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IB chemistry > topic 4/14- bonding and structure > Flashcards

topic 4/14- bonding and structure Flashcards

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1
Q

describe how positively charged ions form

A

metals lose valence electrons to form cations

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2
Q

describe how negatively charged ions form

A

non metals gain valence electrons to form anion

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3
Q

define an ionic bond

A

the electrostatic attraction between oppositely charged ions

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4
Q

state the formula and relative charge of ammonium

A

NH₄⁺

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5
Q

state the formula and relative charge of hydroxide

A

OH¯

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6
Q

state the formula and relative charge of nitrate

A

NO3-

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7
Q

state the formula and relative charge of hydrogencarbonate

A

HCO3-

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8
Q

state the formula and relative charge of carbonate

A

CO₃²⁻

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9
Q

state the formula and relative charge of sulfate

A

SO₄²-

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10
Q

state the formula and relative charge of phosphate

A

PO₄³⁻

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11
Q

describe ionic compounds under normal conditions

A

solids with lattice structures

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12
Q

describe the physical properties of ionic compounds:
- melting/boiling point
- volatility
- electrical conductivity
- solubility

A
  • high melting/boiling points
  • low volatility
  • do not conduct when solid, but conduct when molten or aqueous
  • soluble
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13
Q

explain the high melting/boiling points of ionic compounds

A
  • strong EFOA between oppositely charged ions
  • lot of energy required to overcome them
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14
Q

does magnesium oxide or sodium chloride have a higher melting point? why?

A

magnesium oxide; the EFOA between the 2+ and 2- ions in MgO are much stronger than those between the 1+ and 1- ions in NaCl

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15
Q

explain the low volatility of ionic compounds

A

the strong EFOA between oppositely charged ions take a lot of energy to overcome

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16
Q

explain the electrical conductivity of ionic compounds

A
  • in a solid state, the ions are in fixed positions (ionic lattice) and cannot move.
  • when aqueous or molten, the ions are able to move as the lattice is broken down
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17
Q

describe what happens to the ions in an ionic compound when it dissolves

A
  1. bonds/attraction between the solid particles are broken
  2. new bonds form between sps and water molecules
  3. water molecules surround sps
  4. due to sliding movement of water molecules, sps move through liquid until evenly distributed
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18
Q

why are some substances soluble and others insoluble?

A

there is an energy cost and an energy gain which determines whether a process occurs without an input of energy

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19
Q

a substance will be soluble if the energy gain is —– than the energy cost.

A

greater than

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20
Q

state the two sources of energy cost in dissolving

A
  1. breaking hydrogen bonds in water/solvent
  2. breaking bonds between solute particles
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21
Q

state the two sources of energy gain in dissolving

A
  1. formation of bonds between solute and solvent
  2. increasing entropy
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22
Q

explain the solubility of ionic compounds in polar solvents like water

A

energy released when ions surrounded by H2O molecules (ion-dipole interactions) > energy required to break ionic lattice

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23
Q

explain the insolubility of ionic compounds in non polar solvents

A

energy required to break apart ionic lattice > energy released when non-polar solvent forms interactions with the ions (London forces)

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24
Q

define a covalent bond

A

the electrostatic attraction between a shared pair of electrons and two positively charged nuclei

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25
Q

describe a single covalent bond

A

one shared pair of electrons

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26
Q

describe a double covalent bond

A

two shared pairs of electrons

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27
Q

describe a triple covalent bond

A

three shared pairs of electrons

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28
Q

describe a coordinate covalent bond

A

a covalent bond where one atom has donated two electrons to the shared pair

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29
Q

describe the relationship between bond length and bond strength

A

as bond length increases, bond strength decreases

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30
Q

describe the relationship between number of shared pairs and bond strength/length

A

as number of shared pairs increases, strength of bond increases but length decreases.

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31
Q

explain what happens to the bond strength as you go from HF to HI.

A

bond strength decreases;
- distance of bond pair from nucleus decreases, shielding of bond pair increases
- this outweighs an increase in atomic number

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32
Q

define electronegativity

A

a measure of the tendency of an atom to attract a bonding pair of electrons

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33
Q

what is impure ionic bonding?

