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Chemistry AS Level CIE > Chapter 4 Chemical bonding > Flashcards

Chapter 4 Chemical bonding Flashcards

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1
Q

Electronegativity

A

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

  • This phenomenon arises from the positive nucleus’s ability to attract the negatively charged electrons, in the outer shells, towards itself
  • The Pauling scale is used to assign a value of electronegativity for each atom
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2
Q

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

A

fluorine

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

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3
Q

Electronegativity: Affecting Factors

A
  • Nuclear charge
  • Atomic radius
  • Shielding
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4
Q

Electronegativity: Affecting Factors: Nuclear charge

A
  • Attraction exists between the positively charged protons in the nucleus and negatively charged electrons found in the energy levels of an atom
  • An increase in the number of protons leads to an increase in nuclear attraction for the electrons in the outer shells
  • Therefore, an increased nuclear charge results in an increased electronegativity
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5
Q

Electronegativity: Affecting Factors: Atomic radius

A
  • The atomic radius is the distance between the nucleus and electrons in the outermost shell
  • 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
  • Therefore, an increased atomic radius results in a decreased electronegativity
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6
Q

Electronegativity: Affecting Factors: Shielding

A
  • Filled energy levels can shield (mask) the effect of the nuclear charge causing the outer electrons to be less attracted to the nucleus
  • Therefore, the addition of extra shells and subshells in an atom will cause the outer electrons to experience less of the attractive force of the nucleus
  • —-Sodium (Period 3, Group 1) has higher electronegativity than caesium (Period 6, Group 1) as it has fewer shells and therefore the outer electrons experience less shielding than in caesium

-Thus, an increased number of inner shells and subshells will result in a decreased electronegativity

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7
Q

As nuclear charge increases

A

the nucleus has a greater attractive force on the electrons in shells given that the shielding doesn’t increase.

As a result of this, the atomic radius decreases.

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8
Q

Electronegativity: Trends: down a group

A
  • There is a decrease in electronegativity going down the Group
  • 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

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9
Q

Electronegativity: Trends: Across a period

A
  • Electronegativity increases across a Period
  • The nuclear charge increases with the addition of protons to the nucleus
  • Shielding remains reasonably 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 of the Periodic Table

-This results in smaller atomic radii

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10
Q

Electronegativity & covalent bonds

A

-Single covalent bonds are formed by sharing a pair of electrons between two atoms
—-In diatomic molecules the electron density is shared equally between the two atoms
Eg. H2, O2 and Cl2
-Both atoms will have the same electronegativity value and have an equal attraction for the bonding pair of electrons leading to formation of a covalent bond
-The equal distribution leads to a non-polar molecule

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11
Q

Electronegativity & ionic bonds

A

-When atoms of different electronegativities form a molecule, the shared electrons are not equally distributed in the bond
-The more electronegative atom (the atom with the higher value on the Pauling scale) will draw the bonding pair of electrons towards itself
-A molecule with partial charges forms as a result
The more electronegative atom will have a partial negative charge (delta negative, δ–)
-The less electronegative atom will have a partial positive charge (delta positive, δ+)
-This leads to a polar covalent molecule

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12
Q

If there is a large difference in electronegativity of the two atoms in a molecule

A

the least electronegative atom’s electron will transfer to the other atom

  • This in turn leads to an ionic bond – one atom transfers its electron and the other gains that electron
  • —-The cation is a positively charged species which has lost (an) electron(s)
  • —-The anion is a negatively charged species which has gained (an) electron(s)
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13
Q

Ionic Bonding: Definition

A

-As a general rule, metals are on the left of the Periodic Table and nonmetals are on the right-hand side

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14
Q

ionic bonds

A

-Ionic bonds involve the transfer of electrons from a metallic element to a non-metallic element
Transferring electrons usually leaves the metal and the non-metal with a full outer shell
Metals lose electrons from their valence shell forming positively charged cations
Non-metal atoms gain electrons forming negatively charged anions
Once the atoms become ions, their electronic configurations are the same as a stable noble gas.

