3.1.3 - Bonding Flashcards
(143 cards)
1
Q
Define covalent bonding
A
When two atoms share pairs of electrons
2
Q
What is a dative covalent bond?
A
A dative covalent bond forms when the shared pair of electron in the covalent bond come from only one of the bonding atoms.
3
Q
What is metallic bonding?
A
Metallic bonding is the electrostatic force of attraction between positive metal ions and delocalised electrons.
4
Q
What are the 3 main factors affecting the strength of metallic bonding?
A
- Nuclear Charge
- Number of Delocalised Electrons per Atom / Charge on Ion
- Size of Ion (smaller ions, stronger bond)
5
Q
What structure do ionic structures take?
A
Giant Ionic Lattice
6
Q
What are the properties of ionic compounds?
A
- High melting point and boiling point because of giant lattice of ions with strong electrostatic forces between oppositely charged ions, requires lots of energy to break.
- Poor conductors of electricity when solid / can conduct when molten/aqueous as ions are free to move and carry charge.
7
Q
Explain 3 key properties of metals
A
- High boiling / melting points - strong electrostatic forces of attraction between +ive ions and delocalised electrons
- Good conductors of electricity - delocalised electrons can move through the structure and carry a charge
- Malleable / Ductile - layers of ions can slide over each other, held together by electrostatic forces
8
Q
Describe properties of simple molecules
A
- Low boiling / melting points - due to weak intermolecular forces between molecules e.g van der waals, hydrogen bonds
- Poor conductivity as there aren’t any ions and electrons are localised (fixed in place)
9
Q
How does the presence of lone pairs affect bond angles?
A
2.5° for each lone pair
10
Q
What shape is formed from 4 bp and 1 lp?
A
see-saw
11
Q
What shape is formed from 3 b.p and 2 l.p?
A
T shape
12
Q
What shape is formed from 3 l.p and 2 b.p?
A
Linear
13
Q
What shape is made from 4 bond pairs and 2 lone pairs?
A
Square Planar
14
Q
What is electronegativity?
A
Electronegativity is the power of an atom to attract bonded electrons in a covalent bond.
15
Q
What are the most electronegative atoms?
A
F, O, N, Cl
16
Q
What scale is electronegativity measured on?
A
Pauling scale (ranging from 0 to 4)
17
Q
How does electronegativity change across a period?
A
Electronegativity will increase across a period as the number of protons increases but there is similar shielding. Furthermore, the atomic radius decreases as the electrons in the same shell are pulled in more.
18
Q
How does electronegativity change down a group?
A
Electronegativity will decrease down a group because the distance between the nucleus and bonded electrons increase and the shielding of the inner shell electrons increases.
19
Q
If electronegativity is similar, what type of bonding could be present?
A
If both are <2 in electronegativity, bonding is metallic. If both are >=2 in electronegativity, bonding is non-polar covalent.
20
Q
If electronegativity is different, what type of bonding could be present?
A
If the difference in electronegativity is >0.5: polar covalent. If difference in electronegativity >= 2: Ionic.
21
Q
What’s important to note about the polarity of symmetric molecules?
A
It will not be polar even if the individual bonds within the molecule are polar. This individual dipoles on the bonds cancel out - there is no net dipole moment.
22
Q
What is a dipole moment?
A
The overall effect of polarity of the bonds in a molecule.
23
Q
What are Van Der Waals’ Forces?
A
These are the weakest type of intermolecular force that occurs between all molecular substances and noble gases.
24
Q
What are the factors affecting Van Der Waals’ forces?
A
1) More electrons = stronger VDW forces
2) Bigger surface area = Stronger VDW forces.
25
Why does the b.p of halogens down group 7 increase?
Increasing number of electrons in the bigger molecules hence stronger VDW forces form.
26
Why do long chain alkanes have a higher b.p than spherical shaped alkanes?
There is a larger surface area of contact between chained molecules than there is spherical molecules - hence there are stronger VDW forces.
27
Describe permanent dipole-dipole forces (2)
1. Electrostatic forces between polar molecules
2. Stronger than VDWs so compounds have higher BP.
28
What is hydrogen bonding?
The strongest type of IM force that forms between Hydrogen and F, O, N. The lone pair on these atoms attract a hydrogen atom on another molecule.
29
Use hydrogen bonding to explain why ice floats in water
1) Since ice floats in water, it must be less dense than water.
2) The hydrogen bonds in ice hold the molecules further apart so density is lower whereas in water the hydrogen bonds are constantly breaking and reforming since the molecules move.
30
What are the 3 types of IM forces? List from weakest to strongest.
