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IB Chemistry HL > Chapter 3/ 13 > Flashcards

Chapter 3/ 13 Flashcards

Periodicity

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1
Q

Elements in periodic table

A

Arranged in order of increasing atomic number (Z)

- each successive element has one more proton than the last

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2
Q

Periodicity

A

Repeating patterns of chemical and physical properties

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3
Q

Importance of atomic number

A
  • defining property of an element
  • came from Moseley’s studies of X-rays released when atoms of different metallic elements were bombarded by electrons
  • discovered that frequency of X-rays released showed a direct relationship to atomic number of element
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4
Q

Groups

A

Vertical columns

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5
Q

Periods

A

Horizontal rows

- period no. = no. of occupied main energy levels in the atom

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6
Q

Group 1 elements

A

Alkali metals

  • most reactive group of metals
  • react strongly with cold water
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7
Q

Group 2 elements

A

Alkaline earth metals

  • reactive metals
  • less reactive than group 1
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8
Q

Group 17 elements

A

Halogens

- most reactive group of non-metals

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9
Q

Group 18 elements

A

Noble gases: so called because of their lack of reactivity

- group of very unreactive monoatomic gases

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10
Q

Metalloids

A

Located in a ‘diagonal’ staircase that forms a boundary between metals and non-metals
- have properties of both metals and non-metals

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11
Q

Covalent radius and metallic radius

A

Measured as half the distance between two neighbouring nuclei

  • valent radii and metallic radii can be obtained for most elements (except those that don’t from compounds)
  • used for comparison across a period in the periodic table
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12
Q

van der Waal’s radius

A
  • atomic radii values quoted for the noble gases (group 18)
  • they don’t form compounds
  • larger than covalent radii
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13
Q

Atomic radius trends

A

Across a period:

  • atomic radii of atoms decreases across the period
  • no. of protons increases by one with each successive element, hence, no. of electrons also increases by one
  • extra electron occupies the same main energy level, so, electron shielding remains roughly constant across a period
  • results in stronger attraction between nucleus and valence electrons, pulls them closer to nucleus

Down a group:

  • atomic radii increases down a group
  • each additional period means valence electrons occupy a main energy level further from nucleus
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14
Q

Electron shielding

A

Occurs when outer electrons are shielded from attraction of nucleus by inner electrons (shielding electrons)

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15
Q

Ionic radius

A

Metallic elements tend to lose electron to form cations

Non-metal elements gain electrons to form anions

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16
Q

Isoelectronic species

A

Ions that have different numbers of protons, as they are different elements, but the same number of electrons

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17
Q

Ionic radius trends (positive ions)

A

Positive ions:

  • Lose electrons from valence shell
  • each ion still has same no. of protons in nucleus as parent atom, but are pulling on fewer electrons
  • increases attraction between nucleus and valence electrons
  • hence, positive ions are smaller than their parent atoms
  • So, trend across a period is that ionic radii decrease as nuclear charge increases
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18
Q

Ionic radius trends (negative ions)

A

Negative ions:

  • gain electrons; have more electrons than protons
  • decreases attraction between nucleus and valence electrons; weaker attraction causes ionic radius to increase
  • these extra electrons also increase repulsion between electrons, contributes to increase in ionic radii
  • across period, ionic radii of negative ions decrease as nuclear charge increases
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19
Q

Electrostatic attraction

A

The force of attraction between the positively charged nucleus and negatively charged electrons within an atom

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20
Q

Strength of electrostatic attraction depends on:

A
  1. Atomic radius (distance between nucleus and electrons)

2. Number of shielding electrons within the atom

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21
Q

Shielding electrons

A

Inner electrons that tend to ‘shield’ the outer electrons from the full attraction of the nucleus

Hence, valence electrons don’t feel full attraction from protons in the nucleus

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22
Q

Effective nuclear charge

A

Attraction felt by valence electrons

- is less than the actual nuclear charge of the atom

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23
Q

Trends in effective nuclear charge

A

Across a period:

  • no. of protons increases by one for each successive element, but no. of electrons in inner energy levels doesn’t change
  • effective nuclear charge increases by one until group 18

Down a group:

  • effective nuclear charge remains approx. constant moving down a group
  • attraction from increasing no. of protons in nucleus is offset by increase in occupied energy levels down group
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24
Q

