1
Q
Ionic bonding
A
Involves transfer of one or more electrons from outer shell of one atom to outer shell of another atom
- transfer of these electrons results in formation of positive and negative ions
2
Q
Cations
A
formed when metals lose valence electrons
- hence, have more protons than electrons in the atom overall
- this gives them a positive charge overall, forming a positive ion
3
Q
Anions
A
formed when non-metals gain electrons
- hence, have less protons than electrons in the atom overall
- this gives them a negative charge overall, forming a negative ion
4
Q
Ionic bond
A
Electrostatic attraction between oppositely charged ions
- formed by elements w/ an electronegativity difference of approx. 1.8 units and greater
- ionic bonding is non-directional, force of attraction occurs in all directions around the individual atoms
5
Q
Ionic compounds
A
Composed of a metal and non-metal element (electrically neutral)
- formation of positive and negative ions results in atoms achieving full outer shells of electrons
- oppositely charged ions are attracted to each other, resulting in formation of an ionic bond
- in accordance with the octet rule
Exception to metal and non-metal rule:
Ammonium chloride- has both types of bonding
- ionic bonding between the ions
- covalent bonding in the ammonium ion
6
Q
Octet rule
A
States that atoms are more stable with the electron configuration of a noble gas
7
Q
Structure of ionic compounds
A
Under normal conditions, ionic compounds are usually solids w/ lattice structures
- ions that make up an ionic compound are arranged in a regular crystalline structure, lattice structure
8
Q
Polyatomic ions
A
Ions that consist of more than one atom
- ions that make up the polyatomic ions are bonded by covalent bonds
- but bonding between a polyatomic ion and another ion is ionic
9
Q
Formula of an ionic compound
A
- ratio of the ions that make up the compound, so , overall, the charges cancel out
- referred to as the ‘formula unit’, combination of ions is repeated throughout the lattice structure
10
Q
Physical properties of ionic compounds
A
- MP and BP
- solubility
- electrical conductivity
- brittleness
11
Q
Melting point and boiling point
A
- an ionic compound is held together by electrostatic attractions between oppositely charged ions
- strong attraction between ions results in ionic compounds, relatively high MP and BP
- hence, relatively large amounts of energy are needed to break these forces of attraction between the ions
- hence, ionic compounds are solids under standard conditions
- MP of an ionic compound depends on ionic charge and ionic radius of its component ions
- greater the charge on ion and smaller its ionic radius, greater attraction between oppositely charged ions and higher the MP
12
Q
Volatility
A
How easily a substance evaporates
- ionic compounds have very low volatility because of strong forces of attraction between ions in lattice structure
13
Q
Solubility
A
Not all ionic compounds dissolve in water
- depends on whether forces of attraction between water molecules and ions in lattice are strong enough to overcome attractions between ions
- water molecules are polar due to difference in electronegativity between oxygen and hydrogen
- at surface of ionic lattice, where it’s in contact with water, positive and negative dipoles of water molecule are attracted to oppositely charged ions in lattice structure
- these ions break off from lattice and are surrounded by water molecules (hydration)
- when this happens, the solid dissolves
14
Q
Non-polar solvents
A
Non-polar solvents can’t disrupt lattice structure, hence solubility of ionic substances eg. hexane and propanone is limited
15
Q
Solubility and solvents
A
Polar substances are soluble in polar solvents
Non-polar substances are soluble in non-polar solvents
16
Q
Electrical conductivity
A
Depends on presence of mobile ions
- ionic compounds don’t conduct electricity when solid because ions are held in fixed positions in lattice structure
- when an ionic compound is molten or dissolved, ions are free to move and carry an electric current
17
Q
Brittleness
A
Ionic compounds tend to shatter when a force is applied- are said to be brittle
- fracture across a plane when layers of ions become incorrectly aligned
- occurs because movement of ions within lattice when force is applied places ions of same charge next to each other
- forces of repulsion between ions of same charge cause lattice structure to split and fracture
18
Q
General properties of ionic compounds
A
- generally highly soluble in water
- good electrical conductivity only when molten or dissolved
- solids at room temperature
- high MP and BP
- brittle
19
Q
Why are ionic compounds soluble in water?
A
Water molecules are attracted to oppositely charged ions causing them to dissolve
20
Q
Why are ionic compounds good electrical conductors?
A
When molten/in solution, ions are free to move about
- when solid, ions are held in fixed positions in lattice structure
21
Q
Why are ionic compounds solids at room temperature?