1. Van Der Waals
2. Dipole-Dipole
3. Hydrogen Bonds.
31
What are the 4 types of crystal structures?
Ionic, Metallic, Simple Molecular, Giant Covalent.
32
Describe the properties of simple molecular compounds (4)
1) low m.p/b.p because of weak IM forces (VDW)
2) poor solubility in water
3) poor conductivity of electricity in solid/when molten as there are no ions / electrons are localised
4) generally mostly gases and liquids.
33
Describe properties of macromolecular compounds (5)
1) high mp/bp because of many strong covalent bonds in macromolecular structure, takes a lot of energy to break the many strong bonds.
2) insoluble in water.
3) diamond and sand poor, because electrons are localised / graphite good, as free delocalised electrons between layers.
4) poor conductivity when molten.
5) generally solids.
34
State 5 properties of metals
1) high mp/bp due to strong attraction between +ve ions and sea of delocalised electrons.
2) insoluble.
3) good conductors of electricity.
4) shiny.
5) malleable, as layers of ions can slide over each other.
35
Describe the structure of diamond (2)
1) macromolecular
2) tetrahedral arrangement of carbon atoms.
36
Describe the structure of graphite
1) trigonal planar arrangement of carbon in layers.
2) 3 covalent bonds per atom in each layer, 4th is delocalised.
37
Describe the structure of ice (3)
1) molecular structure
2) tetrahedral arrangement
3) molecules held further apart than in liquid.
38
Describe the structure of Iodine (1)
Regular arrangement of molecules held together by weak VDW forces.
39
Solid: Arrangement and Movement
1. Tightly packed in a regular arrangement
2. Vibrate in fixed positions.
40
Liquid: Arrangement and Movement
1. Tightly packed in a random arrangement
2. Particles move freely and have more energy than in a solid.
41
Gas: Arrangement and Movement
1. Spaced out and in a random arrangement
2. Particles move freely and have lots of energy.
42
Explain, in terms of electronegativity, why the boiling point of H2S2 is lower than H2O2.
1. The electronegativity of S is lower than the electronegativity of O.
2. The difference between H and S electronegativity is less.
3. Hence, S and O have greater delta positive/negative charge, stronger bonds require more energy to break.
4. There is no hydrogen bonding between the H2S2 molecules, only van der Waals forces.
43
State the meaning of the term electronegativity.
The power of an atom or nucleus to withdraw or attract a pair of electrons in a covalent bond.
44
Explain, in terms of its structure and bonding, why titanium has a high melting point.
There is strong attraction between the number of protons and delocalised electrons.
45
Explain, in terms of structure and bonding, why the melting point of carbon is high.
1. Macromolecular structure is giant.
2. Covalent bonds in the structure are strong and require lots of energy to break/overcome.
46
Describe the structure of and bonding in graphite and explain why the melting point of graphite is very high.
1. Layers of C atoms.
2. Are connected by covalent bonds.
3. van der Waals forces between the layers.
4. Strong covalent bonds are what are broken during melting.
47
Explain, in terms of the intermolecular forces present in each compound, why HF has a higher boiling point than HCl.
HF has hydrogen bonding. HCl has permanent dipole-dipole bonding. Hydrogen bonding is stronger.
48
Why do diamond and graphite both have high melting points?
1. Macromolecular structures.
2. Covalent bonds between atoms.
3. These are strong bonds and it requires lots of energy to break them.
49
Why is graphite a good conductor of electricity?
Delocalised electrons can carry charge.
50
Explain why the melting point of magnesium is higher than that of sodium.
1. Mg2+ have a higher charge than Na+.
2. Shorter distance between e- and ions in Mg2+.
3. Hence stronger metallic bonding.
51
What are the 2 conditions for hydrogen bonding to occur?
1. Attraction between lone pair on F,O,N and H.
2. H connected to F, O, N.
52
What is Beryl Chloride (bonding)?
Covalent
53
What is a Bonding pair of electrons?
A pair of electrons shared between 2 atoms.
54
Define metallic bonding.
The attraction between a positive metal ion and a sea of delocalised electrons.
55
Explain the structure of metallic bonding.
1. The attraction between a positive metal ion and a sea of delocalised electrons.
2. Delocalised electrons move throughout the structure where they are malleable and ductile.
3. Layers can slide.
56
Define a dative covalent bond/ co-ordinate bond.
Where one atom donates 2 electrons to form a shared pair of electrons.
57
Why is a covalent bond a covalent bond?
A small difference in electronegativity.
58
Why is an ionic bond an ionic bond?
Large difference in electronegativity.
59
Explain the structure of iodine.
Simple molecular covalent bonds with weak vdw forces between its molecules.
