topic 2 - bonding and structure Flashcards
1
Q
what is the definition of a chemical bond
A
a force of electrostatic attraction between positive nuclei and negative electrons which hold together 2 or more atoms, ions or molecules
2
Q
what is the definition of valency
A
the ability for an atom to form bonds
3
Q
what is the valence shell and valence electrons
A
valence electrons are electrons that form bonds
the valence shell contains the valence electrons, so its the highest energy outermost shell
4
Q
what is the definition of an ionic bond
A
the electrostatic force of attraction between anions and cations
these ions are formed from the loss and gain of electrons
5
Q
in ionic bonding, what can be said about the force of attraction and repulsion
A
the ions are arranged in a lattice structure so that that the electrostatic force of attraction between oppositely charged ions is greater than the electrostatic repulsion between ions with the same charge
6
Q
what can make ionic bonding stronger
A
-the smaller the ionic radius, the stronger the electrostatic force of attraction between oppositely charged ions, therefore more energy is required to break the ionic bond
-the greater the charges on the ions the stronger the ionic bonding
7
Q
explain what happens to atomic radius as the number of protons increase
A
-as proton number increases atomic radius decreases
-because the positive charge of the nucleus will increase, so the electrostatic attraction between the nucleus and electrons increases, so electrons are pulled closer to the nucleus
8
Q
what are the 5 physical properties of ionic compounds
A
high melting point
brittleness
poor conductors of electricity when solid
good conductors of electricity when molten or aqueous
soluble in water
9
Q
explain why ionic compounds have a high melting point
A
-giant regular lattice structure
-containing electrostatic forces of attraction between anions and cations
-these forces are very strong and need a lot of energy to break
10
Q
explain why ionic compound are brittle
A
-when a force is applied to a solid ionic compound the layers of ions slide over eachother
-causing ions of the same charge to be side by side
-they repel each other
-so the solid breaks apart
11
Q
explain why solid ionic compounds do not conduct electricity
A
-the ions are not free to move and there are no delocalised electrons
12
Q
explain why molten or aqueous ionic compounds do conduct electricity
A
-the ions in the giant regular lattice structure are free to move
-the ions will migrate to the electrode with the opposite charge when a potential difference is supplied
-if direct current is used the ionic compound will undergo electrolysis, causing the ions to be discharged at the electrodes
13
Q
explain why ionic compounds are soluble
A
-the energy needed to break apart the lattice structure and separate the ions is supplied by the hydration of separated ions
-the delta negative oxygen in water molecules is attracted to cations in the ionic compound
-delta positive hydrogen in water molecules are attracted to anions in the ionic compound
14
Q
what are isolecetronic ions
A
ions which all have the same number of electrons
but can be different elements
15
Q
why is there variation in ionic radius in a set of isoelectronic ions
A
–as proton number increases ionic radius decreases
-because the positive charge of the nucleus will increase, so the electrostatic attraction between the nucleus and electrons increases, so electrons are pulled closer to the nucleus
16
Q
what is 1 piece of evidence for the existence of ions
A
the ability for an ionic compound to conduct electricity and undergo electrolysis
17
Q
explain the electrolysis of molten sodium chloride
A
-positive sodium cations migrate towards the negative cathode, where they gain electrons (reduction), becoming sodium atoms
-negative chlorine anions migrate towards the positive anode, where they lose electrons (oxidation), becoming chlorine atoms
18
Q
describe how the movement of ions can be demonstrated visually using copper(II) chromate(VI) solution
A
-aqueous copper(II) ions are blue
-aqueous chromate(VI) ions are yellow
-copper(II) cations migrate towards the cathode, and the solution around the cathode turns blue
-chromate(VI) anions migrate towards the anode, and the solution around the anode turns yellow
19
Q
how is a sigma bond formed
A
end on overlap of 2 orbitals which both contain 1 electron only
20
Q
how is a pi bond formed
A
sideways overlap of 2 p-orbitals
pi bonds can only form once a sigma bond has formed
so pi bonds only exist between atoms which have double/triple bonds
21
Q
what type of bond is a single covalent bond between 2 atoms
A
sigma bond
end on overlap of 2 orbitals
22
Q
what are the bonds in a triple covalent bond
A