A
Ions are held in fixed positions by strong electrostatic attractions in lattice structure
22
Q
Why do ionic compounds have high MP and BP?
A
Ions are attracted to each other by strong electrostatic attractions
- large amounts of energy are needed to separate them
23
Q
Why are ionic compounds brittle?
A
Repulsion between ions of same charge causes lattice structure to split and fracture
24
Q
Covalent bonding
A
Occurs between non-metal elements and results in formation of molecules and giant covalent structures
- occurs between elements with a difference in electronegativity of fewer than 1.8 units
25
Hydrogen gas, H2
- exists as a diatomic molecule w/ a single covalent bond between two hydrogen atoms
- atoms are held together by electrostatic attraction that exists between two nuclei and shared pair of electrons
- a hydrogen atom needs two electrons in its valence shell to achieve noble gas structure
- two hydrogen atoms in a hydrogen molecule share pair of bonding electrons, so each atom now has two electrons in its valence shell
- electrons are shared
26
Covalent bond
Electrostatic attraction between a shared pair of electrons and the positively charged nuclei
27
Halogens as diatomic molecules
- all form diatomic molecules
- atoms are bonded via single covalent bond
- atomic radius increases down a group, so length of bond between atoms also increases
28
Different covalent bonds
Single: 2 shared bonding electrons
Double: 4 shared bonding electrons
Triple: 6 shared bonding electrons
29
Bond length
Distance between the nuclei of the bonded atoms
30
Bond strength
Amount of energy needed to break the bond between the atoms
- described in terms of bond enthalpy
- can affect reactivity of a compound
31
Long bonds vs. short bonds
- longer bonds are weaker than shorter bonds due to increased distance between nuclei and shared pairs of bonding electrons
- results in a weaker electrostatic attraction between atoms
32
Multiple bonds
- involve more shared electrons between atoms
- electrostatic attraction between bonded nuclei is greater
- this greater attraction brings nuclei closer together, so multiple bonds are shorter and stronger than single bonds
33
Bond order
Number of bonds between a pair of atoms
- single bonds = bond order of 1
- double bonds = bond order of 2
- triple bonds = bond order of 3
The higher the bond order, the greater the strength of the bond
34
Calculating bond order
sum of individual bond orders/ number of bonding groups
35
Electronegativity and bonding
Measure of attraction of an atom for a shared pair of bonding electrons
- difference in electronegativity determines type of bonding that takes place between atoms
- large difference in electronegativity = ionic bond
- small difference in electronegativity = covalent bond
36
Electronegativity values and bonding
Ionic = greater than or equal to 1.8
Polar covalent = 0.5 - 1.7
Non-polar covalent = 0.0 - 0.4
The greater the electronegativity difference between the atoms, greater the polarity of the bond, greater its ionic character
37
Non-polar covalent bonds
Occur between atoms w/ a small difference in electronegativity (0.0-0.4 units)
- in these molecules, bonding electrons are equally shared
- present in bonds in diatomic molecules composed of the same atom e.g. oxygen, chlorine gas, nitrogen gas
38
Polar covalent bonds
Occurs between atoms w/ a electronegativity difference between 0.5-1.7 units
- in these molecules, bonding electrons aren't shared equally between atoms
- bond w/ greatest polarity = H-F bond (biggest difference in electronegativity between the two atoms)
- H-F bond is still a covalent bond, despite large difference in electronegativity
- bond w/ least polarity = H-I bond (smallest difference in electronegativity between the two atoms)
39
Polarity of hydrogen halides
Polarity of H-X bond decreases as you descend group 7
40
Bond polarity symbols
Bond polarity= results from the difference in electronegativities of the bonded atoms
Bond polarity is designated by the SIGMA+ and SIGMA- signs placed on the molecules
- refers to partial charges called dipoles
- SIGMA- is assigned to the more electronegative element in the bond
41
Factors to consider for polarity of a molecule