60
Explain what bond graphite has?
Covalent macromolecular.
61
Explain why the melting point of graphite is so high and explain its bonding.
1. Graphite is a macromolecular structure, with layers of carbon atoms.
2. Each carbon atom is attached to 3 other carbons.
3. Between the layers there are weak van der Waals forces.
4. Where covalent bonds require a lot of energy to break.
62
State why Iodine is a poor conductor of electricity.
No delocalised electrons.
63
Suggest why graphite is a good conductor of electricity.
Delocalised electrons which move throughout the structure.
64
What type of bond does magnesium have?
Metallic bonding.
65
2 types of covalent bonds:
1. Simple covalent (simple molecular)
2. Macromolecular.
66
Explain the bonding of Diamond.
1. Bonded to 4 carbon atoms strong i.m.f.
2. Cannot conduct electricity.
67
Which molecules have hydrogen bonding?
Oxygen, Fluorine, Nitrogen.
68
What is dipole-dipole interactions?
When there is a large variation in electronegativity between the molecule and the atom.
69
What are the bond angles for a trigonal planar molecule?
120
70
What are the bond angles for a tetrahedral molecule?
109.5
71
What are the bond angles for a trigonal bi-planar molecule?
120 and 90 degrees
72
What are the bond angles for an octahedral molecule?
90 degrees
73
What do the lone pair of electrons do to bonding electrons?
Lone pair electrons repulse more than bonding electrons, reducing the bond angle by 2.5.
74
What are the bond angles for a trigonal bipyramidal molecule?
120 and 90 degrees
75
What do the lone pairs of electrons do to bonding electrons?
Lone pair electrons repulse more than bonding electrons, reducing the bond angle by 2.5
76
What is the bond angle of a bent molecule?
104.5
77
What is the bond angle of a trigonal pyramid?
107
78
What statement about inorganic ions is always correct?
They form giant structures
79
What happens to the dipole of a molecule if it is symmetrical?
The dipole-dipole interaction is no longer the interaction as symmetry cancels out the dipoles.
80
Explain how a lone pair of electrons influences the bond angle.
- Lone pairs repel more than bond pairs
- Bond angle will be lower.
81
Explain how Nitrogen Monoxide is formed.
- Nitrogen and Oxygen from the air react
- At high temperatures.
82
Explain why lots of energy is needed to break the bond between lithium fluoride.
- Lithium fluoride forms ionic bonds between their molecules.
- There is a strong electrostatic attraction
- Where Li+ and F- are oppositely charged.
83
How would you check all the water has been used up?
Re-heat and re-weigh and check that the mass is unchanged
84
Which elements have metallic bonding?
M- Magnesium
A - Aluminum
Ca - Calcium
L - Lithium
S- Strontium
B - Beryllium
A - Aluminum
Ca - Calcium
L - Lithium
S- Strontium
B - Beryllium
85
Define an ionic bond.
86
Most electronegative atoms?
F O N Cl
87
What are the 4 types of crystal structures?
88
Define ionic bonding.
- Ionic bonding is the electrostatic force of attraction between oppositely charged ions formed by the transfer of electrons from metal to non-metal atoms.
89
What factors affect the strength of ionic bonding?
- The strength of ionic bonding increases with smaller ion sizes and higher charges on the ions, which enhance the electrostatic force between them.
90
Explain why MgO has a higher melting point than NaCl.
MgO has a higher melting point than NaCl because Mg2+ and O2- ions are smaller and have higher charges than Na+ and Cl-, resulting in stronger ionic bonds.
91
Define a covalent bond.
A covalent bond is a shared pair of electrons between two atoms, typically non-metals, allowing each atom to attain a stable electronic configuration.
92
What is a dative covalent bond?
A dative covalent bond, also known as a coordinate bond, forms when both electrons shared in the bond come from the same atom.
93
Describe metallic bonding.
Metallic bonding is the electrostatic attraction between positive metal ions and delocalized electrons around them, which allows electrons to move freely, contributing to metal properties like conductivity and malleability.
94
What factors strengthen metallic bonding?
Stronger metallic bonding is due to a greater number of delocalized electrons, a higher nuclear charge (more protons), and smaller metallic ions, which increase the electrostatic forces.
95
Explain why magnesium has a higher melting point than sodium.
Magnesium has a higher melting point than sodium because it has more delocalized electrons and a smaller ionic radius, resulting in stronger metallic bonds.
96
What is the shape and bond angle of a molecule with a tetrahedral structure?
A tetrahedral molecule has a shape where four atoms are symmetrically positioned around a central atom with bond angles of approximately 109.5 degrees.
