outer electrons from the nucleus.
this reduces the effective nuclear charge

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12
Q

why do cations have a greater electrostatic attraction

A

because they have fewer electrons to protons, and one less electron shell, meaning the distance is smaller.

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13
Q

are cations isotonic or not across a period

A

isotonic beacause same number of electrons but an increasing number of protons, so ionic radius decreases.

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14
Q

are anions isotonic or not across a period

A

not because they contain more electrons than protons, so are larger than parent atom. across a period the size decreases because number of electrons remains the same and number of protons increases.

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15
Q

what is first ionization energy

A

amount of energy needed to make a 1+ ion (in gas phase) by removing an electron depends on how strongly that electron is attracted.

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16
Q

what is effective nuclear charge composed of

A

charge on nucleus (number of protons)
shielding from other electrons (shells)

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17
Q

why are paired electrons slightly easier to remove from atoms than unpaired electrons

A

the 2 electrons sharing an orbital will repel each other and have a lower ionisation energy

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18
Q

FIE down a group

A

nuclear charge increases
shielding stays the same
distance from nucleus decreases
INCREASE

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19
Q

IE across a period

A

nuclear charge increases
shielding increases
distance from nucleus increases
DECREASE

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20
Q

why does the type of bonding change across period 3

A

because of the change from metal to non metal oxides, the intermolecualr forces weaken, and they become gas or liquid. the difference in electrnegatiivty between the atoms means the type of bonding cahanges.

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21
Q

amphoteric def

A

can act as both an acid and a base

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22
Q

which period 3 oxides are basic

A

Na2O
MgO

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23
Q

which period 3 oxides are amphoteric

A

Al2O3

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24
Q

which period 3 oxides are acidic

A

SiO2
P4O10 or P4O6
SO3 or SO2
Cl2O7 or Cl2O

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25
Q

what defines basic oxides

A

they dissolve in water to produce basic solutions

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26
Q

what defines acidic oxides

A

they dissolve in water to produce acidic solutions

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27
Q

Na2O (s) + H2O (l) →

A

2NaOH (aq)

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28
Q

MgO (s) + H2O (l)→

A

Mg(OH)2 (aq)

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29
Q

P4O10 (s) + 6H2O (l)→

A

4H3PO4 (aq)

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30
Q

SO3 (g) + H2O (l) →

A

H2SO4 (aq)

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31
Q

3NO2 (g) + H2O (l) →

A

2HNO3 (aq) + NO (g)

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32
Q

Cl2O7 (l) + H2O (l) →

A

2HClO4 (aq)

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33
Q

Cl2O (g) + H2O (l) →

A

2HClO (aq)

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34
Q

alkali metals properties

A

relatively soft
first 3 have low enough densities to float on water
melting and boiling points decrease down the group as the metallic bond gets weaker as ionic radii increase
stored in oil to prevent reactions with air
reactivity increases down the group

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35
Q

low ionisation energies of alkali metals mean they are reactive. what do they form with water

A

metal hydroxides and hydrogen gas

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36
Q

properties of halogens

A

reactive group of non metal elements
reactivity decreases down the group

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37
Q

chlorine properties

A

dense pale green gas
smelly and poisonous
occurs as a chloride in the sea
strong oxidising agents

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38
Q

bromine properties

A

deep red liquid with red brown vapour
smelly and poisonous
occurs as a bromide in the sea

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39
Q

iodine properties

A

grey solid with purple vapour
smelly and poisionous
occurs as iodidies and iodates in some rocks and in seaweed

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40
Q

what is a displacement reaction

A

in which the more reactive halogen displaces the ions of the less reactive halogen from solution.

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41
Q

what happens whrn chlorine is added to a colourless solution of potassium bromide

A

it turns brown due to the formation of aqueous bromine. cyclohexane allows the change to be seen more visibly, as because it is polar, it will forma layer above the reaction mixture

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42
Q

what happens when an aqueous solution of chlorine is added to a solution containing iodide ions

A

the iodide ions have been oxidised eto form aqueous iodide. a purple layer is observed in the layer of cyclohexane because of the formation of iodine with dissolves in cyclohexane.