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15
Q

Electrostatic attractions

A

are formed between the oppositely charged ions to form ionic compounds

  • This form of attraction is very strong and requires a lot of energy to overcome
  • —–This causes high melting points in ionic compounds
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16
Q

Sodium chloride

A
  • Sodium is a Group 1 metal
  • It loses its outer electron to form a sodium ion with a +1 charge (Na+)
  • Chlorine is a Group 7 non-metal
  • It gains 1 electron to form a chloride ion with a -1 charge (Cl–)
  • The oppositely charged ions are attracted to each other by electrostatic forces to form NaCl
  • The final ionic solid is neutral in charge
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17
Q

Magnesium oxide

A
  • Magnesium is a Group 2 metal
  • It loses its 2 outer electrons to form a magnesium ion with a +2 charge (Mg2+)
  • Oxygen is a Group 6 non-metal
  • It gains 2 electrons to form an oxide ion with a -2 charge (O2-)
  • The oppositely charged ions are attracted to each other to by electrostatic forces to form MgO
  • The final ionic solid is neutral in charge
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18
Q

Calcium fluoride

A
  • Calcium is a Group 2 metal
  • It loses its 2 outer electrons to form a calcium ion with a +2 charge (Ca2+)
  • Fluorine is a Group 7 non-metal
  • It gains 1 electron to form a fluoride ion with a -1 charge (F–)
  • As before, the positive and negative ions are attracted to each other
  • To cancel out the 2+ charge of the calcium ion, 2 fluoride ions are needed
  • Calcium fluoride is made when 1 calcium ion and 2 fluoride ions form an ionic bond, CaF2
  • The final ionic solid of CaF2 is neutral in charge
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19
Q

Metallic Bonding: Definition

A

Metal atoms are tightly packed together in lattice structures
When the metal atoms are in lattice structures, the electrons in their outer shells are free to move through the structure
The free-moving electrons are called ‘delocalised electrons’ and they are not bound to their atom

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20
Q

metallic bonding

A
  • When the electrons are delocalised, the metal atoms become positively charged
  • The positive charges repel each other and keep the neatly arranged lattice in place
  • There are very strong electrostatic forces between the positive metal centres and the ‘sea’ of delocalised electrons
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21
Q

The strength of electrostatic attraction can be increased by

A
  • Increasing the number of delocalised electrons per metal atom
  • Increasing the positive charges on the metal centres in the lattice
  • Decreasing the size of the metal ions
  • Due to the delocalised ‘sea’ of electrons, metallic structures have some characteristic properties
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22
Q

features of metallic bonds

A
  • high melting and boiling points

- electrical conductivity

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23
Q

high melting and boiling points why

A
  • positive metal atom centres and the delocalized electrons in a metallic lattice have strong electrostatic forces between them
  • Therefore high energy is needed to overcome the strong forces of attraction
  • As the number of mobile charges increase from left to right of he Periodic Table, the melting and boiling points increase as the electrostatic forces become stronger
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24
Q

electrical conductivity why

A

-When a potential difference is applied to the metallic lattice, the delocalised electrons repel away from the negative terminal and move towards the positive terminal

  • As the numbers of valance electrons increases across the Period, the number of delocalised charges also increases
  • —sodium= 1 valance electron
  • —Magnesium= 2 valance electrons
  • —Aluminum= 3 valances electrons

-Therefore the ability to conduct electricity also increases

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25
Q

Covalent Bonding: Definition

A
  • Covalent bonding occurs between two nonmetals
  • A covalent bond involves the electrostatic attraction between nuclei of two atoms and the bonding electrons of their outer shells
  • No electrons are transferred but only shared in this type of bonding
  • Non-metals are able to share pairs of electrons to form different types of covalent bonds
  • Sharing electrons in the covalent bond allows each of the 2 atoms to achieve an electron configuration similar to a noble gas
  • —–This makes each atom more stable
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26
Q

Dot & cross diagrams

A
  • Dot and cross diagrams are used to represent covalent bonding
  • They show just the outer shell of the atoms involved
  • To differentiate between the two atoms involved, dots for electrons of one atom and crosses for electrons of the other atom are used
  • Electrons are shown in pairs on dot-and-cross diagrams
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27
Q

‘expanding the octet rule’

A

Being able to accommodate more than 8 electrons in the outer shell

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28
Q

‘electron deficient’

A

Accommodating less than 8 electrons in the outer shell means than the central atom

