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Chemistry IB > 3 Periodicity > Flashcards

3 Periodicity Flashcards

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1
Q

In what ways are the elements arranged on the periodic table?

A

elements arranged in order of increasing atomic number, from 1 to 118.

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2
Q

What does the period of an element show?

A

A row of elements thus arranged is called a period. The period number, n, is the outer energy level that is occupied by electrons.

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3
Q

What does the group show?

A

The group shows the number of valence (outer shell) electrons

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4
Q

What 2 elements don’t necessarily fit any groups?

A

Since the electronic configurations of H and He are unusual, they do not fit comfortably into any group. They are thus allocated a group based on similarities in physical and chemical properties with other members of the group

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5
Q

What blocks are there on the periodic table?

A

s,p,d,f blocks

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6
Q

What elements belong to the s-block?

A

The s-block elements are all those with only s electrons in the outer shell

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7
Q

What elements belong to the p-block?

A

The p-block elements are all those with at least one p-electron in the outer shell

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8
Q

What elements belong to the d-block?

A

The d-block elements are all those with at least one d-electron and at least one s-electron but no f or p electrons in the outer shell (up to 5d)

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9
Q

What elements belong in the f block?

A

The f-block elements are all those with at least one f-electron and at least one s-electron but no d or p electrons in the outer shell

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10
Q

What is periodicty?

A

the study of periodic trends in known as periodicity

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11
Q

How would you right the electron configuration of germaninum? (valence shell)

A

4s2 4p2

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12
Q

What is the atomic radius a measure of?

A

The atomic radius of an element is a measure of the size of an atom

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13
Q

Define atomic radius

A

It is the distance between the nucleus of an atom and the outermost electron shell

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14
Q

Is atomic radius precise?

A

It can be quite hard to determine exactly where the boundary of an atom lies, so a variety of approches are taken such as half the mean distance between two adjacent atoms

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15
Q

What factors may affect atomic radius?

A

This will vary depending on the type of structure and bonding, but it gives a comparative value for atoms

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16
Q

How does atomic radius change:

a) down a group
b) across a period

A

a) They generally increase down each group

b) They generally decrease across each period

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17
Q

What can be used to explain periodic trends?

A

These trends can be explained by the electron shell theory

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18
Q

Why do atomic radii decrease as you move across a period?

A

because the atomic number increases (increased positive nuclear charge) but at the same time extra electrons are added to the same principal quantum shell
The larger the nuclear charge, the greater the pull of the nuclei on the electrons which results in smaller atoms

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19
Q

Why do atomic radii increase moving down a group?

A

Atomic radii increase moving down a group as there is an increased number of shells going down the group

The electrons in the inner shells repel the electrons in the outermost shells, shielding them from the positive nuclear charge

This weakens the pull of the nuclei on the electrons resulting in larger atoms

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20
Q

Why does the atomic radius sharply increase at the alkali metals?

A

This is because the alkali metals at the beginning of the next period have one extra principal quantum shell

This increases shielding of the outermost electrons and therefore increases the atomic radius

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21
Q

What is the ionic radius?

A

The ionic radius of an element is a measure of the size of an ion

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22
Q

How does ionic trend change down a group?

A

The trend down a group is the same as atomic radius – it increases as the number of shells increases

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23
Q

What does the trend across a period for ionic radius depend on?

A

The trend across a period is not so straightforward as it depends on whether it is positive or negative ions being considered

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24
Q

How does the trend of ionic radius change with increasing negative charge?

A

Ionic radii increase with increasing negative charge

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25
Q

How does the trend of ionic radius change with increasing positive charge?

A

Ionic radii decrease with increasing positive charge

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26
Q

Why does ionic radius increase with increasing negative charge?

A

Ions with negative charges are formed by atoms accepting extra electrons while the nuclear charge remains the same

The extra electrons experience repulsion with the other valence electrons which increases the ionic radius

The greater the negative charge, the larger the ionic radius

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27
Q

Why does ionic radius decrease with increasing positive charge?

A

Positively charged ions are formed by atoms losing electrons

The nuclear charge remains the same but there are now fewer electrons which undergo a greater electrostatic force of attraction towards the nucleus
which decreases the ionic radius

The greater the positive charge, the smaller the ionic radius

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28
Q

What is the ionisation energy?

A

ionisation energy (IE) of an element is the amount of energy required to remove one mole of electrons from one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

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29
Q

What conditions are ionisation energies measured under?