A

when negative ions are polarised by positive ions.

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34
Q

why does impure ionic bonding occur?

A

the positive ion attracts the electrons of the negative ion

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35
Q

explain why smaller, more highly charged positive ions have a greater polarising effect

A

smaller, more highly charged positive ions have a greater charge density, causing them to have a greater attraction to the anion

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36
Q

explain why larger negative ions are more easily polarised than smaller negative ions

A

larger negative ions have a greater distance and shielding, meaning the electrons are less attracted to the nucleus of the anion

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37
Q

give the formula for the most impure ionic bonding

A

small, highly charged positive ion + large negative ion

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38
Q

explain why electronegativity increases ACROSS a period

A

no. of protons is increasing (nucleus more positive) but distance and shielding are staying mostly the same

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39
Q

explain why electronegativity decreases DOWN a group

A

there is more shielding and an increased distance (so less attraction)

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40
Q

which element has the highest electronegativity?

A

fluorine

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41
Q

what is a pure covalent bond?

A

when the electronegativities of the two atoms are the same

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42
Q

explain the difference between covalent and ionic bonding in terms of electronegativities.

A
  • in a covalent bond, the difference in electronegativity is zero or very small.
  • in an ionic bond the elements have a large difference in electronegativity
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43
Q

describe the relationship between difference in electronegativity and type of bond

A
  • <0.4= covalent bond
  • 0.4-1.8= polar covalent
  • > 1.8= ionic bond
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44
Q

the more electronegative atom has a slightly —– charge
the less electronegative atom has a slightly —– charge

A
  1. negative
  2. positive
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45
Q

a molecule is polar if….

A

there is a net dipole

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46
Q

what does the ‘octet rule’ refer to?

A

the tendency of atoms to gain a valence shell with a total of 8 electrons

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47
Q

state two common exceptions two the octet rule

A

Beryllium and boron; these might form stable compounds with incomplete octets of electrons

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48
Q

what are lewis structures?

A

diagrams that show all the valence electrons in a covalently bonded species

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49
Q

when do resonance structures occur?

A

when there is more than one possible position for a double bond in a molecule

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50
Q

draw the resonance structure for C6H6

A

refer to flashcard

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51
Q

draw the resonance structure for CO₃²⁻

A

refer to flashcard

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52
Q

what is a resonance structure?

A

one of two or more alternative Lewis structures for a molecule or ion that cannot be described fully with one Lewis structure alone

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53
Q

what are the shapes of species determined by?

A

the repulsion of electron pairs according to VSEPR (valence shell electron pair repulsion) theory

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54
Q

describe VSEPR theory

A

pairs of electrons repel each other so that they are as far apart as possible, keeping the force of repulsion to a minimum

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55
Q

how many types of covalent shapes are there that ONLY include bond pairs (no lone pairs)?

A

5

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56
Q

name the 5 covalent molecule shapes with only bond pairs

A
  • linear molecules
  • trigonal planar molecules
  • tetrahedral molecules
  • trigonal bipyramidal molecules
  • octahedral molecules
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57
Q

describe linear molecules (BP only)
- no of bond pairs
- bond angle
- 3 examples
- shape

A
  • 2 bond pairs
  • 180
  • BeCl2, CO2 and all diatomic molecules
  • straight line
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58
Q

describe trigonal planar molecules
- no of bond pairs
- bond angle
- 1 example
- shape

A
  • 3 bond pairs
  • 120
  • BCl3
  • flat peace-sign
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59
Q

describe tetrahedral molecules
- no of bond pairs
- bond angles
- 1 examples
- shape

A
  • 4 bond pairs
  • 109.5
  • methane, CH4
  • Eiffel Tower
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60
Q

describe trigonal bipyramidal molecules
- no of bond pairs
- bond angles
- 1 example
- shape

A
  • 5 bond pairs
  • 90 and 120
  • PF5
  • fidget spinner shot by an arrow
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61
Q

describe octahedral molecules
- no of bond pairs
- bond angles
- 1 example
- shape

A
  • 6 bond pairs
  • 90
  • SF6
  • christian cross that has been shot by an arrow
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62
Q

why does a lone pair have a greater repulsive effect than a bonding pair?