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43
Q

do Al2O3 and SiO2 react well and are soluble

A

no poorly reactive and poorly soluble

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44
Q

structure of Na2O, MgO and Al2O3

A

ionic

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45
Q

structure of SiO2

A

covalent macromolecule (which is why its so insoluble)

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46
Q

structure of P4O10, SO2, SO3

A

simple covalent

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47
Q

which period 3 oxides are alkaline

A

Na2O
mgO

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48
Q

which period 3 oxide is amphoteric

A

Al2O3
alumina

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49
Q

which period 3 oxides are acidic

A

SiO2
P4O10
SO2
SO3

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50
Q

what happens in the reaction of Na2O and water

A

forms sodium hydroxide

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51
Q

what happens in the reaction of MgO

A

forms magnesium hydroxide

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52
Q

what happens in the reaction of Al2O3 with water

A

neutralises the acid

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53
Q

what happens in the reaction of Al2O3 with sodium hydroxide

A

makes a neutral salt (solution remains neutral)

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54
Q

what happens in the reaction of SiO2 with sodium hydroxide

A

neutralises (hence why we classify it as a base)

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55
Q

what happens in the reaction of P4O10 with water

A

forms phosphoric acid H3PO4

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56
Q

what happens in the reaction of SO2 with water

A

forms sulfurous acid
H2SO3

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57
Q

what happens in the reaction of SO3

A

forms hypochlorous acid

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58
Q

what is Z

A

atomic number

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59
Q

what is nuclear charge

A

total chare of all the protons in the nucelus of an atom

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60
Q

group 1 is

A

alkali metals

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61
Q

group 2 is

A

alkaline earth metals

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62
Q

group 17 is

A

halogens

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63
Q

goup 18 is

A

noble gases

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64
Q

what are lanthanides and actinides

A

metallic elements that make up the f block

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65
Q

the s block occupies

A

groups 1 and 2

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66
Q

the p block occupies

A

group 13 to 18

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67
Q

d block occupies

A

group 3 to 12

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68
Q

f block occupies

A

elements at the bottom

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69
Q

what is the covalent radius measured as

A

half the distance between two neighbouring nuclei

70
Q

what does effective nuclear charge depend on

A

The distance between the nucleus and the electrons (the atomic radius).
The number of shielding electrons within the atom (shielding electrons are inner electrons that tend to ‘shield’ the outer electrons from the full attraction of the nucleus).

71
Q

what is the attraction felt by valene electrons called

A

effective nuclear charge

72
Q

what is first ionisation energy

A

the energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions.

73
Q

berylilliym electron configuration

A

1s2 2s2

74
Q

boron electron configuration

A

1s2 2s2 2p1

75
Q

why do electrons in p orbitals require less energy to move

A

they are of higher energy and further from the nucleus than electrons in s orbitals

76
Q

fie increases along a period except for…

A

there is a decrease in IE from beryillium to boron because the one electron in the p orbital in boron takes less energy to remove.

77
Q

how is electronegativity measured

A

on the pauling scale

78
Q

what is first electron affinity

A

energy released when one mole of electrons is added to one mole of gaseous atoms to form one mole of gaseous 1− ions.

79
Q

why are first electron affinities exothermic

A

becuaes energy is released when an electron is added to a neutral atom

80
Q

why do second electron affinities tend to be endothermic

A

because of the rpulsion that occurs when a seecond electron is already added to an already negative ion

81
Q

why do elements in group 17 have the highest electron affinitiees

A

These relatively small atoms can accommodate an electron in their unfilled outer shell and, once there, it is strongly attracted to the nucleus of the atom.

82
Q

how does structure change across a period

A

metallic to giant covalent to molecular covalent

83
Q

where does metallic charayer decrease

A

across a period

84
Q

where does metallic character increase

A

down a group

85
Q

what causes metallic character to decrease along a period

A

The metallic character of an element is related to how easily it loses its valence electrons. The decrease in metallic character across a period is due to the decreasing atomic radius and increasing nuclear charge which results in a stronger attraction between the nucleus and the valence electrons.

86
Q

why do melting points increase down a group

A

due to an increase in London dispersion forces caused by increasing molar mass.

87
Q

whya re d block elements called d block

A

they have valence electrons in the d sub level

88
Q

transiiton metal def

A

defined as an element that has an incomplete d sub-level in its atom or one or more of its ions.

89
Q

properties of transiiton metals

A

They have variable oxidation states.
They form coloured compounds.
The elements or their compounds show catalytic activity.
They form complex ions in solution.
The metals and their complexes show magnetic properties.