A

Ionisation energies are measured under standard conditions which are 298 K and 100 kPa

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30
Q

What are the units of IE?

A

kilojoules per mole (kJ mol-1)

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31
Q

What is the first IE?

A

The first ionisation energy is the energy required to remove the one mole of electrons from one mole the gaseous atoms

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32
Q

How could you represent the first IE of Ca?

A

Ca(g) → Ca+ (g) + e–

1st ∆H I.E. = +590 kJ mol-1

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33
Q

Are the IEs of the group 1 metals higher or lower than the noble gases?

A

As could be expected from their electronic configuration, the group I metals show low IE whereas the noble gases have very high IEs

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34
Q

How does first IE change:

a) down a group
b) across a period

A

a) decreases down a group

b) increases across a period

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35
Q

What 4 factors influence first ionisation energy?

A
  1. size of the nuclear charge
  2. distance of outer electrons from the nucleus
  3. shielding effect of inner electrons
  4. spin-pair repulsion
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36
Q

How does the size of the nuclear charge affect the first ie?

A

the nuclear charge increases with increasing atomic number, which means that there are greater attractive forces between the nucleus and outer electrons, so more energy is required to overcome these attractive forces when removing an electron

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37
Q

How does distance of outer electrons from the nucleus affect first ie?

A

electrons in shells that are further away from the nucleus are less attracted to the nucleus so the further the outer electron shell is from the nucleus, the lower the ionisation energy

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38
Q

how does the shielding effect of inner electrons affect first ie?

A

the shielding effect is when the electrons in full inner shells repel electrons in outer shells preventing them to feel the full nuclear charge so the greater the shielding of outer electrons by inner electron shells, the lower the ionisation energy

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39
Q

How does spin-pair repulsion affect the first ie?

A

paired electrons in the same atomic orbital in a subshell repel each other more than electrons in different atomic orbitals; this makes it easier to remove an electron (which is why the first ionization energy is always the lowest)

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40
Q

What factors cause ionisation energy to increase across a period? 3

A

Across a period the nuclear charge increases

The distance between the nucleus and outer electron remains reasonably constant

The shielding by inner shell electrons remains the same

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41
Q

How does the ionisation energy change from the last element in a period and the first element of the next period?

A

There is a rapid decrease in ionisation energy between the last element in one period and the first element in the next period

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42
Q

What 3 factors cause a rapid decrease between element at end of period and next element at beginning?

A

The increased distance between the nucleus and the outer electrons

The increased shielding by inner electrons

These two factors outweigh the increased nuclear charge

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43
Q

Why is there a slight decrease in 1st I.E between beryllium and boron?

A

as the fifth electron in boron is in the 2p subshell which is further away from the nucleus than the 2s subshell of beryllium

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44
Q

Why is there a decrease in the 1st ie between nitrogen and oxygen?

A

between nitrogen and oxygen due to spin-pair repulsion in the 2p subshell of oxygen

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45
Q

What 2 exceptions are there to the increase of first ie across a period?

A
  • beryllium and boron

- nitrogen and oxygen

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46
Q

How does nuclear charge change down a group?

A

going down a group the nuclear charge increases

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47
Q

What factors cause the first ie to decrease down a group?

A

The distance between the nucleus and outer electron increases

The shielding by inner shell electrons increases

The effective nuclear charge is decreasing as shielding increases

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48
Q

What trend is there in successive ionisation energies and why?

A

The successive ionisation energies of an element increase as removing an electron from a positive ion is more difficult than from a neutral atom

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49
Q

What happens in successive i.e when more electrons are removed?

A

As more electrons are removed the attractive forces increase due to decreasing shielding and an increase in the proton to electron ratio

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50
Q

What is the succesive i.e dependent on?

A

The increase in ionisation energy, however, is not constant and is dependent on the atom’s electronic configuration

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51
Q

Why does the first electron removed have a low i.e?

A

The first electron removed has a low ionisation energy as it is easily removed from the atom due to the spin-pair repulsion of the electrons in the x orbital

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52
Q

Why is the second electron harder to remove? (i.e)

A

The second electron is a little more difficult to remove than the first electron as you are removing an electron from a positively charged ion

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53
Q

Why is the 3rd electron even harder to remove? (i.e)

A

The third electron is much more difficult to remove than the second one corresponding to the fact that the third electron is in a principal quantum shell which is closer to the nucleus (3p)

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54
Q

What does a graph of the logarithms of successive i.es show?