1 sigma bond
2 pi bonds
23
Q
what is bond length
A
the distance between the nuclei of 2 atoms covalently bonded together
24
Q
when comparing covalent bonds formed between 2 atoms of the same element, what is the relationship between bond length and bond strength
A
the shorter the bond the greater the bond strength
because there is an increased electrostatic attraction between the 2 nuclei and the electrons in the overlapping atomic orbitals
25
what can be said about the bond length and bond strength of single, double, and triple covalent bonds
single covalent bonds: longest bond length + weakest strength
double covalent bonds: middle length + middle strength
triple covalent bonds: shortest bond length + strongest strength
the more bonds between 2 atoms of the same element, the shorter the bond length, and the higher the bond strength
26
what is the definition of a covalent bond
the electrostatic attraction between 2 nuclei and the shared pair of electrons between them
27
what is the definition of electronegativity
the ability of an atom to attract a bonding pair of electrons
28
if 2 atoms of the same element are bonded together by the overlap of atomic orbitals, how can the distribution of electron density between the 2 nuclei be described as
-the distribution of electron density will be symmetrical
-because both atoms will have the same electronegativity, so equally share the pair of electrons in the covalent bond
-so the bonding is purely covalent
29
where on the periodic table are the most and least electronegative elements
most electronegative: top right
non-metals are more electronegative than metals
least electronegative: bottom left
30
what is a polar covalent bond
-when 2 atoms bonded together are different elements, so have different electronegativities
-so the distribution of electron density will not be symmetrical
-the more electronegative atom will attract the pair of electrons in the covalent bond more strongly
-so the more electronegative atom will become delta negative
-and the other atom will become more delta positive
-this creates a polare covalent bond
31
how do you calculate to what percentage a bond is covalent or ionic
-work out the difference is electronegativity between the 2 bonded atoms
-find this value in the electronegativity difference column of the table
-read along to the the percentage ionic and covalent character
32
are polar bonds 100% covalent
-no
-ionic and covalent are 2 extremes of a continuum of bonding type
-due to differences in electronegativity between bonded atoms, it creates polar bonds which are not 100% covalent or ionic
32
what are the 5 physical properties of metals
high melting point
good electrical conductors
good thermal conductors
malleable
ductile
33
how is the electrical conductivity of a metal increased
as the number of electrons in the outer shell of the metal cations increases, electrical conductivity increases
34
what is the definition of metallic bonding
the electrostatic force of attraction between the nuclei of metal cations and delocalised electrons
35
describe the structure of metallic bonding
-layers of metal cations arranged in a regular giant lattice structure
-with a sea of delocalised electrons
36
explain why metals have a high melting point
-metals have a giant regular lattice structure where there are many electrostatic forces of attraction between the nuclei of the metal cations and delocalised electrons
-a lot of energy is required to break these forces
-the more delocalised electrons per cation the higher the melting point
-the smaller the cation the larger the melting point
37
what 2 factors can increase the melting point of metals
-the more delocalised electrons per cation the higher the melting point
-group 1 metals have 1 delocalised electron per cation so have a low melting point
-group 2 metals have 2 delocalised electrons per cation so have a higher melting point
-d-block metals have the highest melting point because they have more delocalised electrons per cation
-the smaller the cation the larger the melting point
-because the smaller the cation the closer the delocalised electrons are to the nucleus, so the electrostatic force of attraction between the nuclei of the cations and delocalised electrons increase, so more energy is required to break the force
38
explain why metals are good electrical conductors
when a potential difference is supplied to a metal, the delocalised electrons will be attracted to it, and move towards the positive terminal of the cell
this movement of electrical charge creates an electric current
39
explain why metals are good thermal conductors
-delocalised electrons pass kinetic energy along the metal
-the cations are tightly packed together so when 1 cation vibrates it transfers kinetic energy to the next cation, causing it to vibrate, this continues throughout the metal