A
  • lone pair electrons are localised to an atom, so they are closer to each other
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63
Q

state the hierarchy of repulsion

A

lp to lp> lp to bp> bp to bp

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64
Q

name the 7 types of lone pair inclusive molecule shapes

A
  • trigonal pyramidal molecules
  • v shaped/bent molecules
  • square planar molecules
  • square pyramidal molecules
  • seesaw molecules
  • t/arrow shaped molecules
  • linear molecules
    (two very silly Swiss singers terrify Lav)
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65
Q

describe trigonal pyramidal molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • like tetrahedral molecules but without top
  • 3 bp, 1 lp, 4 ed
  • 107
  • NH3
  • beheaded Eiffel tower
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66
Q

describe v shaped/bent molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example

A

type 1 (SO2 DB):
- like trigonal planar but without top
- 2 bp, 1 lp, 3ed
- 104.5
type 2 (H2O):
- like type 1 but extra pair
- 2 bp, 2 lp, 4 ed
- 104.5

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67
Q

describe square planar molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • like octahedral but without both vertical bits
  • 4 bp, 2 lp, 6 ed
  • 90
  • XeF4
  • cross laid on its side + 2lp above and below
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68
Q

describe square pyramidal
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • like octahedral but without bottom vertical bit
  • 5 bp, 1 lp, 6 ed
  • 85-87.5
  • BrF5
  • cross laid on its side + stick up + lp below
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69
Q

describe see saw molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example

A
  • like trigonal bipyramidal but without left side bit
  • 4 bp, 1 lp, 5 ed
  • 87.5-90 (equatorial-axial), 117 (e-e)
  • SF4
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70
Q

describe t/arrow shaped molecules
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example

A
  • like trigonal bipyramidal but only side 3 + 2 lp
  • 3 bp, 2 lp, 5 ed
  • 87.5-90
  • XeOF2
71
Q

describe linear molecules (WITH lone pairs)
- which bpo shape are they like
- no of bp, lp, ed
- bond angles
- example
- shape

A
  • like trigonal bipyramidal but only 2 vertical bits
  • 2 bp, 3 lp, 5 ed
  • 180
  • I3-
  • straight line
72
Q

define metallic bonding

A

the electrostatic attraction between a lattice of positive ions and delocalised electrons

73
Q

why do metals conduct electricity?

A

the have free-flowing, delocalised electrons which are free to move.

74
Q

why are metals malleable?

A

they have layers which can slide across each other without disrupting the bonding in place.

75
Q

what would be the ratio of ions to delocalised electrons in lithium? (Li is in group 1)

A

1:1 (each Li atom loses ONE electron to become an ion)

76
Q

what would be the ratio of ions to delocalised electrons in magnesium? (Mg is in group 2)

A

1:2 (each Mg atom loses TWO electrons to become an ion)

77
Q

explain why melting and boiling points increase from sodium to magnesium to aluminium.

A
  • slightly smaller radius
  • greater charge
  • more delocalised electrons
    => greater charge density
    => greater ESFOA so more energy required
78
Q

explain why the melting and boiling points of group 1 metal decrease going down the group

A
  • larger radius
  • more shielding
    => lower charge density
    => less ESFOA so less energy required
79
Q

as charge of ion increases, strength of metallic bond —-

A

increases

80
Q

as radius of ion increases, strength of metallic bond —

A

decreases

81
Q

what gives alloys their properties?

A

the fact that they are mixtures of more than one metal.

82
Q

gives two examples of alloys

A

brass (copper + zinc)
bronze (copper + tin)

83
Q

explain the properties of alloys

A
  • contains ions of different sizes
  • this disrupts the structure/regular lattice arrangement
  • the layers slide over each other less easily
    => harder/less malleable
84
Q

what kind of bonding can carbon and silicon form?