A

The graph shows there is a large increase in successive ionisation energy as the electrons are being removed from an increasingly positive ion

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55
Q

What do the big jumps on an i.e graph show?

A

The big jumps on the graph show the change of shell and the small jumps are the change of subshell

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56
Q

What happens when atoms gain electrons?

A

When atoms gain electrons they become negative ions or anions

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57
Q

What is electron affinity?

A

The amount of energy released when one mole of electrons is gained by one mole of atoms of an element in the gaseous state to form one mole of gaseous ions

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58
Q

What is electron affinity the opposite of?

A

ionisation energy

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59
Q

What are the units of EA?

A

The units of EA are kilojoules per mole (kJ mol-1)

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60
Q

What type of reaction is the first ea?

A

The first electron affinity is always exothermic

E.g. the first electron affinity of chlorine is:
Cl (g) + e– → Cl– (g) ∆H = – 349 kJ mol-1

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61
Q

Does the second EA have to be exothermic? Why?

A

no, can be endothermic

This is due to the fact that you are overcoming repulsion between the electron and a negative ion, so energy is required making the process endothermic overall

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62
Q

What trends does EA show and to what other trends are they similar?

A

The pattern is very similar to ionisation energies, except that it is inverted and the minimum points are displaced one element to the right

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63
Q

What does the strongest pull on electrons correlate to?

A

The strongest pull on electrons correlates with the greater amount of energy released when negative ions are formed

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64
Q

What group does not appear on an EA chart?

A

Noble gases do not form negative ions, so they don’t appear in this chart

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65
Q

Where do EA’s reach a peak?

A

The electron affinities reach a peak for group 2 and group 5 elements

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66
Q

How do EA’s change down a group?

A

Electron affinities generally decrease down a group

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67
Q

Why do electron affinities change down a group?

A

As the atoms become larger the attraction for an additional electron is less, since the effective nuclear charge is reduced due to increased shielding

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68
Q

How does the “temp” of the EA’s change down the group?

A

Electron affinity become less exothermic going down the group

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69
Q

What element is an exception?

A

An exception to this is fluorine whose electron affinity is smaller than expected

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70
Q

Why is fluorine an exception?

A

This is because fluorine is such a small atom and an additional electron in the 2p subshell experiences considerable repulsion with the other valence electrons

71
Q

What is electronegativity?

A

Electronegativity is the ability of an atom to attract a pair of electrons towards itself in a covalent bond

72
Q

Why does electronegativity occur?

A

This phenomenon arises from the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself

73
Q

What scale is used to assign a value of electronegativity for each atom?

A

The Pauling scale is used to assign a value of electronegativity for each atom

74
Q

What atom is the most electronegative?

A

Fluorine is the most electronegative atom on the Periodic Table, with a value of 4.0 on the Pauling Scale

75
Q

Why is fluorine so high on the Pauling Scale

A

It is best at attracting electron density towards itself when covalently bonded to another atom

76
Q

Where does attraction occur in an atom?

A

Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom

77
Q

How does an increase in the number of protons affect nuclear attraction?

A

An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells

78
Q

Overall, what does an increased nuclear charge cause relating to EN?

A

Therefore, an increased nuclear charge results in an increased electronegativity

79
Q

How does atomic radius affect attraction?

A

Electrons closer to the nucleus are more strongly attracted towards its positive nucleus
Those electrons further away from the nucleus are less strongly attracted towards the nucleus

80
Q

How does atomic radius affect EN?

A

Therefore, an increased atomic radius results in a decreased electronegativity

81
Q

How does EN change down the group?

A

There is a decrease in electronegativity going down the group

82
Q

How does effective nuclear charge change down a group?

A

We say the effective nuclear charge has decreased down the group

83
Q

Why does electrongeativity decrease down a group? what factors? 3

A

The nuclear charge increases as more protons are being added to the nucleus

However, each element has an extra filled electron shell, which increases shielding

The addition of the extra shells increases the distance between the nucleus and the outer electrons resulting in larger atomic radii

Overall, there is decrease in attraction between the nucleus and outer bonding electrons

84
Q

How does EN change across a period?

A

Electronegativity increases across a period

85
Q

Why does electronegativity increase across a period? (3)

A

The nuclear charge increases with the addition of protons to the nucleus

Shielding remains the same across the period as no new shells are being added to the atoms

The nucleus has an increasingly strong attraction for the bonding pair of electrons of atoms across the period

This results in smaller atomic radii

86
Q

What 6 types of chemical properties can be used to categorise metals/non-metals?