97
Describe the electron pair geometry and molecular shape of ammonia (NH3).
Ammonia has a trigonal pyramidal shape with one lone pair and three bonding pairs of electrons. The bond angle is slightly less than 109.5 degrees, typically around 107 degrees due to lone pair-bond pair repulsion.
98
Define electronegativity.
- Electronegativity is the measure of an atom's ability to attract shared electrons in a chemical bond.
99
How does electronegativity change across a period and down a group in the periodic table?
Electronegativity increases across a period due to a decrease in atomic radius and an increase in nuclear charge. It decreases down a group due to increased electron shielding and greater atomic radius.
100
What is a polar covalent bond?
- A polar covalent bond is a type of covalent bond where electrons are unequally shared between atoms, leading to a partial positive charge on one atom and a partial negative charge on the other.
101
What determines if a molecule is polar?
A molecule is polar if it contains polar bonds and the molecular geometry does not allow for the dipoles to cancel out, resulting in an asymmetrical distribution of electrical charge.
102
Explain the concept of hydrogen bonding and its effects on physical properties.
Hydrogen bonding is a strong type of dipole-dipole interaction that occurs when hydrogen is bonded to highly electronegative atoms like fluorine, oxygen, or nitrogen. This bonding leads to higher boiling and melting points.
103
What are van der Waals forces and how do they affect molecular properties?
Van der Waals forces are weak intermolecular forces caused by temporary fluctuations in electron density, creating temporary dipoles. They are responsible for the physical states of molecular substances at different temperatures.
104
Compare and contrast the properties of ionic, metallic, macromolecular, and simple molecular structures.
- Ionic structures have high melting and boiling points and conduct electricity when molten.
- Metallic structures conduct electricity and heat and are malleable.
- Macromolecular structures (like diamond) have extremely high melting points and are typically hard and brittle.
- Simple molecular structures generally have low melting and boiling points due to weak van der Waals forces.
105
Describe the structure and properties of diamond and graphite.
Diamond has a tetrahedral structure with each carbon atom forming four covalent bonds, making it extremely hard and an excellent insulator. Graphite has layers of carbon atoms arranged in hexagons, with weak forces between layers, making it a good lubricant and conductor of electricity.
106
What role do lone pairs play in molecular shape?
Lone pairs repel bonding pairs more strongly than bonding pairs repel each other, altering bond angles and affecting molecular geometry.
107
Explain the molecular shape and bond angles in water (H2O).
- Water has a bent molecular shape with two bonding pairs and two lone pairs on the oxygen atom, resulting in a bond angle of about 104.5 degrees due to the repulsion between lone pairs.
108
What are the characteristic properties of ionic compounds?
Ionic compounds typically have high melting and boiling points due to strong electrostatic forces between ions. They generally conduct electricity when molten or dissolved because ions are free to move.
109
Define and give examples of molecules formed by simple molecular structures.
- Simple molecular structures are composed of molecules held together by weak intermolecular forces like van der Waals forces. Examples include water (H2O), carbon dioxide (CO2), and methane (CH4).
110
What is a giant ionic lattice?
- A giant ionic lattice is a three-dimensional structure of oppositely charged ions bonded together by strong ionic bonds throughout the entire crystal.
111
Describe the electron configuration and ion formation for magnesium and oxygen.
Magnesium (Mg) has an electron configuration of 1s2 2s2 2p6 3s2 and forms Mg2+ by losing two electrons. Oxygen (O) has an electron configuration of 1s2 2s2 2p4 and forms O2- by gaining two electrons.
112
How does ionic radius change across a period and within a group?
- Ionic radius decreases across a period due to increasing nuclear charge which attracts electrons closer to the nucleus. It increases down a group as ions have more electron shells.
113
Explain the concept of metallic bonding with an example.
Metallic bonding occurs through the attraction between delocalized electrons and positive metal ions within a metal. For example, in magnesium, the electrons from the outer shell are delocalized, contributing to the metallic bond.
114
What influences the boiling and melting points of macromolecular structures?
- Macromolecular structures have high boiling and melting points due to strong covalent bonds throughout the structure, requiring significant energy to break.
115
Discuss the conductivity properties of ionic, molecular, macromolecular, and metallic substances.
- Ionic compounds conduct electricity when molten or in solution.
- Simple molecular substances do not conduct due to lack of free electrons or ions. Macromolecular substances like graphite conduct due to free electrons, while others like diamond do not.
- Metallic substances conduct due to free-moving delocalized electrons.
116
What is the role of lone pairs in molecular geometry?
Lone pairs repel bonding pairs of electrons, often causing changes in bond angles and molecular geometry to minimize repulsion.
