1.6 bonding
DESCRIPTION
as level chemistry bondingTRANSCRIPT
5. Bonding and Intermolecular ForcesTeacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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non-metal atom (e.g. NaCl)
a covalent bond occurs between two non-metal atoms (e.g. I2, CH4)
a metallic bond occurs between atoms in a metal (e.g. Cu)
There are three types of bond that can occur between atoms:
Photo credit (left): Vladimir Zivkovic / shutterstock.com
Microscope shot of salt (NaCl) crystals.
Photo credit (middle): Dr John Mileham
Iodine crystals
Copper wire
Teacher notes
It could be stressed to students that covalent bonding does not have to occur between two atoms of the same element, as is the case in iodine. It should also be pointed out that while a single ionic or covalent bond occurs between just two atoms, multiple ionic bonds exist in an ionic lattice, and multiple covalent bonds occur in covalent compounds containing more than two atoms.
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non-metals gain electrons to form negative ions.
The elements in groups 4 and 8 (also called group 0) do not gain or lose electrons to form ionic compounds.
1
2
3
4
5
6
7
8/0
1+
2+
3+
N/A
3-
2-
1-
N/A
Na+
Al3+
N/A
N/A
N3-
O2-
F-
Mg2+
The number of electrons gained or lost by an atom is related to the group in which the element is found.
Group
Charge
Example
Teacher notes
A hydroxonium ion is sometimes called a hydronium ion or oxonium ion. It is a more accurate way of representing an aqueous H+ ion.
‘Dimerizes’ (as used here in reference to aluminium chloride) means to form a dimer – a molecule made up of two identical subunits (in this case two AlCl3 subunits). Al2Cl3 forms when AlCl3 melts, and can vapourize.
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The strength of metallic bonding depends on two factors:
1. the charge on the metal ions
1. The charge on the metal ions
The greater the charge on the metal ions, the greater the attraction between the ions and the delocalized electrons, and the stronger the metallic bonds. A higher melting point is evidence of stronger bonds in the substance.
2. the size of the metal ions.
Na
Mg
Al
1+
2+
3+
371
923
933
Element
Li
Na
K
Rb
Cs
0.076
0.102
0.138
0.152
0.167
454
371
337
312
302
2. The size of the metal ions
The smaller the metal ion, the closer the positive nucleus is to the delocalized electrons. This means there is a greater attraction between the two, which creates a stronger metallic bond.
Element
point (K)
Teacher notes
The ionic radius of an ion is not fixed, but depends on several factors, such as the co-ordination number of the ion.
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Teacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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Electronegativity values for some common elements. Values given here are measured on the Pauling scale.
In a covalent bond between two different elements, the electron density is not shared equally.
This is because different elements have differing abilities to attract the bonding electron pair. This ability is called an element’s electronegativity.
Teacher notes
Elements in group 8 do not commonly form bonds so electronegativity values have not been measured.
Electronegativity was first proposed in 1932 by the American chemist Linus Pauling (1904-1994). It is after him that the scale used here for measuring electronegativity is named. While the Pauling scale is the most common, there are other measures of electronegativity in use.
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Electronegativity and atomic radius
The electronegativity of an element depends on a combination of two factors:
1. Atomic radius
As radius of an atom increases, the bonding pair of electrons become further from the nucleus. They are therefore less attracted to the positive charge of the nucleus, resulting in a lower electronegativity.
higher electronegativity
lower electronegativity
Teacher notes
See the ‘Trends in Period 3’ presentation for more information about atomic radius.
Note that the inner electron levels have not been included in the diagram.
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2. The number of unshielded protons
The greater the number of protons in a nucleus, the greater the attraction to the electrons in the covalent bond, resulting in higher electronegativity.
However, full energy levels of electrons shield the electrons in the bond from the increased attraction of the greater nuclear charge, thus reducing electronegativity.
greater nuclear charge increases electronegativity…
…but extra shell of electrons increases shielding.
Teacher notes
See the ‘Trends in Period 3’ presentation for more information about shielding.