First ionisation energy

A

Energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions
- measure of attraction between positively charged nucleus and negatively charged outer valence electrons

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25
Q

First ionisation energy trends

A

Across a period:

  • increase across a period
  • due to increase in nuclear charge and decrease in atomic radii
  • results in an increased attraction between nucleus and valence electrons

Down a group:

  • decrease down a group
  • due to increase in atomic radius
  • increased distance results in a weaker attraction between nucleus and valence electrons of an atom
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26
Q

Exceptions to the ionisation energy trend across a period

A

IE decreases from:
- Beryllium to boron
electrons in p orbitals are higher in energy and further from nucleus than electrons in s orbitals, hence less energy is required to remove

  • Magnesium to aluminum
    electrons in p orbitals are higher in energy and further from nucleus than electrons in s orbitals, hence less energy is required to remove
  • Nitrogen to oxygen
    electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by other electron and requires less energy to remove than an electron in a half-filled orbital
  • Phosphorus to sulphur
    electron is removed from a doubly occupied p-orbital. An electron in a doubly occupied orbital is repelled by other electron and requires less energy to remove than an electron in a half-filled orbital (electrons are being removed from 3p sub-level)
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27
Q

Electronegativity

A

Attraction of an atom for a bonding pair of electrons

- measured on Pauling scale

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28
Q

Electronegativity trends

A

Across a period:

  • increases across a period duet to increase in nuclear charge and decrease in atomic radius
  • results in a stronger attraction between nucleus and bonding electrons

Down a group:

  • electronegativity decreases down a group
  • due to increase in atomic radius, increasing distance between nucleus and shared pair of electrons, hence, attraction for these electrons decreases
  • increase in nuclear charge down a group is counteracted by increased shielding caused by extra occupied main energy levels within the atom
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29
Q

Electronegativity and bonding

A

Difference in electronegativity between atoms in a compound determines type of bonding that occurs

  • difference up to 1.7 units = forms covalent bond
  • difference of 1.8 units or more = forms ionic bond
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30
Q

Electron affinity

A

First electron affinity: energy released when one mole of electrons are added to one mole of gaseous atoms to form one mole of gaseous 1- ions
- exothermic

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31
Q

Group 17 elements and electron affinities

A

Group 17 elements have highest electron affinities (most exothermic)

  • relatively small atoms can accommodate an electron in their unfilled outer shell
  • hence, it’s strongly attracted to the nucleus of the atom
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32
Q

Electron affinity trends

A

Across a period:

  • values increase, becoming more exothermic
  • progressing to the filling of the outer valence shell of the atom
  • group 17 atom releases more energy than a group 1 atom on gaining an electron
  • on gaining an electron it achieves a filled valence shell, is more stable

Down a group:

  • values decrease
  • additional electron gained enters an energy level further from nucleus
  • added electron has a weaker attraction to nucleus, releases less energy when added
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33
Q

Metals and ionisation energies

A
  • have low IE

- lose their valence electrons relatively easily to form positive ions

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34
Q

Non-metals and ionisation energies

A
  • high ionisation energies

- gain electrons to form negative ions

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35
Q

Metallic character trends

A

Across a period:

  • metallic character decreases
  • due to increase in IE across a period

Down a group:
- metallic character of elements increases as IE decreases

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36
Q

Metallic character

A
  • strongly related to its ionisation energy
  • IE depends on nuclear charge and atomic radius of an atom
  • higher the nuclear charge, smaller the atomic radius of an atom, lower its metallic character and vice versa
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37
Q

Bonding trends

A

Across a period, bonding and structure change from:

Metallic –> Giant covalent –> Molecular covalent

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38
Q

Period 3 oxides

A

Show gradual trend of decreasing metallic character across the period
- type of bonding changes from ionic to covalent across period, determined by difference in electronegativity between bonding atoms

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39
Q

Metal oxides

A
  • giant ionic structures

- solids under standard conditions due to strong electrostatic attractions between ions

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40
Q

Non-metal oxides

A
  • molecular covalent structures

- gases of liquids under standard conditions

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41
Q

Acid-base properties of period 3 oxides

A
  • change from basic through amphoteric to acidic across a period
  • metal oxides form metal hydroxides when reacted with water
  • non-metal oxides form acidic solutions when reacted with water
42
Q