1. Presence of polar bonds within the molecule
| 2. Molecular geometry of the molecule
42
Non-polar molecules
- Molecules is symmetrical, despite containing polar bonds
- because these molecules are symmetrical, so, bond polarities cancel each other out
- molecule has no net dipole movement
43
Polar molecules
- have a net dipole movement
- presence of polar bonds within molecules, together w/ molecular geometry, means bond polarities don't cancel out
NB/presence of a net dipole movement determines type of intermolecular force that exists between molecules
44
Octet rule
States that the most stable arrangement for an atom is to have 8 electrons in its outermost energy level w/ electron configuration of a noble gas
- it's the tendency of atoms to gain a valence shell w/ a total of 8 electrons
45
Exceptions to the octet rule
- Hydrogen is stable w/ only 2 electrons in its outer shell
- Atoms eg. boron, beryllium and aluminium (in compounds) are stable w/ fewer than 8 electrons in their outer shell
- atoms in period 3 and higher, e.g. sulphur, can form expanded octets w/ up to 12 electrons in their valence shell
46
Incomplete octets
Atoms with less than 8 electrons in their outer shells- said to have incomplete octets or be electron-deficient
47
Electron-deficient molecules
Do not have a full outer shell of electrons
- atoms in molecules that are electron-deficient are able to form coordinate covalent bonds
- eg. Aluminium chloride
48
Example of an incomplete octet
Boron trifluoride (BF3)
- a boron atom only has 3 electrons in its outer energy level- no possibility of it reaching a noble gas structure w/ 8 electrons simply by sharing electrons
- hence, in BF3, boron atom has formed max. no. of bonds that it can under the circumstances- and is stable
49
Forming a molecule
An atom will tend to make as many covalent bonds as possible
50
Example of an electron-deficient molecule
Atoms in molecules that are electron-deficient are able to form coordinate covalent bonds
- example of a molecule that has coordinate covalent bonds is AlCl3 (aluminium chloride)
- at high temp. = aluminium chloride exists as molecules w/ formula AlCl3 (molecule is electron-deficient, still needs 2 electrons to complete octet in outer shell of aluminium atom)
- at lower temp. = 2 molecules of AlCl3 combine to form a dimer w/ formula Al2Cl6
- AlCl3 molecules are able to combine because lone pairs of electrons on 2 of the chlorine atoms form coordinate covalent bonds w/ aluminium atoms
NB/ coordinate covalent bond is represented by use of an arrow
- dots and crosses represent valence electrons of aluminium and chlorine
51
Expanded octets
Elements in period 3 and beyond can have expanded octets- accommodate more than 8 electrons in their valence shell
- molecules w/ central atoms from elements in period 3 can accommodate up to 18 electrons in valence shell
52
Lewis structures
Represent bonding in a molecule, they show both bonding electrons and non-bonding electrons
- also called electron dot diagrams
NB/ remember to draw the lone pairs of electrons so each atom has an octet of electrons
- remember to include square brackets and charge when drawing structure for an ion
53
Coordinate covalent bond
Covalent bond = formed between atoms that both contribute electrons to the bond
A coordinate covalent bond = differs from a 'regular' covalent bond
- both bonding electrons come from one atom
NB/ once formed, a coordinate covalent bond is identical to a 'regular' covalent bond
54
Resonance structures
is one of two or more alternative Lewis structures for a molecule or ion that can't be described fully w/ one Lewis structure alone
- no. of possible resonance structures is equal to no. of different positions for multiple bond
- occurs when molecules contain multiple bonds (double or triple), there is more than one possible Lewis structure that can be drawn
- but none of the resonance structures represents the actual structure of the ion
- the actual structure = resonance hybrid structure
55
Resonance hybrid structure
Intermediate between the forms of resonance structures
| - all bonds are identical, and are intermediate in strength and length between a single and double bond
56
Why do covalent molecules have defined shapes?