90
Q

waht does a catalyst do

A

increase rate of chemical reaction by providing an alternative reaction pathway with a lower activation energy (Ea)

91
Q

Transition metals play an important role as catalysts in the chemical industry by allowing

A

chemical reactions to take place at lower temperatures and pressures than they would otherwise need.

92
Q

transition metals are effective ______________ catalysts

A

heterogenous

93
Q

haber process equation

A

N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

94
Q

hydrogenation converts

A

alkenes to alkanes

95
Q

contact process for making sulfuric acid

A

2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

96
Q

Manganese(IV) oxide (MnO2) for decomposing hydrogen peroxide into water and oxygen equation

A

2H2O2 (aq) → 2H2O (l) + O2 (g)

97
Q

waht is a homogenous catalyst

A

one that has the same physical state as the reactants

98
Q

why are transtion metals also effectice as homogenous catalysts in redox reactions

A

due to their ability to have variable oxidation states

99
Q

Examples of transition metals as homogeneous catalysts include:

A

Enzymes, which are biological catalysts, enable chemical reactions within human body cells to take place at a lower temperature than would otherwise be necessary. Transition metals such as iron and cobalt are important components of these enzymes.
The iron(II) ion is central to the heme group in hemoglobin, which is responsible for carrying oxygen around the body (section B.9.3).
Cobalt(III) ions are found in vitamin B12, which is essential to maintain good health.

100
Q

why are d block elements called that

A

their valence electrons are in the d sub level

101
Q

transition metal

A

is defined as an element that has an incomplete d sub-level in its atom or one or more of its ions.

102
Q

properties of transition metals

A

They have variable oxidation states.
They form coloured compounds.
The elements or their compounds show catalytic activity.
They form complex ions in solution.
The metals and their complexes show magnetic properties.

103
Q

scandium oxidation state

A

+3

104
Q

titanium oxidation state

A

+2,3,4

105
Q

vanadium oxidation state

A

+2,3,4,5

106
Q

chromium oxidation state

A

+2,3

107
Q

maganese oxidation state

A

+2,3,4

108
Q

iron oxidation state

A

+2,3

109
Q

cobalt oxidation state

A

+2,3

110
Q

nickel oxidation state

A

+2

111
Q

copper oxidation state

A

+1,2

112
Q

zinc oxidation state

A

+2

113
Q

why do oxidation states change in transition metals

A

closeness in energy of the 3d and 4s sub levels
there is little change in the atomic radii as we proceed from left to right. This is due to the fact that electrons are being added to the inner sub-level (3d), meaning that the effective nuclear charge experienced by the outer valence electrons in the 4s sub-level remains fairly constant.

114
Q

why si there small range in the value of the first ionisation energy across the first row of the d block

A

The atoms of these d-block elements are of similar size, and the ‘pull’ of the nucleus (or effective nuclear charge) on the outer 4s electrons is also similar.

115
Q

what is a ligand

A

species with lone pairs of electrons that act as Lewis bases in that they donate a lone pair of electrons to the central metal ion.

116
Q

when is acompelx ion formed

A

when ligands bond to a central metal ion
The central metal ion acts as a Lewis acid by accepting the lone pair of electrons from the ligand, forming a coordinate covalent bond in the process.

117
Q

The number of coordinate covalent bonds formed between the ligands and the central metal ion is known as

A

the coordiantion number

118
Q

monodentate ligands

A

form one coordinate b ond

119
Q

bidentate ligands

A

formt wo coordiante bonds

120
Q

hwo many coordiante bonds does EDTA 4- form

A

6 (hexadenate)

121
Q

coordination compound

A

consists of a complex ion together with counter ions that balance out the charge of the complex ion.

122
Q

coordination compounds and charges

A

The negative charges on the counter ions balance out the positive charge of the complex ion. A second example is K3[Fe(CN)6] in which the counter ions, K+, balance out the 3- charge on the complex ion. An important point to note is that species inside the square brackets are ligands (except for the metal ion), whereas those outside the brackets are counter ions.

123
Q

coordination compounds and charges

A

The negative charges on the counter ions balance out the positive charge of the complex ion. A second example is K3[Fe(CN)6] in which the counter ions, K+, balance out the 3- charge on the complex ion. An important point to note is that species inside the square brackets are ligands (except for the metal ion), whereas those outside the brackets are counter ions.