Amphoteric

A

Can act as both an acid and a base

43
Q

Basic oxides

A

Dissolve in water to produce basic solutions

44
Q

Acidic oxides

A

Dissolve in water to produce acidic solutions

45
Q

Properties of alkali metals

A
  • very reactive
  • soft (can be easily cut with a knife; reveals a shiny surface that quickly tarnishes as a layer of oxide is formed)
  • MP and BP are relatively low, decrease down the group (due to metallic bonding getting weaker as ionic radii of metals increase)
  • stored in oil to prevent reaction with oxygen in the air
  • one electron in valence shell, lose relatively easily (low IE) to form 1+ ions with halogens
  • undergo vigorous reactions with water to form a metal hydroxide and hydrogen gas; resulting solution has high pH
  • reactivity increases down the group
46
Q

Properties of halogens

A
  • very reactive group of non-metals
  • characteristic colours
  • physical state changes down the group (related to increasing molar mass of halogens, results in stronger intermolecular forces between molecules)
  • diatomic; they exist as two atoms bonded together, rather than separate atoms
  • MP and BP increases down the group, related to increasing strength of intermolecular forces
  • reactivity of halogens decreases down the group
  • react with group 1 elements to form salts
  • undergo displacement reactions with each other
47
Q

Properties of chlorine

A
  • dense pale-green gas
  • smelly and poisonous
  • occurs as chloride in the sea
48
Q

Properties of bromine

A
  • deep-red liquid with red-brown vapour
  • smelly and poisonous
  • occurs as bromide in the sea
49
Q

Properties of iodine

A
  • grey solid with purple vapour
  • smelly and poisonous
  • occurs as iodides and iodates in some rocks and seaweed
50
Q

d-block elements

A
  • occupy central region of periodic table

- so called because they have valence electrons in d sub-level

51
Q

Transition elements

A

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

  • found in the d-block region (excludes zinc)
  • show and oxidation state of +2 when the s-electrons are removed
52
Q

Scandium as a transition metal

A
  • transition element because of electron configuration of atom [Ar] 3d1 4s2
  • Sc3+ ion has no 3d electrons at all, as atom loses both 4s electrons and one 3d electron when forming the ion
53
Q

Copper as a transition metal

A
  • transition element
  • although Cu and Cu+ have complete 3d sub-levels, copper (II) ion (Cu2+) has electron configuration [Ar]3d9, signifies an incomplete d sub-level
54
Q

Zinc as a transition metal

A
  • not a transition metal because it doesn’t form ions w/ incomplete d-orbitals
  • both atom and ion formed (Zn2+) have a complete 3d sub-level
55
Q

Characteristic properties of transition elements

A

Presence of unpaired electrons in d sub-level gives rise to characteristic properties of transition elements

  • variable oxidation states
  • have coloured compounds
  • elements display catalytic and magnetic properties
  • form complex ions w/ ligands
56
Q

Variable oxidation states

A

Reason:

  • closeness in energy of 3d and 4s sub-levels
  • there’s little change in atomic radii from left to right, as electrons are added to inner sub-level (3d)
  • so, effective nuclear charge experienced by outer valence electrons in 4s sub-level remains constant
  • atoms of d-block elements are of similar size, and effective nuclear charge on outer 4s electrons is also similar
  • results in there being only a small range in values of first IE across first row of d-block

Transition metals display a wide range of oxidation states in their compounds
- contrast to s-block metals, only show oxidation states of +1 (for Group 1) and +2 (for Group 2)

57
Q

Copper (II) sulphate solution

A
  • solution is blue due to presence of hydrated copper (II) ions
  • copper (II) ions in solution are surrounded by 6 water molecules (ligands), results in formation of complex ions
58
Q

Complex ion

A

Made up of a central metal ion surrounded by ligands

59
Q

Ligands

A

Species with lone pairs of electrons that act as Lewis bases, they donate a lone pair of electrons to the central metal ion

  • can be neutral molecules, w/ lone pairs of electrons or charged ions
  • all ligands are lone pair donors (all ligands act as Lewis bases)
60
Q

Central metal ion (in a complex ion)

A

Central metal ion = Lewis acid

- accept lone pair of electron from ligand, forming a coordinate covalent bond

61
Q

Coordinate covalent bond

A

Bond formed between a ligand and the central metal ion

62
Q

Coordination number

A

Number of coordinate covalent bonds formed between the ligands and central metal ion
- no. of coordinate bonds to one central ion