A 
electron arrangement
bonding
electrical conductivity 
type of oxide
reaction with acids
physical characteristics
87
Q

How do non/metals differ via electro arrangement?

A

M - 1-3 (more in periods 5&6) outer shell electrons

N- 4-7 outer shell electrons

88
Q

How do non/metals differ via bonding?

A

M - metallic due to loss of outer shell electrons

n - covalent by sharing of outer shell electrons

89
Q

How do non/metals differ via electrical conductivity?

A

M - good conductors of electricity

N - poor conductors of electricity

90
Q

How do non/metals differ via type of oxide?

A

M - basic oxides (few amphoteric)

N - acidic oxides (some are neutral)

91
Q

How do non/metals differ via reactions with acids?

A

M - many react with acids

N - do not react with acids

92
Q

How do non/metals differ via physical characteristics?

A

M - malleable, high mp and bp

N - flaky, brittle, low mp and bp

93
Q

How are the properties of non/metals explained?

A

The typically properties of metals and non-metals can be explained by reference to their trends in atomic radius, ionic radius, ionisation energy, electron affinity and electronegativity

94
Q

What type of I.E’s and EN’s do metals have? What does this affect?

A

The low ionisation energies and low electronegativities of metals can account for the ability of their valence electrons to move away from the nucleus

This is known as ‘delocalisation‘ of the electrons

95
Q

How do EN’s and I.E’S change from left to right?

A

These properties increase from left to right as you transition from metal to metalloid to non-metal

96
Q

What do the EN’s and I.E’s of non-metals account for?

A

The high electronegativity and electron affinity of non-metals can be related to their tendency to share electrons and form covalent bonds, either with themselves or other non-metal elements

97
Q

What explains the behaviour of metalloids?

A

The similarities in electronegativities of the diagonal band of metalloids which divides the metals from the non-metals explains the behaviour of metalloids

98
Q

What can provide evidence of changing chemical trends?

A

acid-base character of oxides

99
Q

How does the nature of oxides change across a period?

A

The broad trend is that oxides change from basic through amphoteric to acidic across a period

100
Q

What type of oxide is aluminium?

A

Aluminium oxide is amphoteric

101
Q

What does it mean if the oxide is amphoteric?

A

which means that it can act both as a base (and react with an acid such as HCl) and an acid (and react with a base such as NaOH)

102
Q

What does it mean if the oxide is amphoteric?

A

which means that it can act both as a base (and react with an acid such as HCl) and an acid (and react with a base such as NaOH)

103
Q

What can be used to describe the acidic or basic nature of oxides across a period?

A

The acidic and basic nature of the Period 3 elements can be explained by looking at their structure, bonding and the Period 3 elements’ electronegativity

104
Q

Where is the difference between electronegativity the biggest?

A

oxygen and Na, Mg and Al is the largest

105
Q

What type of bond will the metals in period 3 form?

A

Electrons will therefore be transferred to oxygen when forming oxides giving the oxide an ionic bond

106
Q

How will Si, P and S form bonds with oxygen?

A

The oxides of Si, P and S will share the electrons with the oxygen to form covalently bonded oxides

107
Q

Why do metals form alkaline solutions?

A

The oxides of Na and Mg which show purely ionic bonding produce alkaline solutions with water as their oxide ions (O2-) become hydroxide ions (OH–)

108
Q

Why do non-metals form acidic solutions?

A

The oxides of P and S which show purely covalent bonding produce acidic solutions with water because when these oxides react with water, they form an acid that donates H+ ions to water

109
Q

What do metallic oxides form when they react with water?

A

The metallic oxides form hydroxides when they react with water

110
Q

What do non-metallic oxides form when they react with water?

A

The non-metallic oxides form oxoacids when they react with water

111
Q

Practice equations chemical trends

A

ayy

112
Q

What type of property of an element can be predicted by its location on the periodic table?

A

metallic and non-metallic behaviour

113
Q

What type of structure does bonding between metal and non-metals usually have/

A

Metal and non-metal elements generally form ionic compounds so the elements Na to Al have giant ionic structures

114
Q

How does the ionic nature of compounds change down a group?

A

The oxides become more ionic as you go down the group as the electronegativity decreases

115
Q

How does the ionic nature of compounds change across a period?