40
explain why metals are malleable and ductile
malleable-can be pressed into different shapes
ductile-can be made into wires
-cations are arranged in a regular giant lattice structure
-when a force is applied to a metal, the layers of cations slide over eachother
-the delocalised electrons are free to move, so move with the cations, preventing strong forces of repulsion forming between layers of cations
41
what is the definition of a discrete molecule
an electrically neutral group of two or more atoms held together by chemical bonds
42
how is a dative covalent bond formed
when an empty orbital of one atom overlaps with an orbital containing a non-bonding/lonec pair of electrons in another atom
43
what are the 3 points the valence shell electrons repulsion theory states
-the shape of a molecule/ion is caused by repulsion between electron pairs that surround the central atom
-electron pairs are arranged around the central atom to achieve minimum repulsion
-lone pair-lone pair repulsion is greater than lone pair-bond pair repulsion is greater than bond-bond pair repulsion
44
what rules are used to determine the shape of a molecule/ion
-the shape of a molecule/ion is caused by repulsion between electron pairs that surround the central atom
-electron pairs are arranged around the central atom to achieve minimum repulsion
45
what rule is used to determine bond angles
-lone pair-lone pair repulsion is greater than lone pair-bond pair repulsion is greater than bond-bond pair repulsion
46
how do you determine the shape of a molecule which contains double/triple covalent bonds
treat each double/triple covalent bond the same as a single covalent bond
47
why is a hydrogen molecule non-polar
-its a diatomic molecule
-the 2 hydrogen atoms in the molecule have the same electronegativity
-the distribution of electron density in the bonding electrons is symmetrical
-so the bonds are non-polar making the molecule non-polar
48
is hydrogen chloride polar molecule
why
-yes
-chlorine is more electronegative than hydrogen so the electron density within the bonding pair of electrons is closer to the chlorine atoms
-making the chlorine delta negative and hydrogen delta positive
-so as the only bond is polar the molecule is polar
49
what is the name of a carbon dioxide molecule
is it polar
-linear
-both bonds are polar
-carbon is less electronegative than oxygen
-making carbon delta positive and oxygen delta negative
-but as both bonds are symmetrical (in that the electron densities face away from each other, spreading out the charge, electron densities are symmetrical) the dipoles cancel each other out
-so CO2 is not a polar molecule
50
what is the name of a boron chloride molecule
is it polar
-trigonal planar
-all 3 B-Cl bonds are polar
-chlorine is more electronegative than boron
-making chlorine delta negative and boron delta positive
-the molecule is symmetrical (in that the electron densities face away from each other, spreading out the charge) so the dipoles cancel each other out
-boron chloride is non-polar
51
what is the name of a water molecule
is it polar
-v-shaped
-both O-H bonds are polar
-the dipoles reinforce each other (electron density points towards the oxygen)
-so water is a polar molecule
52
what is the name of a tetrachloromethane molecule
is it polar
-tetrahedral
-all 4 C-Cl bonds are polar
-the molecule is symmetrical (in that the electron densities face away from each other, spreading out the charge) so the dipoles cancel each other out
-its non-polar
53
what is the name of a trichloromethane molecule
is it polar
-tetrahedral
-all 4 bonds are polar, however in the H-C bond carbon is delta negative but in the C-Cl bond chlorine is delta negative (the electron density faces downwards so dipoles don't cancel each other out, electron densities aren't symmetrical)
-the dipoles reinforce each other
-the molecule is polar
54
what is the name of a ammonia molecule
is it polar
-trigonal pyramidal
-all three N-H bonds are polar
-it is not a symmetrical molecule because of the lone pair causing the bond angles to be 107 degrees
-the dipoles reinforce each other (electron density points towards the nitrogen atom, electron densities aren't symmetrical)
-so ammonia is a polar molecule
55
explain how London forces are formed
-a molecule can have polar bonds, but because the electron densities are symmetrically distributed the dipoles cancel each other out, so the molecule is non-polar
-electrons in the charge cloud are constantly moving
-in a moment of time the electron density in the charge cloud may not be symmetricallly distributed
-causing a temporary dipole in molecule A where one side to be delta positive and the other side delta negative