-Some examples of this occurring can be seen with Period 3 elements

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29
Q

Dative Bonding: Definition

A
  • In simple covalent bonds the two atoms involved shares electrons
  • Some molecules have a lone pair of electrons that can be donated to form a bond with an electron-deficient atom
  • —An electron-deficient atom is an atom that has an unfilled outer orbital
  • So both electrons are from the same atom
  • This type of bonding is called dative covalent bonding or coordinate bond
  • An example of a dative bond is in an ammonium ion
  • —The hydrogen ion, H+ is electron-deficient and has space for two electrons in its shell
  • —The nitrogen atom in ammonia has a lone pair of electrons which it can donate to the hydrogen ion to form a dative covalent bond
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30
Q

Aluminium chloride is also formed using

A

dative covalent bonding

  • At high temperatures aluminium chloride can exist as a monomer (AlCl3)
  • —The molecule is electron-deficient and needs to electrons complete the aluminium atom’s outer shell
  • At lower temperatures the two molecules of AlCl3 join together to form a dimer (Al2Cl6)
  • —-The molecules combine because lone pairs of electrons on two of the chlorine atoms form two coordinate bonds with the aluminium atoms
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31
Q

Bond overlap in covalent bonds

A

-A single covalent bond is formed when two nonmetals combine
-Each atom that combines has an atomic orbital containing a single unpaired electron
-When a covalent bond is formed, the atomic orbitals overlap to form a combined orbital containing two electrons
—–This new orbital is called the molecular orbital
The greater the atomic orbital overlap, the stronger the bond

  • Sigma (σ) bonds are formed by direct overlap of orbitals between the bonding atoms
  • Pi (π) bonds are formed by the sideways overlap of adjacent above and below the σ bond
32
Q

σ bonds

A
  • Sigma (σ) bonds are formed from the end-on overlap of atomic orbitals
  • S orbitals overlap this way as well as p orbitals
  • The electron density in a σ bond is symmetrical about a line joining the nuclei of the atoms forming the bond
  • The pair of electrons is found between the nuclei of the two atoms
  • The electrostatic attraction between the electrons and nuclei bonds the atoms to each other
  • single bonds
33
Q

π bonds

A
  • Pi (π) bonds are formed from the sideways overlap of adjacent p orbitals
  • The two lobes that make up the π bond lie above and below the plane of the σ bond
  • This maximises overlap of the p orbitals
  • A single π bond is drawn as two electron clouds one arising from each lobe of the p orbitals
  • The two clouds of electrons in a π bond represent one bond containing two electrons by

Double bond between atoms

34
Q

Examples of sigma & pi bonds: Hydrogen

A
  • The hydrogen atom has only one s orbital

- The s orbitals of the two hydrogen atoms will overlap to form a σ bond

35
Q

Examples of sigma & pi bonds: Ethene

A
  • Each carbon atom uses three of its four electrons to form σ bonds
  • Two σ bonds are formed with the hydrogen atoms
  • One σ bond is formed with the other carbon atom
  • The fourth electron from each carbon atom occupies a p orbital which overlaps sideways with another p orbital on the other carbon atom to form a π bond
  • This means that the C-C is a double bond: one σ and one π bond
36
Q

Examples of sigma & pi bonds: Ethyne

A
  • This molecule contains a triple bond formed from two π bonds (at right angles to each other) and one σ bond
  • Each carbon atom uses two of its four electrons to form σ bonds
  • One σ bond is formed with the hydrogen atom
  • One σ bond is formed with the other carbon atom
  • Two electrons are used to form two π bonds with the other carbon atom
37
Q

Examples of sigma & pi bonds: Hydrogen cyanide

A
  • Hydrogen cyanide contains a triple bond
  • One σ bond is formed between the H and C atom (overlap of an sp C hybridised orbital with the 1s H orbital)
  • A second σ bond is formed between the C and N atom (overlap of an sp C hybridised orbital with a p orbital of N)
  • The remaining two sets of p orbitals of nitrogen and carbon will overlap to form two π bonds at right angles to each other
38
Q