124
Q

what causes types of magnetism on trnasiiton elemetns

A

The presence of unpaired electrons in the 3d sub-level of transition metal atoms or ions gives rise to two characteristic properties of these elements: their usefulness as catalysts and their magnetic properties.
wo electrons in the same atomic orbital have opposite spins and are represented by single-headed arrows, one pointing up and one pointing down. Electron spins give rise to magnetic effects

125
Q

Substances with only paired electrons are classified as being

A

diamagnetic
If an atomic orbital contains two electrons, the opposite spins cancel out the magnetic effect and the substance is diamagnetic. Diamagnetic materials show a weak repulsion in an external magnetic field; however, we normally think of these substances as being non-magnetic. Pyrolytic carbon (similar to graphite) is diamagnetic.

126
Q

Substances with half-filled atomic orbitals (containing only one electron) exhibit

A

paramagnetism. Paramagnetic materials are attracted to an external magnetic field, however they do not retain their magnetic properties when the external field is removed. The greater the number of unpaired electrons, the stronger the attraction.

127
Q

the unpaired electrons become aligned with an external magnetic field and this alignment persists even after the external field is removed – the object itself becomes magnetised.

A

ferromagnetism (the nromal type and only certian metalds can exhibit it)

128
Q

Transition metals are effective heterogeneous catalysts. In heterogeneous catalysis, the catalyst is

A

in a different state to that of the reactants.

129
Q

Iron (Fe) in the Haber process, producing ammonia:

A

N2 (g) + 3H2 (g) ⇌ 2NH3 (g)

130
Q

Vanadium(V) oxide (V2O5) in the Contact process for making sulfuric acid. Specifically, it is utilised to increase the rate at which SO2 is oxidised to SO3 in the first step of the process:

A

2SO2 (g) + O2 (g) ⇌ 2SO3 (g)

131
Q

Manganese(IV) oxide (MnO2) for decomposing hydrogen peroxide into water and oxygen:

A

2H2O2 (aq) → 2H2O (l) + O2 (g)

132
Q

homogeneous catalyst

A

one that is in the same physical state as the reactatns

133
Q

transition elelment def

A

an element that has an incomplete d sub-level in its atom or one or more of its ions.

134
Q

why is copper considered to be a transition element

A

when copper forms the Cu2+ ion, it has an incomplete d sub-level and is therefore considered to be a transition element.

135
Q

why is zinc not considered to be a transition element

A

as both zinc atoms and zinc ions (Zn2+) have complete d sub-levels.

136
Q

properties of transititon elements

A

They have variable oxidation states.
They form coloured compounds.
The elements or their compounds show catalytic activity.
They form complex ions in solution.
The metals and their complexes show magnetic properties.

137
Q

why do all transition elements have a +2 ox state except scandium

A

they lose their 4s electrons to form 2+ ions (chromium and copper lose one electron from the 4s sub level and the second from the 3d sub level when they form 2_ ions because they only have one electron in the 4s sub level

138
Q

why is there a large increase in ioisation energy between the second nand third ionisation energies in magnesium

A

The increase is due to the third electron being removed from the 2p sub-level, which is closer to the nucleus than the 3s sub-level. This large increase in energy is the reason why magnesium does not form a Mg3+ ion.

139
Q

why does Titanium shows a large increase between the 4th and 5th ionisation energy

A

due to the 5th electron being removed from the 3p sub-level. Because of the closeness in energy between the 4s and 3d sub-levels, titanium can have an oxidation state of +2, +3 and +4. However, the extra energy required to remove the 5th electron from the 3p sub-level means that titanium does not have an oxidation state of +5.

140
Q

ligans are

A

species with lone pairs that act as lewis bases (donat ea lone pair to central metal ion

141
Q

the central metal ion acts as a

A

lewis acid by accepting the lone pair from the ligand forming a coord covalne tbond in the process

142
Q

what is the coordination number

A

The number of coordinate covalent bonds formed between the ligands and the central metal ion

143
Q

A coordination compound consists of a

A

complex ion together with counter ions that balance out the charge of the complex ion.

144
Q

dimagnetism

A

only paired electrons .
weak repulsion in magnetic fields so basically non magnetic

145
Q

parapmagentism

A

half filled orbitals. attracted to magnetic fields but do not retain magnetic properties when field is removed. greater the number of unparied electrons, the stronger the attraction

146
Q

ferromagnetism

A

unparied electrons become aligned with an external magnetic field and this alignment persists even after the external field is removed.