-the delta positive end of molecule A is closer to molecule B, so the electron density of molecule B is pulled to the delta positive side of molecule A, creating an induced dipole in molecule B
-the delta positive end of a temporary dipole in a molecule is attracted to the delta negative end of an induced dipole in another molecule, this force is a London force
56
in what type of molecules do London forces occur
London forces occur between all molecules
polar and non-polar
57
out of London forces, permanent dipole attractions and hydrogen bonding, which is the strongest and weakest force
London forces are the weakest type of intermolecular force
hydrogen bonding is the strongest intermolecular force
58
does the attractive force of London forces change depending on the molecule
-yes
-as the number of electrons in the molecule increases the attractive force increases
59
explain why as the size of a molecule increases the strength of London forces increase
-the larger the molecule the greater the number of electrons
-the more electrons there are in the molecule the greater the fluctuation in electron density
-therefore the larger the instantaneous dipoles and induced dipoles
-so the London force from the delta positive side of an instantaneous dipole will more strongly attract the delta negative side of an induced dipole
60
what 2 factors affect the size of London forces
-the number of electrons within the molecule
-shape and size of the molecule
61
how does the shape and size of a molecule affect the size of London forces
the more points of contact between molecules the greater the London forces
62
explain how permanent dipole-permanent dipole attractions are formed
-occurs between polar molecules
-the delta positive atom in one molecule will attract the delta negative atom in another molecule when the dipoles are aligned favourably
-this attractive force between the 2 molecules is a permanent dipole-permanent dipole attraction
-random movements of molecules mean that the dipoles aren't always aligned to produce a favourable attraction
63
why are London forces more frequently formed than permanent dipole-permanent dipole attractions
-induced dipoles are formed so that the attraction between the delta negative side of the instantaneous molecule and the delta positive side of the induced molecule are favourable,, so London forces always form
-permanent dipole-permanent dipole attractions only occur when when the dipoles are aligned favourably
64
hydrogen bonding occurs when hydrogen is bonded to what
-when hydrogen is bonded to small highly electronegative atoms, so hydrogen is always delta positive and the other atom is delta negative
-oxygen, nitrogen and fluorine
-the hydrogen has to be directly bonded to these atoms
65
between what type of molecules does hydrogen bonding occur
only between polar molecules containing a hydrogen atom bonded to oxygen, fluorine or nitrogen
66
explain how hydrogen bonds are formed when the hydrogen atom is bonded to oxygen
-oxygen atom is delta negative and the hydrogen is delta positive
-on the oxygen atom there is a lone pair of electrons, this lone pair is used to form a partial bond with the delta positive hydrogen on another molecule
-the bond angle between the oxygen on one molecule and the hydrogen on another molecule is 180 degrees
67
explain how hydrogen bonds are formed when the hydrogen atom is bonded to nitrogen
-the nitrogen atom is delta negative and the hydrogen is delta positive
-on the nitrogen atom there is a lone pair of electrons, this lone pair is used to form a partial bond with the delta positive hydrogen on another molecule
-the bond angle between the nitrogen on one molecule and the hydrogen on another molecule is 180 degrees
68
explain how hydrogen bonds are formed when the hydrogen atom is bonded to fluorine
-the fluorine atom is delta negative and the hydrogen is delta positive
-on the fluorine atom there is a lone pair of electrons, this lone pair is used to form a partial bond with the delta positive hydrogen on another molecule
-the bond angle between the fluorine on one molecule and the hydrogen on another molecule is 180 degrees
69
what are 4 examples of giant covalent lattices
diamond
graphite
graphene
silicon(IV) oxide
70
describe the structure of diamond
-allotrope of carbon
-each carbon atom forms 4 sigma bonds
-giant 3-D tetrahedral shape
-all bond angles are 109.5 degrees
71
what are the properties of diamond
-very hard
-very high melting point
-doesn't conduct electricity
72
explain why diamond is very hard and has a very high melting point
-it has a high number of strong C-C bonds
-these covalent bonds need a lot of heat energy to break
73
describe the structure of graphite
-allotrope of carbon