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Electronegativity increases across a period because:
1. The atomic radius decreases.
2. The charge on the nucleus increases without significant extra shielding. New electrons do not contribute much to shielding because they are added to the same principal energy level across the period.
Teacher notes
Fluorine has, with a Pauling electronegativity value of 4.0, the highest electronegativity of all the elements.
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Electronegativity trends: down a group
2. Although the charge on the nucleus increases, shielding also increases significantly. This is because electrons added down the group fill new principal energy levels.
Electronegativity decreases down a group because:
1. The atomic radius increases.
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Non-polar bonds
If the electronegativity of both atoms in a covalent bond is identical, the electrons in the bond will be equally attracted to both of them.
This results in a symmetrical distribution of electron density around the two atoms.
Bonding in elements (for example O2 or Cl2) is always non-polar because the electronegativity of the atoms in each molecule is the same.
both atoms are equally good at attracting the electron density
cloud of electron density
Effect of electronegativity on polarization
The greater the electronegativity difference between the two atoms in a bond the greater the polarization of the bond.
decreasing polarization
This can be illustrated by looking at the hydrogen halides:
H
F
Cl
Br
I
© Boardworks Ltd 2009
Ionic or covalent?
Rather than saying that ionic and covalent are two distinct types of bonding, it is more accurate to say that they are at the two extremes of a scale.
Less polar bonds have more covalent character.
increasing polarization
More polar bonds have more ionic character. The more electronegative atom attracts the electrons in the bond enough to ionize the other atom.
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Molecules containing polar bonds are not always polar.
If the polar bonds are arranged symmetrically, the partial charges cancel out and the molecule is non-polar.
Non-polar molecules
If the polar bonds are arranged asymmetrically, the partial charges do not cancel out and the molecule is polar.
Polar molecules
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Teacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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hydrogen bonds – for example, found between H2O molecules in water.
permanent dipole–dipole forces – for example, found between HCl molecules in hydrogen chloride.
van der Waals forces – for example, found between I2 molecules in iodine crystals.
There are three main types of intermolecular force:
The molecules in simple covalent substances are not entirely isolated from one another. There are forces of attraction between them. These are called intermolecular forces.
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Strength of van der Waals forces
The strength of van der Waals forces increases as molecular size increases.
Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce.
0
50
100
150
200
-50
-100
-150
-200
Br2
This is illustrated by the boiling points of group 7 elements.
F2
Cl2
I2
element
Strength of van der Waals forces
Straight chain alkanes can pack closer together than branched alkanes, creating more points of contact between molecules. This results in stronger van der Waals forces.
butane (C4H10)
boiling point = 261 K
The points of contact between molecules also affects the strength of van der Waals forces.
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Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the molecules may align so there is electrostatic attraction between the opposite charges on neighbouring molecules.
Permanent dipole–dipole forces (dotted lines) occur in hydrogen chloride (HCl) gas.
The permanent dipole–dipole forces are approximately one hundredth the strength of a covalent bond.
Teacher notes
Students could be reminded that molecules need to be polar, i.e. have an overall dipole, for permanent dipole–dipole forces to occur.
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Teacher notes
phosphine (PH3) = exhibits permanent dipole–dipole forces
hydrogen sulfide (H2S) = exhibits permanent dipole–dipole forces
chlorine (Cl2) = does not exhibit permanent dipole–dipole forces
ozone (O3) = does not exhibit permanent dipole–dipole forces
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What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than in other polar bonds.
This is because these atoms are highly electronegative due to their high nuclear charge and small size. When these atoms bond to hydrogen, electrons are withdrawn from the H atom, making it slightly positive.
Hydrogen bonds are therefore particularly strong examples of permanent dipole–dipole forces.
The H atom is very small so the positive charge is more concentrated, making it easier to link with other molecules.
Teacher notes
Hydrogen bonds are about one tenth the strength of a covalent bond.
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Hydrogen bonding
In molecules with OH or NH groups, a lone pair of electrons on nitrogen or oxygen is attracted to the slight positive charge on the hydrogen on a neighbouring molecule.