Examples of sigma & pi bonds: Nitrogen

A
  • Nitrogen too contains a triple bond
  • The triple bond is formed from the overlap of the s orbitals on each N to form a σ bond and the overlap of two sets of p orbitals on the nitrogen atoms to form two π bonds
  • These π bonds are at right angles to each other
39
Q

Hybridisation

A
  • The p atomic orbitals can also overlap end-on to form σ bonds
  • In order for them to do this, they first need to become modified in order to gain s orbital character
  • The orbitals are therefore slightly changed in shape to make one of the p orbital lobe bigger
  • This mixing of atomic orbitals to form covalent bonds is called hybridisation
  • —Mixing an s with three p orbitals is called sp3 hybridisation (each orbital has ¼ s character and ¾ p character)
  • —Mixing an s with tw p orbital is called sp2 hybridisation
  • —And mixing an s with one p-type orbitals formssp hybridised orbitals
40
Q

Bond energy

A
  • The bond energy is the energy required to break one mole of a particular covalent bond in the gaseous states
  • —Bond energy has units of kJ mol-1
  • The larger the bond energy, the stronger the covalent bond is
41
Q

Reactivity of covalent molecules

A
  • The reactivity of a covalent bond is greatly influenced by:
  • —-The bond polarity
  • —-The bond strength
  • —-The bond type (σ/π)
42
Q

Bond length

A
  • The bond length is internuclear distance of two covalently bonded atoms
  • It is the distance from the nucleus of one atom to another atom which forms the covalent bond
  • The greater the forces of attraction between electrons and nuclei, the more the atoms are pulled closer to each other
  • This decreases the bond length of a molecule and increases the strength of the covalent bond
  • Triple bonds are the shortest and strongest covalent bonds due to the large electron density between the nuclei of the two atoms
  • This increase the forces of attraction between the electrons and nuclei of the atoms
  • As a result of this, the atoms are pulled closer together causing a shorter bond length
  • The increased forces of attraction also means that the covalent bond is stronger
43
Q

VSEPR: Theory

A
  • The valence shell electron pair repulsion theory (VSEPR) predicts the shape and bond angles of molecules
  • Electrons are negatively charged and will repel other electrons when close to each other
  • In a molecule, the bonding pair of electrons will repel other electrons around the central atom forcing the molecule to adopt a shape in which these repulsive forces are minimised
44
Q

When determining the shape and bond angles of a molecule, the following VSEPR rules should be considered:

A
  • Valence shell electrons are those electrons that are found in the outer shell
  • Electron pairs repel each other as they have similar charges
  • Lone pair electrons repel each other more than bonded pairs
  • Repulsion between multiple and single bonds is treated the same as for repulsion between single bonds
  • Repulsion between pairs of double bonds are greater
  • The most stable shape is adopted to minimize the repulsion forces
45
Q

Different types of electron pairs have different repulsive forces

A
  • Lone pairs of electrons have a more concentrated electron charge cloud than bonding pairs of electrons
  • The cloud charges are wider and closer to the central atom’s nucleus
  • The order of repulsion is therefore: lone pair – lone pair > lone pair – bond pair > bond pair – bond pair
46
Q

Hydrogen bonding

A

is the strongest form of intermolecular bonding

  • —Intermolecular bonds are bonds between molecules
  • —Hydrogen bonding is a type of permanent dipole – permanent dipole bonding
  • For hydrogen bonding to take place the following is needed:
  • —A species which has an O or N (very electronegative) atom with an available lone pair of electrons
  • —A species with an -OH or -NH group
  • When hydrogen is covalently bonded to an electronegative atom, such as O or N, the bond becomes very highly polarised
  • The H becomes so δ+ charged that it can form a bond with the lone pair of an O or N atom in another molecule

-For hydrogen bonding to take place, the angle between the -OH/-NH and the hydrogen bond is 180o

47
Q

The number of hydrogen bonds depends on

A
  • The number of hydrogen atoms attached to O or N in the molecule
  • The number of lone pairs on the O or N
48
Q

Properties of water

A

-Hydrogen bonding in water, causes it to have anomalous properties such as high melting and boiling points, high surface tension and anomalous density of ice compared to water