-each carbon atom is covalently bonded to 3 other carbon atoms by sigma bonds
-forming layers of hexagonal rings
-the 4th electron on each carbon atom is in a p-orbital
-p-orbitals from different carbon atoms overlap to produce a cloud of delocalised electrons above and below the plane of rings
74
what are the properties of graphite
a lubricant
good conductor of electricity
high melting and boiling point
75
explain why graphite is a lubricant
there are weak intermolecular forces between layers of graphite
76
explain why graphite is a good conductor of electricity
delocalised electrons are free to move through out the structure when a potential difference is applied
but can only conduct electricity parallel to its layers because the delocalised electrons aren't free to move between layers
77
what is the structure of graphene
a 1 atom thick layer of graphite
each carbon forms sigma bonds with 3 other carbon atoms
forming interlocking hexagonal rings
78
what are 2 examples of solid molecular lattices
ice and iodine
79
describe the structure and forces within iodine
-simple molecular
-iodine exists as diatomic molecules
-the structure is described as "face centred cubic"
-molecules of iodine are held together by London forces
80
what are the properties of molecular solids
-low melting and boiling point
-generally insoluble, but may dissolve if hydrogen bonding is possible (e.g. sucrose) or if the substance reacts with water (e.g. Cl2)
-don't conduct electricity
-covalent bonding
-intermolecular forces of attraction are present
81
explain why molecular solids have a low melting point
molecules are held together by intermolecular forces, which require little energy to break
(the covalent bonds don't need to break)
82
what structure does metallic bonding have
giant lattice
83
what structure does ionic bonding have
giant lattice
84
what structure does covalent bonding have
either giant lattice, molecular, or macromolecular
85
what are the properties of giant metallic lattices
-metallic bonding
-no intermolecular forces of attraction present
-high melting and boiling points
-good conductors of electricity when solid and molten
-insoluble unless the metal reacts with water
86
what are the properties of giant ionic lattices
-ionic bonding
-no intermolecular forces of attraction present
-high melting and boiling points
-don't conduct electricity when solid, but do when molten/aqueous
-soluble
87
what are the properties of giant ionic lattices
-covalent bonding
-no intermolecular forces of attraction present
-high melting and boiling points
-don't conduct electricity when solid or molten
-insoluble
88
for unbranched alkanes what is the relationship between relative molecular mass and boiling temperature
as relative molecular mass increases boiling temperature increases
89
what intermolecular forces are present in alkanes
London forces only
90
explain why as the relative molecular mass of alkanes increases, boiling temperature also increases
-as molecular mass increases, the number of electrons per molecule increases, so the strength of the instantaneous and in induced dipoles increase
-as the length of the carbon chain increases, teh number of points of contact between adjacent chainsmolecules increase, because instantaneous dipole-induced dipole forces exist between each point of contact between molecules, the greater the overall London force of attraction
-both of these points mean more energy is needed to break London forces
91
explain why even if 2 different alkanes have the same relative molecular mass, the non-branched alkane will always have a high melting point
-the more ethyl groups in the molecule, the fewer points of sontact between adjacent molecules (they don't pack together as well)
-so there's a lower overall intermolecular force of attraction between molecules
-so less energy required to break them
92
explain why methanol and ethane have different boiling temperatures, even though both molecules have to same number of electrons in them
-both molecules have similar chain lengths and the molecules contain the same number of electrons
-both molecules have London forces
-but methanol also has hydrogen bonding because of the -OH group
-more energy is required to break hydrogen bonds compared to London forces
-
93
is the statement "the predominant bonding in alcohols is hydrogen bonding" true or false
-for the first few alcohols this can be true
-but as the carbon chain increases in length the percentage contribution of hydrogen bonding to the enthalpy required to vaporise the alcohol (completely separate the molecules, therefore breaking all intermolecular forces) decreases
-so the statement isn't true for all alcohols
94
what is the trend in boiling point for hydrogen halides
-from HF to HCl boiling point decreases