Hydrogen bonding makes the melting and boiling points of water higher than might be expected. It also means that alcohols have much higher boiling points than alkanes of a similar size.
hydrogen bond
lone pair
Teacher notes
See the ‘Alcohols’ presentation for more information about the structure and properties of alcohols.
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Boiling points of the hydrogen halides
The boiling point of hydrogen fluoride is much higher than that of other hydrogen halides, due to fluorine’s high electronegativity.
0
20
40
-20
-40
-60
-80
-100
HF
HCl
HBr
HI
The means that hydrogen bonding between molecules of hydrogen fluoride is much stronger than the permanent dipole–dipole forces between molecules of other hydrogen halides. More energy is therefore required to separate the molecules of hydrogen fluoride.
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Teacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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Teacher notes
co-ordinate bond – A covalent bond formed by the donation of a lone pair of electrons from one atom into an empty orbital on another atom. Also called a dative covalent bond.
covalent bond – The attraction between two non-metal atoms caused by the overlapping of electron orbitals resulting in the sharing of one or more pairs of electrons between the atoms.
dative covalent bond – A covalent bond formed by the donation of a lone pair of electrons from one atom into an empty orbital on another atom. Also called a co-ordinate bond.
dipole – The unequal sharing of electron density in a molecule resulting in regions of partial positive charge and regions of partial negative charge. Dipoles may be permanent, temporary or induced.
electron density – A measure of how likely it is that an electron will be found at a specific location within an orbital.
electronegativity – The power of an atom to withdraw electron density and attract the pair of bonding electrons in a covalent bond.
hydrogen bonding – A special example of a permanent dipole–dipole intermolecular force that occurs when hydrogen is bonded to nitrogen, oxygen or fluorine. A hydrogen bond is the attraction between a lone pair on a nitrogen, oxygen or fluorine atom on one molecule, and the δ+ hydrogen atom (bonded to nitrogen, oxygen or fluorine) on a neighbouring molecule.
induced dipole – The result of a non-polar molecule moving close to a molecule with a dipole. The dipole in one molecule distorts the electron density in the other, creating a dipole in that molecule. This occurs in the formation of van der Waals forces.
intermolecular force – An attraction between two or more covalent molecules. The three main types are van der Waals forces, permanent dipole–dipole forces and hydrogen bonds.
ionic bond – The electrostatic attraction between oppositely-charged ions, which are formed by the transfer of electrons from a metal to a non-metal. This force holds the ions together in an ionic lattice.
ionic lattice – A regular arrangement of positively and negatively-charged ions held together by electrostatic forces (ionic bonds) acting between them.
lone pair – A pair of electrons that is not involved in a covalent bond.
metallic bond – The attraction between positively-charged metal ions and a sea of electrons from their highest energy levels that have become delocalized. This force holds metal atoms together in a giant metallic lattice.
metallic lattice – A regular arrangement of positively-charged metal ions surrounded by a sea of delocalized electrons and held together by the attraction between them (metallic bonds).
non-polar bond – A covalent bond in which the electron density is distributed symmetrically around both atoms.
non-polar molecule – A molecule in which either there are no polar bonds, or the polar bonds are arranged symmetrically so that the partial charges cancel out and the molecule has no overall permanent dipole.
permanent dipole–dipole forces – Intermolecular forces between the opposite charges on molecules with permanent dipoles.
polar bond – A covalent bond between two atoms with different electronegativities, which causes an asymmetrical distribution (unequal sharing) of the electron density. This, in turn, creates a slight positive charge on one atom and a slight negative charge on the other.
polar molecule – A molecule containing polar bonds arranged asymmetrically so that the partial charges do not cancel out. The molecule therefore has an overall dipole.
temporary dipole – A spontaneous and transient (constantly changing) dipole resulting from the chance uneven distribution of electron density of a molecule.
van der Waals forces – Weak intermolecular forces of attraction that occur when temporary dipoles cause induced dipoles in neighbouring atoms.
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In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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non-metal atom (e.g. NaCl)
a covalent bond occurs between two non-metal atoms (e.g. I2, CH4)
a metallic bond occurs between atoms in a metal (e.g. Cu)
There are three types of bond that can occur between atoms:
Photo credit (left): Vladimir Zivkovic / shutterstock.com
Microscope shot of salt (NaCl) crystals.