Electrons in a covalent bond are located in a specific position within a molecule
- hence, covalent molecules have a defined shape, determined by the orientation of these 'fixed' or 'localised' bonds in space
57
Delocalised electrons
involves e- that are shared by/between all atoms in a molecule or ion as opposed to being localised between a pair of atoms
- give greater stability to a molecule or polyatomic ion
- exist in any molecule where there is more than one possible Lewis structure- hence, more than one position for a multiple bond
58
VSEPR theory
Valence shell electron pair repulsion theory
- theory that electron pairs in molecules repel each other and orientate themselves as far away from each other as possible
- molecule adopts shape that minimises repulsion between bonding and non-bonding electrons (lone pairs)
- enables us to predict shapes of molecules
- theory proposes that molecular shape adopted is that where strengths of repulsion between electron pairs in valence shell of central atom are minimised
59
Bonding domains
Bonding electrons in single, double or triple bonds are known as bonding domains
60
Non-bonding domains
Non-bonding electrons are called non-bonding domains
61
Electron domains
Both bonding domains and non-bonding domains are collectively known as electron domains
- single bonds, double bonds and triple bonds each count as one electron domain, as do non-bonding pairs of electrons
62
Molecular geometry
Shape of a molecule
- deduced by looking at Lewis structure for molecule
- count no. of electron domains around central atom
63
Electron domain geometry
Total no. of electron domains (both bonding and non-bonding) around the central atom
- for some atoms, molecular geometry is different to electron domain geometry
64
Electron domain geometry vs. molecular geometry
Electron domain geometry:
- total no. of electron domains around the central atom
Molecular geometry:
- shape of a molecule
- no. of electron domains around the central atom
- ALSO takes into account the extra repulsion between bonding and non-bonding domains
65
Linear geometry
- molecules w/ two bonding domains around the central atom- linear molecular geometry
- bond angle in a linear molecule = 180 degrees
- electron domain geometry is also linear
66
Trigonal planar geometry
- molecules/ polyatomic ions w/ 3 bonding domains
- bonding domains are at 120 degrees to each other in the same plane
- 3 bonding domains around the central atom
- electron domain geometry is the same as molecular geometry
2 bonding domains + 1 non-bonding = shape is bent or V-shaped
67
Tetrahedral geometry
- molecules w/ 4 bonding domains around central atom
- bond angle = 109.5 degrees
- electron domain geometry is same as molecular geometry- because, all 4 electron domains are bonding domains
68
Non-bonding electrons
Cause slightly more repulsion than bonding pairs of electrons
69
Order of repulsion between non-bonding domains (NBD) and bonding domains (BD)
NBD-NBD > NBD-BD > BD-BD
Greatest repulsion = between non-bonding domains
Least repulsion = between bonding domains
70
Purpose of Lewis structures
show all the valence electrons in a covalently bonded species
71
Giant covalent structures
Substances that don't exist as discrete molecules
| - instead, exist as many atoms bonded together in a regular arrangement
72
Diamond
- each carbon atom in diamond is covalently bonded to 4 other carbon atoms in a tetrahedral arrangement w/ a bond angle of 109.5 degrees
- strong covalent bonds between carbon atoms mean that it's a very hard substance w/ a high MP and BP
- precious gemstone, used in jewellery
- poor electrical conductor- no delocalised electrons within its structure
73
Silicon
- group 14 element
- each silicon atom is bonded to 4 other silicon atoms in a tetrahedral arrangement, bond angle of 109.5 degrees
- bonds between silicon atoms are strong covalent bonds- electrons are held in fixed positions, can't move within the structure- silicon is a poor conductor of electricity at low temp.
- electrical conductivity can be improved by doping (addition of small amounts of elements eg. phosphorus or boron to pure silicon)- process used in production of photovoltaic cells
74
Silicon dioxide
- each silicon atom is bonded by strong covalent bonds to 4 oxygen atoms in a tetrahedral arrangement w/ bond angle of 109.5 degrees
- Si - O - Si bonds have a bent arrangement caused by lone pairs of electrons on oxygen atoms
- strong covalent bonds between atoms are responsible for properties of silicon dioxide
- a very hard substance, a poor conductor of electricity, high MP and BP
75
Allotropes of carbon
Carbon exists as 3 allotropes:
- diamond
- graphite
- fullerenes
These allotropes have different properties eg. electrical conductivity and hardness
- because of different bonding within structures
76
Allotropes
Different forms of the same element in the same physical state
77
Graphite
- each carbon atom is covalently bonded to 3 other carbon atoms
- has a layered structure consisting of carbon atoms arranged in fused hexagonal rings
- a graphite crystal is composed of many such planar sheets of hexagonally arranged carbon atoms, stacked on top of one another
- layers are held together by relatively weak London dispersion forces
- each carbon atom has an electron which becomes delocalised across the plane
- delocalised electrons explains why graphite can conduct electricity along the plane of the crystal when a voltage is applied
- graphite is soft, hence, used as 'lead' in pencils (layers are able to slide over one another due to weak intermolecular forces between layers, making graphite soft and slippery)
78
Properties and uses of diamond
1. colourless transparent crystals that sparkle in light - used in jewellery and ornamental objects
2. Hardest natural substance - used in drill bits and diamond saws
3. Doesn't conduct electricity
79