63
Q

Monodentate ligands

A

Ligands that form just one coordinate covalent bonds

64
Q

Bidentate ligands

A

Ligands that form two coordinate bonds from each ion or molecule to the central metal ion

65
Q

Overall charge on a complex ion

A

Sum of the charges of the central metal ion and the charges contributed by the ligands

66
Q

Shapes of complex ions

A

Depend on no. and size of ligands involved

look at notes for shapes

67
Q

Catalytic behaviour

A

Presence of unpaired electrons in 3d sub-level of transition metal atoms or ions gives rise to two characteristic properties of these elements

  • catalysts
  • magnetic properties
68
Q

Catalyst

A

Increases rate of a chemical reaction by providing an alternative reaction path w/ a lower activation energy (Ea)

69
Q

Transition elements as catalysts

A

Play an important role as catalysts in chemical industry

- allow chemical reactions to take place at lower temp. and pressures than they would otherwise need

70
Q

Heterogenous catalysis

A

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

  • Transition elements are effective heterogeneous catalysts
  • eg. use of Iron in Haber process
71
Q

Haber process

A

Reactants nitrogen and hydrogen are in gaseous state w/ iron being in solid state

  • iron provides a surface on which reactant molecules can adsorb
  • hence, come together w/ correct orientation to react
72
Q

Homogeneous catalyst

A

A catalyst that is in the same physical state as the reactants
- transition metals are also effective as homogeneous catalysts in redox reactions due to their ability to have variable oxidation states

73
Q

Examples of transition metals as homogeneous catalysts

A
  • enzymes (biological states) enable chemical reactions within human body cells to take place at lower temp. than would be otherwise necessary (iron and cobalt are important components of these enzymes)
  • iron (II) ion is central to haeme group in haemoglobin, responsible for carrying oxygen around the body
  • cobalt (III) ions are found in vitamin B12, essential to maintain good health
74
Q

Electrons and orbitals

A

Two electrons in same atomic orbital have opposite spins- represented by single-headed arrows

  • electron spins give rise to magnetic effects
  • classified as: dimagnetism, paramagnetism or ferromagnetism
  • if an atomic orbital contains two electrons, opposite spins cancel each other out
75
Q

Dimagnetism

A

Substances with any no. of paired electrons
- diamagnetic materials show a weak repulsion in an external magnetic field (these are normally considered as non-magnetic)

76
Q

Pyrolytic carbon

A
  • similar to graphite

- is diamagnetic

77
Q

Paramagnetism

A

Substances w/ half-filled atomic orbitals (contain only one electron)

  • paramagnetic materials are attracted to an external magnetic field
  • greater no. of unpaired electrons, stronger the attraction
78
Q

Ferromagnetism

A
  • the largest effect
  • unpaired electrons become aligned with an external magnetic field
  • this alignment persists even after the external field is removed- object itself becomes magnetised
  • only certain metals, some rare earth elements and their alloys are ferromagnetic
79
Q

Transition metal and paramagnetism

A
  • transition metal complexes have unpaired electrons

- attracted to a magnetic field, showing paramagnetic properties

80
Q

Dimagnetism vs. Paramagnetism

A 
Dimagnetic = substances with only paired electrons
Paramagnetic = substances with unpaired electrons
81
Q

Complex ions and d-orbitals

A

d-orbitals have same energy in an isolated atom

  • but split into two sub-levels in a complex ion
  • electric field of ligands may cause d-orbitals in complex ions to split
  • so energy of an electronic transition between them corresponds to a photon of visible light
82
Q

Coloured compounds

A

Formation of coloured compounds in solution is a characteristic property of transition elements
- formation of coloured compounds is due to presence of partially-filled orbitals in 3d sub-level

83
Q

Electromagnetic spectrum

A
  • made up of a range of wavelengths/frequencies of electromagnetic radiation
  • very little is ‘seen’ by the human eye
  • small portion of spectrum that we can see, visible spectrum; 400-700nm
84
Q

White light in solution

A
  • white light passes through a solution containing transition metal ions
  • certain wavelengths of visible spectrum are absorbed, certain wavelengths are transmitted
85
Q