Photo credit (middle): Dr John Mileham
Iodine crystals
Copper wire
Teacher notes
It could be stressed to students that covalent bonding does not have to occur between two atoms of the same element, as is the case in iodine. It should also be pointed out that while a single ionic or covalent bond occurs between just two atoms, multiple ionic bonds exist in an ionic lattice, and multiple covalent bonds occur in covalent compounds containing more than two atoms.
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non-metals gain electrons to form negative ions.
The elements in groups 4 and 8 (also called group 0) do not gain or lose electrons to form ionic compounds.
1
2
3
4
5
6
7
8/0
1+
2+
3+
N/A
3-
2-
1-
N/A
Na+
Al3+
N/A
N/A
N3-
O2-
F-
Mg2+
The number of electrons gained or lost by an atom is related to the group in which the element is found.
Group
Charge
Example
Teacher notes
A hydroxonium ion is sometimes called a hydronium ion or oxonium ion. It is a more accurate way of representing an aqueous H+ ion.
‘Dimerizes’ (as used here in reference to aluminium chloride) means to form a dimer – a molecule made up of two identical subunits (in this case two AlCl3 subunits). Al2Cl3 forms when AlCl3 melts, and can vapourize.
7849.unknown
7850.unknown
The strength of metallic bonding depends on two factors:
1. the charge on the metal ions
1. The charge on the metal ions
The greater the charge on the metal ions, the greater the attraction between the ions and the delocalized electrons, and the stronger the metallic bonds. A higher melting point is evidence of stronger bonds in the substance.
2. the size of the metal ions.
Na
Mg
Al
1+
2+
3+
371
923
933
Element
Li
Na
K
Rb
Cs
0.076
0.102
0.138
0.152
0.167
454
371
337
312
302
2. The size of the metal ions
The smaller the metal ion, the closer the positive nucleus is to the delocalized electrons. This means there is a greater attraction between the two, which creates a stronger metallic bond.
Element
point (K)
Teacher notes
The ionic radius of an ion is not fixed, but depends on several factors, such as the co-ordination number of the ion.
* of 43
Teacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
* of 43
Electronegativity values for some common elements. Values given here are measured on the Pauling scale.
In a covalent bond between two different elements, the electron density is not shared equally.
This is because different elements have differing abilities to attract the bonding electron pair. This ability is called an element’s electronegativity.
Teacher notes
Elements in group 8 do not commonly form bonds so electronegativity values have not been measured.
Electronegativity was first proposed in 1932 by the American chemist Linus Pauling (1904-1994). It is after him that the scale used here for measuring electronegativity is named. While the Pauling scale is the most common, there are other measures of electronegativity in use.
* of 43
Electronegativity and atomic radius
The electronegativity of an element depends on a combination of two factors:
1. Atomic radius
As radius of an atom increases, the bonding pair of electrons become further from the nucleus. They are therefore less attracted to the positive charge of the nucleus, resulting in a lower electronegativity.
higher electronegativity
lower electronegativity
Teacher notes
See the ‘Trends in Period 3’ presentation for more information about atomic radius.
Note that the inner electron levels have not been included in the diagram.
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2. The number of unshielded protons
The greater the number of protons in a nucleus, the greater the attraction to the electrons in the covalent bond, resulting in higher electronegativity.
However, full energy levels of electrons shield the electrons in the bond from the increased attraction of the greater nuclear charge, thus reducing electronegativity.
greater nuclear charge increases electronegativity…
…but extra shell of electrons increases shielding.
Teacher notes
See the ‘Trends in Period 3’ presentation for more information about shielding.
* of 43
Electronegativity increases across a period because:
1. The atomic radius decreases.
2. The charge on the nucleus increases without significant extra shielding. New electrons do not contribute much to shielding because they are added to the same principal energy level across the period.
Teacher notes
Fluorine has, with a Pauling electronegativity value of 4.0, the highest electronegativity of all the elements.