117
What is hydrogen bonding and why is it significant?
- Hydrogen bonding is a strong type of dipole-dipole interaction occurring when hydrogen is bonded to highly electronegative elements like nitrogen, oxygen, or fluorine. It significantly affects physical properties like boiling point and solubility.
118
How do Van der Waals forces vary among molecules?
Van der Waals forces vary in strength based on the number of electrons in a molecule; more electrons lead to stronger forces and thus higher boiling points.
119
What is a dative covalent bond, and can you provide an example?
- A dative covalent bond occurs when both electrons shared in the bond come from one atom. An example is the bond between the nitrogen and hydrogen in the ammonium ion (NH4+).
120
Describe the effect of ionic size and charge on the melting point of ionic compounds.
- Smaller ionic size and higher charge increase the electrostatic forces between ions, leading to higher melting points for ionic compounds.
121
Explain why metals are generally malleable and good conductors of electricity.
Metals are malleable because the layers of atoms can slide over each other without breaking bonds. They are good conductors due to the presence of delocalized electrons that can move freely and carry electrical current.
122
Differentiate between polar and non-polar molecules with examples.
- Polar molecules, like water (H2O), have an uneven distribution of electron density leading to dipole moments. Non-polar molecules, like methane (CH4), have symmetrical electron distribution which cancels out dipole moments.
123
What determines the solubility of substances in water?
- Solubility in water is determined by the ability to form favorable interactions with water molecules. Ionic and polar compounds are generally soluble due to their ability to interact with the polar water molecules.
124
Explain the structure and properties of graphite and how they differ from diamond.
- Graphite consists of layers of carbon atoms arranged in hexagons with weak forces between layers, allowing layers to slide and making graphite a good lubricant and conductor.
- Diamond consists of a rigid tetrahedral structure of carbon atoms, making it extremely hard and an insulator.
125
How does molecular shape influence boiling and melting points?
- Molecular shape can influence intermolecular forces; for example, linear molecules have less surface contact and weaker van der Waals forces compared to branched molecules, resulting in lower boiling and melting points.
126
Describe the factors that affect the strength of metallic bonds.
- The strength of metallic bonds is influenced by the number of delocalized electrons, the charge and size of metal ions, and the electron density.
- More delocalized electrons and smaller, more charged ions strengthen the bond.
127
Why does electronegativity increase as you go across a period?
The number of protons increases. The atomic radius decreases because the electrons in the same shell are pulled in more.
128
What are sigma bonds?
Where the atomic orbitals overlap directly along the internuclear axis. All single bonds are sigma bonds.
129
What are pi bonds?
Where the atomic orbitals overlap above and below the internuclear axis. All double bonds contain a sigma and a pi bond.
130
Why does electronegativity increase as you go across a period?
The number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more.
131
What are Sigma Bonds?
Where the atomic orbitals overlap directly along the internuclear axis. All single bonds are sigma bonds.
132
What are Pi Bonds?
Where the atomic orbitals overlap above and below the internuclear axis. All double bonds contain a sigma and a pi bond. All triple bonds contain a sigma bond and 2 pi bonds.
133
What are some factors which affect the strength of VDW forces?
The more electrons there are in the molecule, the higher the chance temporary dipoles will form. This means VDW forces will be stronger and boiling points will be greater.
134
Draw the structure of diamond.
135
Draw the structure of graphite.
136
Which molecule is not able to form a co-ordinate bond with another species?
CH4
137
Which species has a square planar shape?
XeF4
138
Which substance contains delocalised electrons?
Graphite
139
Which change occurs when water is vaporised?
Intermolecular forces are overcome.
140
Explain the difference in boiling points between ethanol and methoxymethane.
Ethanol has higher boiling points due to hydrogen bonding, which is stronger than the van der Waals forces present in methoxymethane.
141
Deduce the type of intermolecular forces in SiF4 and explain how this type of force arises.
SiF4 has van der Waals forces due to temporary dipoles created by the uneven distribution of electrons in one molecule inducing a dipole in a neighboring molecule.
142
Give the meaning of the term electronegativity and explain how permanent dipole-dipole forces arise between hydrogen chloride molecules.
Electronegativity is the ability of an atom to attract a pair of electrons in a covalent bond. In hydrogen chloride, the difference in electronegativity between hydrogen and chlorine creates a permanent dipole, with δ+ on hydrogen and δ- on chlorine, leading to permanent dipole-dipole interactions.
143
In terms of the intermolecular forces involved, suggest why iodine requires more heat energy for melting than chlorine does.
- Iodine is a bigger molecule.
- Therefore, it has stronger induced dipole-dipole forces.