Colour wheel and complementary colours

A
  • use a colour wheel to explain relationship between absorbed and transmitted light
  • each colour in wheel has a complementary colour
  • for coloured solutions, light that we observe is the transmitted light- this is the complementary colour of the light that is absorbed

NB/ colour of complex ion seen is the complementary colour to that light which is absorbed

86
Q

Exceptions to forming coloured compounds

A
  • Cu+ and Sc3+ don’t form coloured compounds

- in these ions, d sub-level is either completely full or completely empty

87
Q

Splitting d sub-level

A
  • 5 d-orbitals in an isolated transition metal atom/ion are degenerate
  • ligands bond to central metal ion
  • repulsion between electrons in ligands and those in d orbitals in transition metals causes 5 d orbitals to split into 2 sets of different energy (non-degenerate orbitals)
88
Q

Delta E (change in E)

A

Difference in energy between two sets of d orbitals

  • exact energy difference between non-degenerate d orbitals in a transition metal is determined by several factors
  • one being the identity of ligands that surround the transition metal ion
89
Q

Transition metals and d orbitals

A

Transition metal ions have at least one partially filled d orbital

  • hence, electrons can transition from lower set to higher set of d orbitals by absorbing energy
  • energy absorbed corresponds to wavelength of visible light, w/ complementary colour of colour that is absorbed being transmitted
  • amount of energy absorbed during the electron transition (delta E) is calculated using E=hv
90
Q

Why are compounds of transition elements coloured?

A

Colour doesn’t occur due to transitions between principal energy levels but because some colour is absorbed when electrons are excited between split d-orbitals
- complexes of d-block elements are coloured because light is absorbed when an electron is excited between the d-orbitals

91
Q

Spectrochemical series

A

Effect of different ligands on degree of d orbital splitting in an octahedral complex, hence size of the change in energy is given in this series

92
Q

Presence of ammonia ligands

A
  • causes a difference in colour, due to presence of ammonia ligands, cause greater splitting of d orbitals
  • means that size of deltaE increases, resulting in a greater amount of energy being absorbed when electrons transition to higher set of d orbitals
  • hence, a shorter wavelength of light is absorbed, and solution appears a darker blue colour
93
Q

Factors that determine colour of complex ions

A
  • identity of metal ion at centre of complex ion
    • this affects no. of electrons involved in d orbitals, oxidation state of metal and nuclear charge at centre of ion
  • oxidation states of the metal ion
    • higher the oxidation state, higher the charge, lower the no. of e-
    • higher the electron repulsion between the ligand and d electrons
    • hence, higher the energy
  • nuclear charge
    • higher the no. of protons, higher the nuclear charge
    • this increases the electrostatic attraction between donated pairs of e- and nucleus
    • thus, change in energy is greater, resulting in a higher wavelength of light being emitted
  • ligand identity
    • higher the ligand on the spectrochemical series, higher the charge density
    • higher the charge density, higher the split in d orbitals, due tot increased repulsion within orbitals
    • higher the split, higher the energy
94
Q

Crystal field theory

A

Used to explain colour in transition metal complexes

  • theory accounts for colour of complex ions as arising from d-d electron transitions caused by splitting of d sub-level orbitals by repulsion effect of ligands present
  • type of splitting that takes place depends on no. of ligands attached as they interact w/ electron distributions in space of 5 d-orbitals
  • theory is based on assumption that ligands are point charges
95
Q

Anion ligands

A

Should have greatest splitting effect

- but anion ligands are found at low end of spectrochemical series, causing the least splitting

96
Q

Weakness of crystal field theory

A
  • doesn’t take account of covalent bonding between ligand and central metal cation
  • these are explained by alternative ligand field theory, model based on molecular orbital (MO) theory
97
Q

Transition metal ions in solution

A
  • have a high charge density

- thus, act as Lewis acids and attract species rich in e- (ligands)

98
Q

Ligands

A

neutral molecules or anions that contain one or more non-bonding pairs of electrons

  • ligands form covalent bonds w/ central transition metal ion to form a complex ion
  • if a ligand is higher on the spectrochemical series than another, it can displace that ligand
99
Q

Complex ion

A

has a central metal ion at its centre w/ a no. of other molecules surrounding it

100
Q

Polydentate ligands

A

ligands that can utilise 2 or more pairs of lone electrons to form a coordinate bond

IB Chemistry HL (9 decks)

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