-HF has a much greater boiling point because it has hydrogen bonding and London forces
-from HCl to HBr to HI there is a steady increase in boiling point
-because the number of electrons in the molecule also increase, which means greater instantaneous dipole-induced dipole attractions, so London forces increase
95
what are the 2 anomalous properties of water
-relatively high melting point and boiling point, that would be expected for a molecule with such few electrons
-the density of ice at 0°C is less than that of water at°C
96
explain why water has a high melting and boiling point
-hydrogen bonding requires a lot of heat energy to break
-water molecules form 2 hydrogen bonds per molecule
-which in comparison to NH3 and HF which make 1 hydrogen bond per molecule
97
explain why the density of ice at 0°C is less than that of water at 0°C
-when water molecules are in liquid state the molecules are held together much closer together than in solid state
-when water starts to freeze,, water molecules move away from eachother and form rings of 6 molecules held together by hydrogen bonds
-this strucutre means there are large areas of open space
-so there are less water molecules in a given area, decreasing density in solid state
98
what are 2 requirements for a solute to dissolve
-solute particles must be separated from each other, then surrounded by solvent particles
-the forces of attraction between the solute and solvent particles must be strong enough to overcome the solvent-solvent forces and solute-solute forces
99
how does an ionic solid dissolve in water
-the ions in the lattice are hydrated
-the delta negative oxygen on the water molecule attracts the cations, removing them from the lattice
-the cations become surrounded by the delta negative oxygen atoms in the water molecules
-the delta positive hydrogen on hte water molecule attracts the anion, removing it from the lattice
-the anion becomes surrounded by hydrogen atoms on the water molecules
-the interaction between ions in the lattice and water molecules is an ion-dipole interaction
-this hydration process releases hydration energy
100
explain what effects the solubility of alcohols
-as the hydrocarbon chain of alcohols increase in length, the solubility of alcohols decreases
-because London forces are more predominant in the bonding
-alcohols are only soluble because they form hydrogen bonds with water molecules
101
why do halogenoalkanes not dissolve in water
halogenoalkanes cannot form hydrogen bonds with water
102
if a solute contains the same intermolecular forces as a solvent, the solute is likely to dissolve in the solvent
true or false
true
103
like dissolves like
what does this mean
-if a solute contains the same intermolecular forces as a solvent, the solute is likely to dissolve in the solvent
-if they are both polar or both non-polar the solute is likely to dissolve in the solvent
104
what is a dimer
a molecule consisting of 2 identical molecules bonded together by a dative covalent bond
105
what is an example of a dimer
aluminium chloride Al2Cl6
2 AlCl3 molecules join together by a lone pair of electrons on one of the chloride ions forming a dative covalent bond
106
what are 3 examples of molecules formed by dative covalent bonding
hydroxonium ion H3O+
ammonium ion NH4+
aluminium chloride Al2Cl6
107
Ammonia and boron trifluoride react to form a compound NH3BF3 which contains a dative
covalent bond. Each of the molecules, NH3 and BF3, has a different feature of its electronic
structure that allows this to happen. Use these two different features to explain how a dative
covalent bond is formed.
(2 marks)
1 - Nitrogen donates is lone pair of electrons
2 - to the boron atom which has 6 electrons in its outer shell, so is electron deficient and can accept 2 more electrons
108
Explain why both water and carbon dioxide molecules have polar bonds but only
water is a polar molecule.
(4 marks)
1 - oxygen is more electronegative than carbon and hydrogen
2 - creating polar bonds where oxygen is delta negative and carbon and hydrogen are delta positive
3 - due to the linear shape the CO2 molecule is symmetrical so the dipoles cancel each other out
4 - H2O is v-shaped because of the lone pairs of electrons on the oxygen atom so the dipoles don't cancel each other out
109
explain the shape of a PCl3 molecule
bond angels aren't required (3 marks)
trigonal pyramidal
3 bonding pairs and 1 lone pair around the central phosphorous atom
electron pairs repel to positions of maximum separation and minimal repulsion
110
Explain why phosphorus forms PCl5 but nitrogen does not form NCl5
(2 marks)
phosphorous can expand its octet to hold 10 electrons because it has a 3d orbital available
nitrogen can only hold 8 electrons in its outer shell