A

The oxides become less ionic as you go across a period as the electronegativity increases

116
Q

What do oxides of non-metals form?

A

The oxides of non-metals such as S, N and P form molecular covalent compounds

117
Q

Why are the group 1 meta;s called alkali metals?

A

The group 1 metals are called the alkali metals because they form alkaline solutions with high pH values when reacted with water

118
Q

What configuration do all group 1 metals end with?

A

They all end in the electron configuration ns1

119
Q

What are 4 physical properties of group 1 metals?

A

Are soft and easy to cut, getting softer and denser as you move down the group

Have shiny silvery surfaces when freshly cut

Conduct heat and electricity

They all have low melting points and low densities and the melting point decreases going down the group as the atomic radius increases and the metallic bonding gets weaker

120
Q

What will alkaline metals react with?

A

oxygen and water vapour (even in air)

121
Q

What do alkaline metals for, when reacting with water?

A

reacting vigorously to produce an alkaline metal hydroxide solution and hydrogen gas

122
Q

What is the general equation (use m for the metals) between a group 1 metal and water?

A

2M + 2H2O –> 2MOH + H2

123
Q

What is formed when an alkali metals reacts with a halogen?

A

alkali metal halide salt

124
Q

What is the general equation for the reaction between a group 1 metal and a halogen?

A

2M + Cl2 –> 2MCl

125
Q

How does the reactivity of the alkaline metals change down the group?

A

becomes increasingly vigorous going down group 1

126
Q

What are the 3 reasons for the reactivity becoming more vigorous down group 1?

A

The atoms of each element get larger going down the group

This means that the ns1 electron gets further away from the nucleus and is shielded by more electron shells.

The further an electron is from the positive nucleus, the easier it can be lost in reactions

127
Q

Are halogens poisonous?

A

yes

128
Q

What is special about halogens?

A

Halogens are diatomic, meaning they form molecules of two atoms

129
Q

How many valence electrons do halogens have?

A

All halogens have seven electrons in their outer shell

130
Q

What type of ions (special name) do halogens form?

A

They form halide ions by gaining one more electron to complete their outer shells

131
Q

How does the density and mp and bp of the halogens change as you go down the group?

A

The density and melting and boiling points of the halogens increase as you go down the group

132
Q

How does reactivity change going down group 17?

A

Reactivity of group 17 non-metals decreases as you go down the group

133
Q

What does the halogen e- configuration end in?

A

ns2np5

134
Q

What must halogens do to form ions?

A

Each outer shell contains seven electrons and when they react, they will need to gain one outer electron to get a full outer shell of electrons

135
Q

How does electron affinity change going down the group? How is atomic radius affected?

A

Going down the group, the electron affinity decreases and the atomic radius increases

136
Q

How does shielding change while going down the group?

A

As you go down group 17, the number of shells of electrons increases so shielding also increases

137
Q

Why does reactivity decrease going down group 17? (simple)

A

This means that the outer electrons are further from the nucleus so there are weaker electrostatic forces of attraction that attract the extra electron needed

The electron is attracted less readily, so the lower down the element is in Group 17 the less reactive it is

138
Q

When does a halogen displacement react occur?

A

A halogen displacement occurs when a more reactive halogen displaces a less reactive halogen from an aqueous solution of its halide

139
Q

What colour does chlorine have in an aqueous solution?

A

very pale green, but usually appears colourless as it is very dilute

140
Q

What colour does bromine have in an aqueous solution?

A

orange but will turn yellow when diluted

141
Q

What colour does iodine have in an aqueous solution?

A

brown

142
Q

What happens when you add chlorine solution to potassium bromide? Colour change

A

If you add chlorine solution to colourless potassium bromide solution, the solution becomes orange as bromine is formed

Chlorine is above bromine in group 17 so it is more reactive

Chlorine will therefore displace bromine from an aqueous solution of a metal bromide

143
Q

What is the equation for the reaction between chlorine and potassium bromide?

A

2KBr (aq) + Cl2 (aq) → 2KCl (aq) + Br2(aq)

144
Q

What is a transition element?

A

an element whose atom has an incomplete d sub-shell or which can give rise to cations with an incomplete d sub-shell. Ions of transition elements have characteristic properties due to their partially filled d-subshell.

145
Q

Why is Zn not a transition metal?