A

giant covalent/network covalent structures

85
Q

Graphite, diamond, buckminsterfullerene and graphene are ? of carbon

A

allotropes

86
Q

describe the structure of graphite

A
  • hexagonal layers of carbon
  • each atom bonded to 3 others
  • London forces between layers
  • 120
87
Q

describe the structure of diamond

A
  • continuous tetrahedral arrangement of C atoms
  • each atom bonded to 4 others
  • 109.5
88
Q

describe the structure of buckminsterfullerene

A
  • 60 carbon atoms
  • each atom bonded to 3 others (4th ed)
  • interlocking hexagons and pentagons
89
Q

describe structure of graphene

A
  • hexagonal sheet
  • each atom bonded to 3 others
  • 1 layer thick
90
Q

describe structure of silica, or silicon dioxide

A
  • continuous tetrahedral structure
  • each Si shared between 4 O
  • each O shared between 2 Si
91
Q

explain the properties of diamond

A
  • high melting/boiling point: strong covalent bonds
  • non conductor: no de
  • hard: each atom forms 4 bonds
92
Q

explain the properties of graphite

A
  • high melting/boiling point: strong covalent bonds within layers
  • conductor: each atom forms 3 bonds so there is 1 de
  • brittle: weak London forces between layers
93
Q

explain the properties of graphene

A
  • high melting/boiling point: strong covalent bonds
  • conductor: each atom forms 3 bonds so there is 1 de
94
Q

explain the properties of buckminsterfullerene

A
  • low melting/boiling point: weak intermolecular forces between molecules
  • semi-conductor: each carbon forms 3 bonds so there is 1de
95
Q

explain the properties of silica, or silicon dioxide

A
  • high melting/boiling point: strong covalent bonds
  • non-conductor: no de
96
Q

what do covalent bonds arise from?

A

the overlap of atomic orbitals

97
Q

describe sigma bonds

A
  • formed by any head-on/axial overlap of orbitals
  • s-s/s-p/p-p
98
Q

describe pi bonds

A
  • formed by the lateral overlap of two p orbitals
99
Q

5 differences between sigma and pi bonds

A
  • only 1 s bond exists between two atoms; can be more than 1 pi bond between two atoms
  • e density is maximum about the bond axis; e density high above and below bond axis
  • free rotation about s bond possible; not possible
  • s can be independently formed; formed after sigma bond
  • relatively strong; weak
100
Q

define formal charge

A

the charge an atom would have if all the atoms in the molecule had the same electronegativity

101
Q

state the equation for formal charge

A

FC= no of valence electrons - 0.5(no of bonding electrons) - no of non bonding electrons

102
Q

what Lewis structure is preferred?

A

the one with the atoms having FC (formal charge) values closest to 0

103
Q

When does delocalisation of electrons in a molecule arise?

A

when electrons are shared between all atoms in a molecule or ion

104
Q

what wavelength of light can break down ozone?

A

UV-B light, 280-320nm

105
Q

give the equation for the splitting of ozone

A

O3–> O2 + O’

106
Q

which type of UV has the highest energy and shortest wavelength?

A

UV-C

107
Q

why is a specific wavelength of light required to dissociate oxygen and ozone?

A

as a certain amount of energy is needed in order for the bonds to be broken

108
Q

how does the ozone layer protect the earth from UV radiation?

A

most of UV-C and UV-B are absorbed by oxygen and ozone molecules. therefore, only UV-A reaches the earth’s surface, which is less harmful and lower energy.

109
Q

oxygen (O2) absorbs

A

UV-C

110
Q

ozone (O3) absorbs

A

UV-B

111
Q

state the equation for the dissociation of oxygen

A

O2–> 2O’

112
Q

state the mechanisms by which O3 is broken down by CCl3F

A
  1. CCl3F -(UV)-> ‘CCl2F + Cl’ (UV causes C-Cl bond to break)
  2. Cl’ + O3 –> ClO’ +O2 (Cl radical attacks O3)
  3. ClO’ + O’ –> O2 + Cl’
113
Q

state the mechanisms by which O3 is broken down by NOx

A

formation of nitrogen oxide;
- N2O + O’ –> 2NO
- NO2 + O’ –> NO + O2
breakdown of O3:
- ‘No + O3 –> NO2 + O2
- NO2 + O3 –> 2O2 + NO

114
Q

London forces occur between….