Properties and uses of graphite
1. Dark grey shiny solid
2. Soft- used in pencils and as a lubricant
3. Good electrical conductor- used as electrodes and in electric motors
80
C60 fullerene
- a simple molecular substance even though it contains so many bonded carbon atoms
- structure is made of carbon atoms bonded together in 20 hexagons (6 carbon rings) and 12 pentagons (5 carbon rings), truncated icosahedron
- each carbon atoms is covalently bonded to 3 other carbon atoms
- presence of delocalised electrons within structure means that it does conduct electricity, poorer electrical conductor than graphite
81
Graphene
- single- layered material made of individual sheets of graphite
- composed of single layers of graphite w/ a bond angle of 120 degrees between carbon atoms in a trigonal planar arrangement
- high tensile strength
- high electrical and thermal conductivity
- useful in energy and biomedical industries
- behaves as a semi-metal, it's suitable for electronic devices
NB/ of the different allotropes of carbon, graphene is the most chemically reactive
- because of the reactive edges of the structure, where there are carbon atoms with unoccupied bonds
82
Intramolecular forces
Atoms within a molecule are bonded by covalent bonds- bonds that are within the molecule
83
Intermolecular forces
Forces that exist between molecules
| - responsible for physical properties eg. BP, MP, solubility of a substance
84
3 types of intermolecular forces
- London dispersion forces
- dipole-dipole forces
- hydrogen bonding
NB/ London dispersion, dipole- induced dipole and dipole-dipole forces are collectively known as van der Waal's forces
85
London dispersion forces
refers to instantaneous induced induced dipole-induced dipole forces that exist between any atoms or groups of atoms
- should be used for non-polar entities
- weakest type of intermolecular force
NB/ have significant effects on physical properties of molecular covalent substances
86
Induced dipoles
A molecule with a temporary dipole can induce a dipole in a neighbouring molecule- an induced dipole
- temporary dipole in one atom induces a dipole in adjacent atoms
87
Temporary dipoles
Temporary dipole- caused by changes in electron density within an atom or molecule
- at a certain moment in time, electrons may be concentrated on one side of an atom or molecule- gives this side a slight -ve charge, and opp. side a slight +ve charge
- these opposite charges are known as dipoles
88
The strength of London dispersion forces depends on:
- ease with which the electrons in an atom or molecule form a temporary or induced dipole (their polarizability)
- surface area of the molecule
89
Strength of London dispersion forces and polarisability
In general, as molar mass of a molecule increases, so does its polarizability
- results in a larger temporary or induced dipole being formed within the molecule, leads to stronger London dispersion forces between the molecules
- as London forces between molecules increase, BP of substance also increases
90
Halogens and BP
- as we move down group, molar mass of molecules increases
- results in an increase in strength of London dispersion forces between molecules
- increase in BP
91
Strength of London dispersion forces and surface area
- molecules that have a greater area between them have stronger London dispersion forces and higher BP
92
Dipole-dipole forces
- only exist between polar molecules that have a permanent dipole eg. HCl
93
Hydrogen chloride as a permanent dipole
- permanent dipole arises due to difference in electronegativity between hydrogen and chlorine
- hydrogen, w/ lower electronegativity value, has a positive dipole while chlorine, w/ higher electronegativity value has a negative dipole
- force of attraction is between -ve dipole of chlorine atom of one molecule and +ve dipole of hydrogen atom on another
94
Hydrogen bonding
Occurs between molecules that have an electronegative nitrogen, oxygen or fluorine atom directly bonded to a hydrogen atom (H-N, H-O, H-F)
- stronger type of dipole-dipole attraction, responsible for high BP of water
95
Hydrogen bonding in water
- Hydrogen bonding is responsible for high BP of water
| - hydrogen bonds occur between lone pairs of electrons on oxygen atom and hydrogen atom on a nearby water molecule
96
Relative strength of intermolecular forces
Hydrogen bonding > dipole-dipole forces > London dispersion forces
97
Solubility
- water is the universal solvent- able to dissolve many different substances
- many ionic compounds are soluble in water due to its polar nature
- certain covalent substances are also soluble, able to form hydrogen bonds w/ water molecules
- polar substances are soluble in polar solvents
- non-polar substances are soluble in non-polar solvents
98
Metallic structure
In metals:
- valence electrons are free to move throughout the metallic structure
- electrons are delocalised ('sea of delocalised electrons')
- metal atoms are ionised, forming positive ions that are arranged in a lattice structure
- bonding in metals = metallic bond
99
Properties of metals
- good conductors of heat and electricity
- ductile (can be made into wires)
- malleable (can be bent into shape)
- shiny when polished
100
Metallic bond
The electrostatic attraction between lattice of positive metal ions and 'sea' of delocalised electrons
- non-directional- force of attraction occurs in all directions between positive ions and delocalised electrons within lattice structure
101
Metals' malleability and ductileness
- if sufficient force is applied to metal, one layer of metal ions can slide over each other without disrupting metallic bond- bond remains intact
- explains why metals are malleable and ductile
- bending a metal into shape doesn't break the metallic bond, instead, layers of metal ions slide over each other without disrupting metallic bond
102
Why can metals conduct electricity?