147
Q

are transition metals magnetic

A

they have unpaired electrons and are attracted to a amgnetic firled, so yes paramagnetic

148
Q

what type of catalyst are transition metals

A

heterogenous and homogenous

149
Q

what is a heterogenous catalyst

A

the catalyst is in a different state to that of the reactants.

150
Q

examples of transition elements as homogenous catalysts

A

Enzymes, which are biological catalysts, enable chemical reactions within human body cells to take place at a lower temperature than would otherwise be necessary. Transition metals such as iron and cobalt are important components of these enzymes.
The iron(II) ion is central to the heme group in hemoglobin, which is responsible for carrying oxygen around the body (section B.9.3).
Cobalt(III) ions are found in vitamin B12, which is essential to maintain good health.

151
Q

examples of transtiion metals as heterogenous catalysts

A

iron in the haber process - the reactants hydrogen and nitrogen are in the gaseous state with the iron being in the solid state. The iron provides a surface on which the reactant molecules can adsorb and so come together with the correct orientation to react

nickel in the conversion of alkenes to alkanes

vanadium oxide in the contact process for making sulfuric acid

152
Q

what is ligand exchange

A

If some of the water molecules are replaced by ammonia molecules, then the wavelength of light absorbed changes. occurs when excess aqueous ammonia is added to a solution of copper(II) sulfate. The [Cu(NH3)4(H2O)2]2+ complex ion is produced which is a darker blue colour in solution than the [Cu(H2O)6]2+ complex ion. This is because it absorbs light of shorter wavelength, which corresponds to yellow and green light (this is discussed in more detail later in this section).

153
Q

what dos degenerate mean

A

d orbitals are of equal energy

154
Q

E = hv where e is

A

energy in joules

155
Q

E = hv where h is

A

plancks constant (6.63 x 10 -34)

156
Q

E = hv where v is

A

frequency of light in /s

157
Q

diagram of degenerate 3d orbitals splitting

A
158
Q

what is the spectrochemical series

A

the energy difference between non degenerate d orbitals in a transition metal ion is determined by several factors. one of these factors is the identity of the ligands that surround hte transition metal ion. The effect of different ligands on the degree of d orbital splitting in an octahedral complex, and therefore the size of the ∆E, is given in the spectrochemical series. Ligands higher up in the spectrochemical series, such as CN– and CO, produce a larger splitting of the d orbitals and ligands lower in the spectrochemical series, such as I– and Br–, produce a smaller splitting of the d orbitals.

159
Q

which ligands replace which

A

stronger field ligands will replace weaker field ligands in solution

160
Q

weakest field ligands

A

cause least splitting of d orbitals
absorb longer wavelengths of light

161
Q

strongest field ligands

A

cause greatest splitting of d orbitals
absorb shorter wavelengths of light

162
Q

a larger energy gap corresponds to

A

a shorter wavelength of light being absorbed and the solution appears a darker blue colour

163
Q

factors that determine the colour of complex ions are

A

The identity of the metal ion at the centre of the complex. This affects the number of electrons involved in the d orbitals, the oxidation state of the metal and also the nuclear charge at the centre of the ion.
The geometry of the complex ion. This affects the strength of the repulsion between the ligands and the d orbital electrons.
The oxidation state of the transition metal ion. For example, a complex ion with Fe2+ as the central metal ion will appear a different colour to a complex ion with Fe3+ as the central metal ion (assuming that the geometry and ligands are the same). This is because of the different oxidation state in iron(II), which is +2, compared to iron(III), which is +3.

164
Q

concentration against time graphs for zero, first and second order reactants

A
165
Q

rate against concentration for zero, first and second order reactants

A
166
Q

what does the molecualrity of a reacion tell you

A

The molecularity of a reaction tells us the number of reactant particles in an elementary step.

167
Q

fif one particle is involved in the elementary step

A

unimolecular

168
Q

if two reactant particles are involved it is

A

bimolecular

169
Q

if three reactant particles are involved

A

termolecular

170
Q

moelcualrity diagram

A
171
Q

how to deduce rds from a graph

A

transition state 1 has the highest activation energy, therefore, this is the rate-determining step (the step with the highest activation energy). This transition state is unstable and can either decompose to reform the reactants or go on to form the intermediate. The dip in the energy level profile shows the formation of the intermediate, which then goes on – via transition state 2 – to form the products. Although the transition states and intermediate are short-lived, only the intermediate can be isolated because it has bonds that are fully formed, as opposed to the partial bonds present in the transition state.

IB Chemistry HL (12 decks)

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