49
Q

High melting & boiling points of water

A
  • Water has high melting and boiling points which is caused by the strong intermolecular forces of hydrogen bonding between the molecules
  • In ice (solid H2O) and water (liquid H2O) the molecules are tightly held together by hydrogen bonds
  • A lot of energy is therefore required to break the water molecules apart and melt or boil them
50
Q

High surface tension of water

A
  • Water has a high surface tension
  • Surface tension is the ability of a liquid surface to resist any external forces (i.e. to stay unaffected by forces acting on the surface)
  • The water molecules at the surface of liquid are bonded to other water molecules through hydrogen bonds
  • These molecules pull downwards on the surface molecules causing the surface them to become compressed and more tightly together at the surface
  • This increases water’s surface tension

All because of hydrogen bonding between water molecules

51
Q

Density

A
  • Solids are denser than their liquids as the particles in solids are more closely packed together than in their liquid state
  • In ice however, the water molecules are packed in a 3D hydrogen-bonded network in a rigid lattice
  • Each oxygen atom is surrounded by hydrogen atoms
  • This way of packing the molecules in a solid and the relatively long bond lengths of the hydrogen bonds means that the water molecules are slightly further apart than in the liquid form
  • Therefore, ice has a lower density than liquid water
52
Q

The enthalpy changes increase going from H2S to H2Te

A

due to the increased number of electrons in the Group 16 elements

  • This causes an increased instantaneous dipole – induced dipole forces as the molecules become larger
  • Based on this, H2O would have a much lower enthalpy change (around 17 kJ mol-1)
  • However, the enthalpy change of vaporisation is almost 3 times larger which is caused by the hydrogen bonds present in water but not in the other hydrides
53
Q

enthalpy of vaporisation

A

energy required to boil a substance

54
Q

Electronegativity

A

is the ability of an atom to draw a pair of electrons towards itself in a covalent bond
-Electronegativity increases across a Period and decreases going down a Group

55
Q

Polarity

A
  • When two atoms in a covalent bond have the same electronegativity the covalent bond is nonpolar
  • When two atoms in a covalent bond have different electronegativities the covalent bond is polar and the electrons will be drawn towards the more electronegative atom
  • As a result of this:
  • –The negative charge centre and positive charge centre do not coincide with each other
  • –This means that the electron distribution is asymmetric
  • –The less electronegative atom gets a partial charge of δ+ (delta positive)
  • –The more electronegative atom gets a partial charge of δ- (delta negative)
56
Q

Dipole moment

A
  • The dipole moment is a measure of how polar a bond is
  • The direction of the dipole moment is shown by the following sign in which the arrow points to the partially negatively charged end of the dipole:

-The sign shows the direction of the dipole moment and the arrow points to the delta negative end of the dipole

57
Q

Assigning polarity to molecules

A

-To determine whether a molecule with more than two atoms is polar, the following things have to be taken into consideration:

  • –The polarity of each bond
  • –How the bonds are arranged in the molecule

-Some molecules have polar bonds but are overall not polar because the polar bonds in the molecule are arranged in such way that the individual dipole moments cancel each other out

58
Q

van der Waals’ Forces

A
  • Covalent bonds are strong intramolecular forces
  • Molecules also contain weaker intermolecular forces which are forces between molecules
  • These intermolecular forces are called van der Waals’ forces
  • There are two types of van der Waals’ forces:
  • —Instantaneous (temporary) dipole – induced dipole forces also called London dispersion forces
  • —Permanent dipole – permanent dipole forces
59
Q

Instantaneous dipole – induced dipole (id – id)

A
  • Instantaneous dipole – induced dipole forces or London dispersion forces exist between all atoms or molecules
  • The electron charge cloud in non-polar molecules or atoms are constantly moving
  • During this movement, the electron charge cloud can be more on one side of the atom or molecule than the other
  • This causes a temporary dipole to arise
  • This temporary dipole can induce a dipole on neighbouring molecules
  • When this happens, the δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other
  • Because the electron clouds are moving constantly, the dipoles are only temporary