- when a potential difference is applied across a metal, a direction is imposed on movement of delocalised electrons
- they're repelled from negative electrode and move to positive electrode
- this orderly flow of delocalised electrons in a given direction constitutes flow of an electric current
- delocalised electrons can also conduct heat,
103
Why can metals conduct heat?
- delocalised electrons can conduct heat
- electrons move through the metal, carrying kinetic energy (in form of vibrations) from hotter part of metal to colder part
- these delocalised electrons are also responsible for shininess of metals, because they reflect wavelengths of visible light
104
The strength of the metallic bond depends on:
- charge on the metal ion
- ionic radius of the metal ion
The MP can be taken as approx. measure of strength of metallic bond
- stronger the bond, higher the MP
105
Alloys
Homogeneous mixtures composed of two or more metals or a metal and a non-metal
- have different properties to the metal they're made from
- they're often stronger and more resistant to corrosion than their component elements
106
How are alloys made?
- alloys are produced by adding one or more metal or non-metal elements to another metal in its molten state so different atoms can mix
- as mixture solidifies, atoms of different metals or non-metals are scattered through the lattice structure
107
Why are metal malleable?
Due to layers in metallic lattice being able to slide over each other without breaking the metallic bond
108
Why are alloys harder than component metals?
Addition of different sized atoms in alloy means that layers can't slide over each other as easily
- results in alloys being harder than its component metals
- because of different packing of atoms in metallic lattice, alloys have properties that are different from their component elements
- alloys are often stronger, more chemically stronger and more resistant to corrosion
109
Properties of Brass
Composition: 70% copper, 30% zinc
- harder than pure copper
- used to make musical instruments
110
Properties of Bronze
Composition: 90% copper, 10% tin
- harder than pure copper
- used to make statues
111
Properties of Mild steel
Composition: 99.7% iron, 0.3% carbon
| - stronger and harder than pure iron
112
Properties of Stainless steel
Composition: 74% iron, 18% chromium, 8% nickel
- increased resistance to corrosion
- used to make cutlery
113
Properties of Solder
Lead solder: 50% tin, 50% lead
Lead-free solder: tin w/ various metals
- lower MP than either tin or lead
- used in electrical circuit boards
114
Resonance structures
exist for molecules that have more than one possible Lewis structure
115
Resonance hybrid structures
- a hybrid of all the possible resonance structures of that compound
- has bonds of intermediate length and strength
- each resonance structure contributes to the hybrid structure depending on its energy
- resonance structure w/ lowest energy contributes the most to the hybrid structure
NB/ resonance hybrid is always more stable than any of the possible resonance structures
Dashed lines indicate that each bond is identical in terms of length and strength
116
Equivalent resonance structures
- when all resonance structures are of equal energy
| - thus, make equal contributions to resonance hybrid
117
Resonance energy
difference in energy between the hybrid structure and that of the most stable resonance structure
118
Equivalent Lewis structures
each resonance structure contains same no. of single and double bonds as another
119
Non-equivalent Lewis structure
each resonance structure contains different no. of multiple bonds
120
Formal charge
used to determine which Lewis structure is the preferred one when there is more than one possibility
- the charge an atom would have if all atoms in the molecule had the same electronegativity
- assumes that a bond is 100% covalent, w/ no ionic character
121
Equation used to calculate formal charge
FC = (no. of valence e-) - 1/2(no. of bonding e-) - (no. of non-bonding e-)
Use:
V= valence e-
B= bonding e-
L= non-bonding e-
FC = V - 1/2B - L
NB/ preferred Lewis structure is the one w/ the formal charge closes to 0 on each atom
- for neutral compounds, the sum of the formal charges must be equal to 0 eg. CO2
- sum of formal charges must = charge on the ion
122
What happens when there's more than one Lewis structure w/ formal charge closest to 0?