* of 43
Electronegativity trends: down a group
2. Although the charge on the nucleus increases, shielding also increases significantly. This is because electrons added down the group fill new principal energy levels.
Electronegativity decreases down a group because:
1. The atomic radius increases.
* of 43
Non-polar bonds
If the electronegativity of both atoms in a covalent bond is identical, the electrons in the bond will be equally attracted to both of them.
This results in a symmetrical distribution of electron density around the two atoms.
Bonding in elements (for example O2 or Cl2) is always non-polar because the electronegativity of the atoms in each molecule is the same.
both atoms are equally good at attracting the electron density
cloud of electron density
Effect of electronegativity on polarization
The greater the electronegativity difference between the two atoms in a bond the greater the polarization of the bond.
decreasing polarization
This can be illustrated by looking at the hydrogen halides:
H
F
Cl
Br
I
© Boardworks Ltd 2009
Ionic or covalent?
Rather than saying that ionic and covalent are two distinct types of bonding, it is more accurate to say that they are at the two extremes of a scale.
Less polar bonds have more covalent character.
increasing polarization
More polar bonds have more ionic character. The more electronegative atom attracts the electrons in the bond enough to ionize the other atom.
* of 43
Molecules containing polar bonds are not always polar.
If the polar bonds are arranged symmetrically, the partial charges cancel out and the molecule is non-polar.
Non-polar molecules
If the polar bonds are arranged asymmetrically, the partial charges do not cancel out and the molecule is polar.
Polar molecules
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Teacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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hydrogen bonds – for example, found between H2O molecules in water.
permanent dipole–dipole forces – for example, found between HCl molecules in hydrogen chloride.
van der Waals forces – for example, found between I2 molecules in iodine crystals.
There are three main types of intermolecular force:
The molecules in simple covalent substances are not entirely isolated from one another. There are forces of attraction between them. These are called intermolecular forces.
* of 43
Strength of van der Waals forces
The strength of van der Waals forces increases as molecular size increases.
Atomic radius increases down the group, so the outer electrons become further from the nucleus. They are attracted less strongly by the nucleus and so temporary dipoles are easier to induce.
0
50
100
150
200
-50
-100
-150
-200
Br2
This is illustrated by the boiling points of group 7 elements.
F2
Cl2
I2
element
Strength of van der Waals forces
Straight chain alkanes can pack closer together than branched alkanes, creating more points of contact between molecules. This results in stronger van der Waals forces.
butane (C4H10)
boiling point = 261 K
The points of contact between molecules also affects the strength of van der Waals forces.
* of 43
Permanent dipole–dipole forces
If molecules contain bonds with a permanent dipole, the molecules may align so there is electrostatic attraction between the opposite charges on neighbouring molecules.
Permanent dipole–dipole forces (dotted lines) occur in hydrogen chloride (HCl) gas.
The permanent dipole–dipole forces are approximately one hundredth the strength of a covalent bond.
Teacher notes
Students could be reminded that molecules need to be polar, i.e. have an overall dipole, for permanent dipole–dipole forces to occur.
* of 43
Teacher notes
phosphine (PH3) = exhibits permanent dipole–dipole forces
hydrogen sulfide (H2S) = exhibits permanent dipole–dipole forces
chlorine (Cl2) = does not exhibit permanent dipole–dipole forces
ozone (O3) = does not exhibit permanent dipole–dipole forces
7939.unknown
What is hydrogen bonding?
When hydrogen bonds to nitrogen, oxygen or fluorine, a larger dipole occurs than in other polar bonds.
This is because these atoms are highly electronegative due to their high nuclear charge and small size. When these atoms bond to hydrogen, electrons are withdrawn from the H atom, making it slightly positive.
Hydrogen bonds are therefore particularly strong examples of permanent dipole–dipole forces.
The H atom is very small so the positive charge is more concentrated, making it easier to link with other molecules.
Teacher notes
Hydrogen bonds are about one tenth the strength of a covalent bond.
* of 43
Hydrogen bonding
In molecules with OH or NH groups, a lone pair of electrons on nitrogen or oxygen is attracted to the slight positive charge on the hydrogen on a neighbouring molecule.