Weakest attraction

London force is part of all bonding

60
Q

Hydrogen Bonding as a Permanent Dipole

A
  • Hydrogen bonding is an intermolecular force between molecules with an -OH/-NH group and molecules with an N/O atom
  • Hydrogen bonding is a special case of a permanent dipole – dipole force between molecules
  • —Hydrogen bonds are stronger forces than pd – pd forces
  • The hydrogen is bonded to an O/N atom which is so electronegative, that almost all the electron density from the covalent bond is drawn towards the O/N atom
  • This leaves the H with a large delta positive and the O/N with a large delta negative charging resulting in the formation of a permanent dipole in the molecule
  • A delta positive H in one molecule is electrostatically attracted to the delta negative O/N in a neighbouring molecule

Strongest attraction

61
Q

Permanent dipole – permanent dipole (pd – pd)

A
  • Polar molecules have permanent dipoles
  • The molecule will always have a negatively and positively charged end
  • Forces between two molecules that have permanent dipoles are called permanent dipole – permanent dipole forces

-The δ+ end of the dipole in one molecule and the δ- end of the dipole in a neighbouring molecule are attracted towards each other

Medium strength
Forms when two elements bond which have different electro-negativity

62
Q

For small molecules with the same number of electrons, pd – pd forces are stronger than id – id

A
  • –Butane and propanone have the same number of electrons
  • –Butane is a nonpolar molecule and will have id – id forces
  • –Propanone is a polar molecule and will have pd – pd forces
  • –Therefore, more energy is required to break the intermolecular forces between propanone molecules than between butane molecules

So, propanone has a higher boiling point than butane

63
Q

Intramolecular forces

A
  • Intramolecular forces are forces within a molecule
  • Ionic bonding is the electrostatic attraction between positive (cations) and negative (anions) ions in an ionic crystal lattice
  • —These ions are formed by transferring the electrons from one species to the other
  • Covalent bonds are formed when the outer electrons of two atoms are shared
  • Metallic bonding is the electrostatic attraction of positively charged metal ions and their delocalised electrons in a metal lattice

Form strong covalent bonds

64
Q

Intermolecular forces

A

-Intramolecular forces are forces between molecules and are also called van der Waals’s forces

Intermolecular forces are responsible for melting and boiling point by of simple covalent bonds

65
Q

intramolecular forces are

A

stronger than intermolecular forces

66
Q

The strengths of the types of bond or force are as follows: Strongest > weakest

A
  • Metallic bonding
  • ionic bonding
  • covalent
  • hydrogen bonding
  • permanent dipole - permanent dipole forces
  • Instantaneous dipole - induced dipole
67
Q

Dot & Cross Diagrams

A

are diagrams that show the arrangement of the outer-shell electrons in an ionic or covalent compound or element
The electrons are shown as dots and crosses

68
Q

In a dot-and-cross diagram

A
  • –Only the outer electrons are shown
  • –The charge of the ion is spread evenly which is shown by using brackets
  • –The charge on each ion is written at the top right-hand corner
69
Q

Ionic compounds atoms try achieve a ….

A
  • Ionic bonds are formed when metal transfer electrons to a non-metal to form a positively charged and negatively charged ion
  • The atoms achieve a noble gas configuration
70
Q

Coordinate bonding

A

-Coordinate bonding or also called dative covalent bonding is formed when one atom provides both the electrons needed for a covalent bond

  • In a displayed formula, the dative covalent bond is represented by an arrow
  • The head of the arrow points away from the lone pair that forms the bond
71
Q

Electrostatic force: ionic bond

A

Proportional to the change of ions

Inversely proportional to the square of distance between ions

E ∝ (Q^+) + (Q^-) // d^2

Attraction between cations and anions

72
Q

Electrostatic force: covalent bond

A

Electrostatic force between neighbouring nuclei and shared electron

No electron transfer but there is electron sharing

73
Q

Exception to bonding: extended- octet

A

From period 3 onwards the d orbital can form part of the bonding with leads to more then 8 electrons in the outer most valance shell

Why? Empty d orbital

74
Q

Factors affecting London forces

A

1) size of molecules
- size bigger = more electrons
- easier to from a temporary dipole
- more energy required to overcome L.F
2) contact point
- greater contact points = higher London forces between the molecules

75
Q

Factors affecting London forces

A

1) size of molecules
- size bigger = more electrons
- easier to from a temporary dipole
- more energy required to overcome L.F
2) contact point
- greater contact points = higher London forces between the molecules

Chemistry AS Level CIE (21 decks)

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