Assign -ve formal charge to the more electronegative atom
123
Sum of formal charges
- neutral molecule
- polyatomic ion
Neutral molecule: sum of formal charges must = 0
Polyatomic ion: sum of formal charges must = overall charge on the ion
124
Expanded octets
- elements in period 3 and beyond can accommodate more than 8 e- in their valence shells
- due to availability of d orbitals which can be used for bonding
- expanded octets results in formation of molecules w/ 5 and 6 electron domains around the central atom
125
Electron domains
repel each other to be as far apart as possible
126
Electron domain geometries for 5 and 6 electron domains
5 electron domains: trigonal bipyramidal
| 6 electron domains: octahedral
127
Electron domain geometry
total no. of electron domains (both bonding and non-bonding) around the central atom
128
Molecular geometry
- total no. of electron domains (both bonding and non-bonding) around the central atom
- takes into account the extra repulsion between bonding and non-bonding e-
NB/ for some molecules, molecular geometry is different to e- domain geometry
129
Molecules w/ 6 electron domains
Electron domain geometry: Octahedral
Molecular geometry:
1. Octahedral:
- 6 bonding domain, 0 non-bonding domains
- bond angles are 90 degrees and 180 degrees
2. Square pyramidal:
- 5 bonding domains and 1 non-bonding domain
- bond angle of less than 90 degrees
- reduced bond angle is due to extra repulsion between bonding and non-bonding domains around central atom
3. Square planar:
- 4 bonding domains and 2 non-bonding domains
- bond angle of 90 degrees
130
Molecules w/ 5 electron domains
Electron geometry: Trigonal bipyramidal
Molecular geometry:
1. Trigonal bipyramidal:
- 5 bonding domains, 0 non-bonding
2. See-saw:
- 4 bonding domains, 1 non-bonding domain
- bond angles are 90 degrees and less than 120 degrees
3. T-shaped:
- 3 bonding domains and 2 non-bonding domains
- bond angle is less than 90 degrees
4. Linear:
- 2 bonding domains and 3 non-bonding domains
- bond angle is 180 degrees
NB/ non-bonding domains occupy equatorial positions to minimise effect of repulsive forces between non-bonding and bonding domains
131
Covalent bonding
- single
- double
Covalent bond: formed by the sharing of valence e- between atoms
Single: involves sharing of 2 e-
Double: involves sharing of 4 e-
Triple bond: involves sharing of 6 e-
132
Delocalised pi electrons
Electrons that are shared between more than 2 nuclei
- they are present in molecules and polyatomic ions for which there's more than one possible Lewis structure
- originate from overlap of pi bonds in a molecule or ion
133
Single covalent bond
composed of 1 sigma bond
134
Double covalent bond
composed of 1 sigma and 1 pi bond
135
Triple covalent bond
composed of 1 sigma and 2 pi bonds
136
Bond strengths between pi and sigma bonds
- pi bond isn't as strong as a sigma bond
- extra strength of sigma bonds comes from the greater overlap of atomic orbitals in the bond
- in a pi bond, atomic orbitals can't overlap as much, results in a weaker bond
137
Properties of the sigma bond
- formed by head-on axial overlap of atomic orbitals
- electron density is concentrated between the nuclei of the bonding atoms
- a sigma bond can be present on its own, or in combinate w/ pi bonds
- there's free rotation of atoms around a sigma bond
Formed by:
- head-on overlap of s and p orbitals
- head-on overlap of two p orbitals
138
Properties of pi bond
- formed by sideways overlap of atomic orbitals
- electron density is concentrated above and below the plane of the nuclei of the bonding atoms
- it isn't formed independently, it's only formed in the presence of a sigma bond
- there is no free rotation around a pi bond because it would involve breaking the bond
139
Ozone
- also called trioxygen
- composed of 3 oxygen atoms bonded together w/ a V- shaped or bent molecular geometry
- central oxygen atom has 3 electron domains- 2 bonding, 1 non-bonding
- extra repulsion from non-bonding pair of e- gives a bond angle of approx. 116 degrees
- ozone molecule has delocalised pi e-
- these e- are spread over both bonding positions in the molecule- results in both bonds being identical in both bond length and bond strength
- these intermediate bonds are represented by the dashed lines, resonance hybrid structure
- ozone molecule is polar (permanent dipole) due to different formal charges on each oxygen atom
140
Calculation of wavelength of light required to dissociate oxygen and ozone
1. Divide oxygen's bond enthalpy by Avogadro
2. Use wavelength = (hc)/E
3. Convert to nm by multiplying by 10^9
141
Chain reaction
- CFC's are highly stable compounds
- means that CFC molecules released into lower atmosphere could remain intact and reach upper atmosphere
- when exposed to UV radiation, compounds such as trichlorofluromethane decompose to produce chlorine free radicals