Hydrogen bonding makes the melting and boiling points of water higher than might be expected. It also means that alcohols have much higher boiling points than alkanes of a similar size.
hydrogen bond
lone pair
Teacher notes
See the ‘Alcohols’ presentation for more information about the structure and properties of alcohols.
* of 43
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Boiling points of the hydrogen halides
The boiling point of hydrogen fluoride is much higher than that of other hydrogen halides, due to fluorine’s high electronegativity.
0
20
40
-20
-40
-60
-80
-100
HF
HCl
HBr
HI
The means that hydrogen bonding between molecules of hydrogen fluoride is much stronger than the permanent dipole–dipole forces between molecules of other hydrogen halides. More energy is therefore required to separate the molecules of hydrogen fluoride.
* of 43
Teacher notes
In ‘Slide Show’ mode, click the name of a section to jump straight to that slide.
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Teacher notes
co-ordinate bond – A covalent bond formed by the donation of a lone pair of electrons from one atom into an empty orbital on another atom. Also called a dative covalent bond.
covalent bond – The attraction between two non-metal atoms caused by the overlapping of electron orbitals resulting in the sharing of one or more pairs of electrons between the atoms.
dative covalent bond – A covalent bond formed by the donation of a lone pair of electrons from one atom into an empty orbital on another atom. Also called a co-ordinate bond.
dipole – The unequal sharing of electron density in a molecule resulting in regions of partial positive charge and regions of partial negative charge. Dipoles may be permanent, temporary or induced.
electron density – A measure of how likely it is that an electron will be found at a specific location within an orbital.
electronegativity – The power of an atom to withdraw electron density and attract the pair of bonding electrons in a covalent bond.
hydrogen bonding – A special example of a permanent dipole–dipole intermolecular force that occurs when hydrogen is bonded to nitrogen, oxygen or fluorine. A hydrogen bond is the attraction between a lone pair on a nitrogen, oxygen or fluorine atom on one molecule, and the δ+ hydrogen atom (bonded to nitrogen, oxygen or fluorine) on a neighbouring molecule.
induced dipole – The result of a non-polar molecule moving close to a molecule with a dipole. The dipole in one molecule distorts the electron density in the other, creating a dipole in that molecule. This occurs in the formation of van der Waals forces.
intermolecular force – An attraction between two or more covalent molecules. The three main types are van der Waals forces, permanent dipole–dipole forces and hydrogen bonds.
ionic bond – The electrostatic attraction between oppositely-charged ions, which are formed by the transfer of electrons from a metal to a non-metal. This force holds the ions together in an ionic lattice.
ionic lattice – A regular arrangement of positively and negatively-charged ions held together by electrostatic forces (ionic bonds) acting between them.
lone pair – A pair of electrons that is not involved in a covalent bond.
metallic bond – The attraction between positively-charged metal ions and a sea of electrons from their highest energy levels that have become delocalized. This force holds metal atoms together in a giant metallic lattice.
metallic lattice – A regular arrangement of positively-charged metal ions surrounded by a sea of delocalized electrons and held together by the attraction between them (metallic bonds).
non-polar bond – A covalent bond in which the electron density is distributed symmetrically around both atoms.
non-polar molecule – A molecule in which either there are no polar bonds, or the polar bonds are arranged symmetrically so that the partial charges cancel out and the molecule has no overall permanent dipole.
permanent dipole–dipole forces – Intermolecular forces between the opposite charges on molecules with permanent dipoles.
polar bond – A covalent bond between two atoms with different electronegativities, which causes an asymmetrical distribution (unequal sharing) of the electron density. This, in turn, creates a slight positive charge on one atom and a slight negative charge on the other.
polar molecule – A molecule containing polar bonds arranged asymmetrically so that the partial charges do not cancel out. The molecule therefore has an overall dipole.
temporary dipole – A spontaneous and transient (constantly changing) dipole resulting from the chance uneven distribution of electron density of a molecule.
van der Waals forces – Weak intermolecular forces of attraction that occur when temporary dipoles cause induced dipoles in neighbouring atoms.
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