A

all molecules and atoms of noble gases

115
Q

Dipole-dipole interactions occur between…

A

polar molecules

116
Q

hydrogen bonding occurs when…

A

a hydrogen atom is bonded to Nitrogen, Fluorine or Oxygen atoms with a lone pair of electrons

117
Q

Explain how London (dispersion forces) arise

A

London forces are caused by temporary dipoles which arise in atoms due to uneven distribution of electrons

118
Q

Explain the trend in boiling points of Group 7 elements

A
  • bpt increases as you go down the group
  • more electrons in atoms; bigger dipole
  • London forces stronger
  • take more energy to overcome
119
Q

As chain length of alkanes increases, boiling point ?
Why?

A

increases- more points of contact with each adjacent molecule, increasing London forces

120
Q

Explain how permanent dipole-dipole interactions arise

A

molecules with permanent dipole:
- e- spend more time with atom that has higher electronegativity, giving it a negative dipole which attracts the positive dipole on the nearby molecule

121
Q

explain how hydrogen bonding arises

A
  1. the polarised H atom has its only electrons withdrawn, exposing the proton in it’s nucleus
  2. this high positive charge density attracts the lone pairs of electrons on F/N/O atoms of neighbouring molecules.
122
Q

what is the rule with drawing hydrogen bonding?

A

the covalent bond of the hydrogen’s molecule and the hydrogen bond must be parallel.

123
Q

state the relative strengths of the 3 intermolecular forces

A

London (dispersion) forces< dipole-dipole forces < hydrogen bonds

124
Q

why would ethanol (CH3CH2OH) have a higher melting point than methoxymethane (CH3OCH3)?

A

ethanol has hydrogen bonding between -O-H bonds, methoxymethane only has London forces.

125
Q

why would pentan-1-ol have a higher boiling point than pentan-3-ol?

A
  • hydrogen bonds in pentan-3-ol are weaker as longer chains have more points of contact
126
Q

when does a hybrid orbital occur?

A

a hybrid orbital results from the mixing of different types of atomic orbitals on the same atom

127
Q

define hybridisation

A

intermixing orbitals of slightly different energies so as to redistribute their energies and form new sets of orbitals of equivalent energy and shape.

128
Q

the hybridized orbitals are always equivalent in ——and ——

A

energy and shape

129
Q

the number of hybridized orbitals formed is equal to —————

A

the number of orbitals that get hybridized

130
Q

state the 3 types of hybridisation

A

sp, sp2, sp3

131
Q

describe the formation of sp3 hybrid orbitals in methane

A
  • four valence orbitals of the carbon (one 2s and three 2p orbitals) combine
  • form four equivalent hybrid orbitals
  • these overlap with the 1s orbitals of hydrogen to form bonds
132
Q

explain why sp3 orbitals result in a tetrahedral arrangement

A
  • each sp3-hybridized orbital bears an electron, and electrons repel each other
  • to minimize the repulsion between electrons, the four sp3-hybridized orbitals arrange themselves so that they are as far away as possible from each other
  • resulting in the tetrahedral arrangement predicted by VSPER (109.5’)
133
Q

describe the formation of sp2 hybrid orbitals in ethene

A
  • three valence orbitals of the carbon (one 2s and two 2p orbitals) combine
  • form three equivalent hybrid orbitals
  • the remaining electron goes into the unhybridized 2p orbital (different energy=2nd bond)
134
Q

describe the formation of sp hybrid orbitals in ethyne

A
  • two valence orbitals of the carbon (one 2s and one 2p) combine
  • form two equivalent hybrid orbitals
  • remaining 2 electrons go into the unhybridized 2p orbitals (different energy=2nd and 3rd bond)
135
Q