- chlorine free radicals act as catalysts- break down ozone to form diatomic oxygen molecules
- free oxygen atoms are also present due to action of powerful UV radiation on molecular oxygen
- they're part of the normal generation of ozone that takes place in upper atmosphere
142
Ozone depletion
- other compounds that cause ozone depletion: nitrogen oxides eg. nitrogen monoxide and nitrogen dioxide
- both molecules have an unpaired e- and are free radicals
- nitrogen oxides are produced in internal combustion engines- high temp. allow atmospheric nitrogen and oxygen to react
- in this reaction nitrogen oxide acts as a catalyst in the reaction
143
Carbon in the ground state
- each carbon atom forms 4 covalent bonds in many compounds
- its electron config. shows that it only has 2 half-filled atomic orbitals available for bonding (2pz orbital is unoccupied)
- in its excited state, an e- has been promoted from a doubly occupied 2s orbital to empty 2pz orbital- excitation
- energy required to promote an e- to a higher energy orbital is made up for when carbon atom forms 4 new bonds
Electron config. of excited carbon atom:
- has 4 half-filled orbitals (one 2s and three 2p orbitals)
- but, if carbon were to form 4 covalent bonds w/ this electron config., bonds wouldn't be equal
- when carbon forms 4 bonds, all bonds are equal
144
Hydbridisation
involves mixing of atomic orbitals to form hybrid orbitals to be used for bonding
- orbitals of the same atom, which are of different energy, can overlap to form hybrid orbitals of equal energy
- no. of hybrid orbitals formed is equal to that no. of atomic orbitals involved in the process of hybridisation that has taken place
3 main types:
- sp
- sp^2
- sp^3
145
sp^3 hybridisation
- involves mixing of one s orbital and three p orbitals to form four sp^3 orbitals
- these adopt a tetrahedral arrangement
- the hybrid orbitals repel each other and arrange themselves in a config. that minimises e- repulsion by spreading themselves as far apart as possible
146
sp^2 hybridisation
- one 2s orbital mixes w/ tow of the 2p orbitals to form three sp^2 hybrid orbitals
- one 2p orbital doesn't participate in the hybridisation process, leaving one unhybridised 2p orbital
147
sp hybridisation
- one 2s orbital mixes w/ a 2p orbital to form two sp hybrid orbitals
- one of the 2p orbitals is involved in the hybridisation, so there are two unhybridised 2p orbitals remaining
148
2 electron domains
Electron domain geometry: linear
Molecular geometry: linear
Hybridisation: sp
149
3 electron domains
1. Electron domain geometry: trigonal planar
Molecular geometry: trigonal planar
Hybridisation: sp^2
2. Electron domain geometry: trigonal planar
Molecular geometry: V-shaped
Hybridisation: sp^2
150
4 electron domains
1. Electron domain geometry: tetrahedral
Molecular geometry: tetrahedral
Hybridisation: sp^3
2. Electron domain geometry: tetrahedral
Molecular geometry: trigonal pyramidal
Hybridisation: sp^3
3. Electron domain geometry: tetrahedral
Molecular geometry: V-shaped
Hybridisation: sp^3
151
What determines how many electrons are lost or gained?
The no. of e- lost or gained is determined by the electron configuration of the atom
152
Increasing the bond length and strength
Bond length decreases and bond strength increases as the no. of shared electrons increases
153
Deducing the polar nature of a covalent bond
- polar bonds result from unequal sharing of electrons
- the more electronegative atom exerts a greater pulling power on the shared electrons
- hence this holds the electrons more than the other
- hence, electron distribution is unsymmetrical
- this is because one end of the molecule has more electrons than the other
- the more electronegative atom has the greatest electron density when bonded
154
Molecular polarity
- If the dipoles in the BONDS cancel out then the MOLECULE will be non-polar
- If the net BOND dipoles are non-zero then the MOLECULE will be polar - electronegativity values are needed to calculate bond dipoles and molecular geometry to see if these cancel
155
Non-directional bonding in alloys
production of alloys is possible because of the non-directional nature of the delocalised electrons
- the lattice can accomodate ions of different sizes
156
Alloys vs. regular metals
Alloys are usually more stronger than regular metals
- because if different atoms are present, the regular network of +ve ions is disturbed
- atoms of a different size make it harder for layers of +ve ions to slide over each other
- thus prevent bending or denting of the metal
157
Covalent bonding in terms of pi and sigma bonds
- a covalent bond results from the overlap of atomic orbitals
158
Resonance
using two or more Lewis structures to represent a particular molecule or ion
IB Chemistry HL
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