A

Zn, though part of the d-block, is not a transition metal. The common oxidation state is Zn2+ which does not have a partially filled d-subshell. Instead, both Zn and Zn2+ have a completely filled d-subshell.

146
Q

What causes transition metals to have their unique properties?

A

the characteristic properties of transition metals that

arise from their partially filled d-subshell.

147
Q

What charge ion can all transition metals form and why?

A

Every transition metal can form ions with charge +2 owing to the fact that 4s electrons are lost before any 3d electrons. But many of the transition metals can also occur in other oxidation states: e.g. Fe
Complex ions
2+ 3+ + 2+ and Fe , Cu and Cu
.

148
Q

What is a complex ion?

A

a number of ligands form dative covalent bonds to a positive (metal) ion

149
Q

What is a ligand?

A

a molecule or anion that donates a non-bonding pair of electrons to form a dative covalent bond with a metal ion

150
Q

What is the coordination number?

A

the number of ligands around the central ion

151
Q

What type of complex ions do most transition metals form in water? Give an example with aluminium?

A

Most of the transition metal ions and some non-transition metal ions form hexahydrated complexes in water, such as [Al(H2O6] 3+

152
Q

how are hexahydrated complexes formed?

A

Six water molecules donate one electron paireach, forming six dative covalent bonds

153
Q

What type of structure do hydrated complex ions have?

A

complexes with a coordination number of 6 have an octahedral geometry.

154
Q

What necessary criterion does the ligand need to bond with the metal ion?

A

so long as the ligand can donate an electron pair

155
Q

What ligand would displace water and why?

A

ammonia readily displaces water ligands, because it can form stronger dative covalent bonds than water.

156
Q

What is the spectrochemical series?

A

The order of ligand bond strength is called the spectrochemical series

157
Q

List the spectrochemical series

A

I−

158
Q

Describe the energy levels in a free ion (5-d orbital specifically)

A

In a free ion, the energy levels of each of the 5 d-orbitals are the same, they are said to be degenerate.

159
Q

What causes d-orbitals splitting?

A

But when bonds with a particular geometry form such as described above, the d-electrons closer to the ligands will have a higher energy than those further away, which results in the d-orbitals splitting in energy.

160
Q

Why is d-orbital splitting important for tm?

A

This splitting is what gives transition metals their characteristic properties.

161
Q

Which ligand would cause the greatest d-orbital splitting and why?

A

the CN– ligands cause a larger splitting of the d-orbital energy level because they form stronger bonds

162
Q

What happens when electrons absorb light?

A

electrons can absorb light, thereby moving into an excited state.

163
Q

in a tm, where does an excited e- move from and to?

A

When light is absorbed, an electron moves from a d-orbital with lower energy to a d-orbital with higher energy.

164
Q

why can we “see” the electrons moving between the split d-orbitals?

A

The energy difference between the split d-orbitals is typically in the range of visible light.

165
Q

Why do tm have coloured compounds?

A

So while [Fe(H2O)6]2+ may absorb red photons, [Fe(CN)6]4– will absorb photons with more energy such as blue light, because the energy levels are split further apart.

166
Q

What colour will a tm compound appear?

A

The colour that is observed is the complementary colour to the light that is absorbed, so a solution containing [Fe(H2O)6]2+ will appear green.

167
Q

What does the light absorption of a tm compound depend on? (3)

A

The colour of light absorption (i.e. the amount of splitting of the d-orbitals) thus depends on the identity and oxidation state of the central ion,
the identity of the ligand and the coordination number/geometry around the central ion.

168
Q

How can a SMALL magnetic field be formed? (atoms)

A

When orbitals are partially filled, often (some of) the electrons in the orbitals are unpaired. Each electron has a spin, which when two electrons are paired cancel each other out.

169
Q

How can a bigger magnetic field be created?

A

When electrons are unpaired each produces a tiny magnetic field. When those spins line up in a material they can produce a substantial magnetic field.

170
Q

what does it mean if a metal is ferromagnetic?

A

Fe,Ni,Co show ferromagnetism - permanently magnetic

171
Q

When is a substance paramagnetic?

A

when electrons are unpaired (weak ligand splitting

172
Q

What does it mean if a substance is paramagnetic?

A

slightly attracted by a magnet

173
Q

When is a substance diamagnetic?

A

no unpaired electrons (strong ligand field splitting)

174
Q

What does it mean if a substance is diamagnetic

A

slightly repelled by a magnet

Chemistry IB (13 decks)

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