explain why sp2 orbitals result in a trigonal planar arrangement

A
  • each sp2-hybridized orbital bears an electron, and electrons repel each other
  • to minimize the repulsion between electrons, the three sp2-hybridized orbitals arrange themselves so that they are as far away as possible from each other
  • resulting in the trigonal planar arrangement predicted by VSPER (120’)
136
Q

explain why sp orbitals result in a linear arrangement

A
  • each sp-hybridized orbital bears an electron, and electrons repel each other
  • to minimize the repulsion between electrons, the three sp-hybridized orbitals arrange themselves so that they are as far away as possible from each other
  • resulting in the linear arrangement predicted by VSPER (180’)
137
Q

draw ammonium and ammonia

A

refer elsewhere

138
Q

draw SO4 2- with lines

A

refer elsewhere

139
Q

draw dot and cross for OH-

A

refer elsewhere

140
Q

draw lewis and line for NO3-

A

refer elsewhere

141
Q

draw CO3 2- with lines

A

refer elsewhere

142
Q

draw PO4 3- with lines

A

refer elsewhere

143
Q

draw HCO3- with lines

A

refer elsewhere

144
Q
A
145
Q

macroscopic level of gases

A

properties

146
Q

molecular level of gases

A
  • structure
  • dynamics (motion)
  • intermolecular forces
147
Q

equation and definition for pressure

A

P=F/A
= force exerted per area

148
Q

SI units for pressure

A

Pa = N/m^2
kPa=10^3 Pa

149
Q

1atm =

A

101.3 kPa, 760mmHg (Torr)

150
Q

1 bar =

A

100.0 kPa

151
Q

equation for pressure with density, g, and h

A

p x h x g, where g = 9.8m/s^2

152
Q

what does the manometer lead to?

A

an equalisation of pressure:
P(atm) = P(Hg) = p(Hg)gh(Hg)

153
Q

if the Hg in a manometer was replaced with water, would the height of the water be greater or less than Hg?

A

greater, due to water’s lower density (P constant, p goes down, g constant, so h goes up)

154
Q

describe kinetic molecular theory - the ideal gas

A
  1. gases made of tiny particles moving completely randomly
  2. total volume of particles very small compared to size of container
  3. particles do not interact with each other
  4. particle collisions are elastic (no energy lost)
  5. kinetic energy (KE) increases with temperature
155
Q

describe (5) KE increases with temperature

A

for a large collection of molecules:
- at a given temperature, all gases have the same distribution of kinetic energy
- each molecule: KE = 1/2mv^2

156
Q

what is the effect of temperature on average kinetic energy?

A

the average kinetic energy increases with temperature

157
Q

state 2 equations relating kinetic energy with temperature

A

KE(avg) = 3RT/2
Temperature is in Kelvin
KE is in J/mol

KE (avg) = 3RT/2Na
for a single gas molecule

158
Q

units for R

A

J/molK

159
Q

3RT/2Na =

A

1/2mu^2

160
Q

state the 2 equations relating u, R, T, m, Na, and M

A
161
Q

how is kinetic energy dependent on the mass of the particles?

A

at the same temperature, more massive molecules move slower

162
Q

what does pressure result from?

A

gas particles colliding with container walls

163
Q

Boyle’s law

A

at constant T and fixed n, volume is inversely proportional to pressure, or V=k/P

164
Q

low Pext = Pgas

A

-> high V

165
Q

high Pext = Pgas

A

-> low V

166
Q

Charles Law

A

at same P and fixed n, volume is proportional to temperature, or V= kT

167
Q

Avogadro’s law

A

V=kn

168
Q

what equation describes an ideal gas?

A

PV = nRT
P - atm
V - L
T - K

169
Q

P1V1/n1T1 =

A

P2V2/n2T2

170
Q

mole fraction of gas A

A

Xa = na/ntotal = Pa/Ptotal

171
Q

law of partial pressures

A

Pa + Pb + Pc +…. = PTotal

172
Q

define effusion

A

escape of a pas through a hole into a vacuum

173
Q

define diffusion

A

movement of one gas through another

IB chemistry (29 decks)

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