Chapter 4 ~chemical Bonding Flashcards
Definitions and key elements (96 cards)
1
Q
What are the 3 types of intramolecular forces?
A
- Ionic bonds
- Covalent bonds
- Metallic bonds
2
Q
Intramolecular definition :
A
Forces that hold atoms together within a molecule
3
Q
Definition of intermolecular forces :
A
The weak forces that exist between molecules
4
Q
The 3 types of intermolecular forces
A
- Van Der Waal’s forces (also called dispersion, London, temporary dipole-induced dipole forces)
- Permanent dipole-dipole forces
- Hydrogen bonds
5
Q
How are ions formed?
A
When atoms gain or lose electrons
6
Q
Definition of an ionic bond
A
°the electrostatic attraction between oppositely charged ions (happens between non-metals and metals)
7
Q
Definition of electrovalent
A
°it is another name for an ionic bond which is the electrostatic attraction between oppositely charged ions
8
Q
Definition of dot and cross diagrams
A
A diagram showing the arrangement of the outer shell electrons in an ionic or covalent element or compound. The electrons are shown as dots or crosses to show their origin
9
Q
Definition of a covalent bond
A
A bond formed by the sharing of pairs of electrons between 2 atoms (non metal and non metal)
10
Q
Definition of double covalent bond plus 3 examples
A
°atoms sharing 2 pairs of electrons
Examples : oxygen molecule, carbon dioxide (between the oxygen and carbon atom) and ethene (between the 2 carbons)
11
Q
Definition and one example of a triple covalent bond
A
° when atoms bond together by sharing 3 pairs of electrons
Examples: nitrogen molecule
12
Q
Definition of a co-ordinate bond or dative covalent bond and 2 examples
A
°a covalent bond in which both electrons in the bond come from the same atom
Or
It is formed when 1 atom provides both the electrons needed for a covalent bond
Examples :
- Ammonium ion NH4+ (the hydrogen is electron deficient)
- Al2Cl6 (the AlCl3 is electron deficient)
13
Q
For a dative covalent bond we need two things :
A
- 1 atom having a line pair or electrons
2. A 2nd atom having an unfilled orbital to accept the lone pair (an electron deficient compound)
14
Q
Electron deficient compound definition and example
A
° a noble gas configuration is not achieved and the compound too little electrons
Example : boron trifluoride BF3 (only has 6 e-)
15
Q
What is an expanded octet and an example of one
A
° too many electrons to achieve a gas configuration
Example: sulfur hexaflouride SF6 (has 12 e-)
16
Q
Aluminum chloride at high and low temperatures
A
High temperature : AlCl3 (electron deficient - still needs 2e-)
Low temperature : Al2Cl6 (the AlCl3 molecules bond bec of lone pairs of 2 Cl atoms & form a dative covalent bond with Al atoms
17
Q
What is bond energy/ bond enthalpy ?
A
°the energy required to break 1 mole of a particular (given) bond in 1 mole of gaseous molecules
18
Q
What is bond length?
A
The average distance between nuclei of 2 bonded atoms in a molecule
19
Q
Why are double bonds shorter than single bonds?
A
Bec double bonds have a greater quantity of negative charge between the 2 atomic nuclei.
The greater force of attraction between the e-s and the nuclei pulls the atoms closer together.
Stronger bond
20
Q
3 things that determine chemical reactivity
A
- Bond strength
- Polarity
- Whether bond is sigma or pi bond
21
Q
A reaction on happens when…
A
A successful collision 💥 occurs with enough energy to break bonds in either or both molecules
22
Q
Why is nitrogen unreactive?
A
It has a triple bond that requires a lot of energy to break the nitrogen atoms apart
23
Q
What is the relationship between bond energy and bond length
A
Inversely proportional. As bond length decreases bond strength (bond energy) increase
24
Q
Electron repulsion theory
A
× E-s are all negative and therefore repel one another
× The repulsion forces the pairs of e-s apart until their repulsive forces are minimized
×order of repulsion:
Lone pair - lone pair (most repulsion) >lone pair - bond pair >bond pair - bond pair (least)
25
The shape and bond angles of a covalently bonded molecule depend on 2 things
1. Number of pairs of electrons around each atom
| 2. Whether these pairs are lone or bonding pairs
26
Order of repulsion
| Most to least
Lone pair - lone pair > lone pair - bond pair > bond pair- bond pair
27
Why does a lone pair repel more than a bond pair?
Lone pairs have a more concentrated e- charge cloud than bonding pair of e-.
Their cloud charges are:
×wider
×Slightly closer to the nucleus of the central atom
28
Does bond length increase or decrease down group 17?
Increase (the atoms get larger so, so does the distance)
29
Are diatomic molecules polar or non polar
Non polar
30
Name the 7 types of shapes of molecules, their bond angles, an example and amount of lone pairs
1. Linear (180°) e.g. CO2.... Each oxygen molecule has 2 lone pairs
2. Trigonal planar (120°) e.g. BF3.. no lone pairs
3. Tetrahedral (109.5°) e.g. CH4.... No lone pairs
4. Trigonal pyramidal (107.5°) e.g. NH3... One lone pair
5. Non-linear V-shaped (104.5°) e.g. H2O... 2 lone pairs
6. Octahedral (90°) e.g. SF6..... No lone pairs
7. Trigonal bipyrammidal (90°) e.g. PCl5/ PF5.....
No lone pairs
31
How are sigma bonds formed?
When hybridised orbitals overlap linearly (end on) and a covalent bond forms
32
Is the e- density of each sigma bond symmetrical or not symmetrical?
Symmetrical
33
How are pi bonds formed?
Sideway overlap of p orbitals
34
Is a pi bond's electron density symmetrical or not?
Not symmetrical
35
2 clouds of electrons in a pi bond represents
One bond consisting of 2 e-
36
Shape of S orbital
Spherical
37
Shape of p orbital
Dumb-bell shaped
38
Definition of hybridisation
The process of fusing atomic orbitals
39
The 3 types of hybrids and their characters
1. sp~3 :each orbital has 1/4 s and 3/4 p
2. sp~2 : 1 s orbital and 2 p- orbitals are hybridised
3. sp : 1 s orbital and 1 p orbital is hybridised
40
What is electron density?
The measure of the probability of an e- being present in a specific location
41
What are delocalised electrons?
Electrons that are not associated with any one particular atom or bond.
They can move between 3 or more adjacent atoms
42
Why are metallic bonds strong?
Bec the ions are held together by strong electrostatic attraction between their positive charges and the negative charges of the delocalised e-
This electrostatic attraction acts in all directions.
43
The strength of a metallic bond increases with 3 things
1. Increase in positive charge on ions in metal lattice ( bigger lattice =stronger electrostatic forces of attraction )
2. Decreasing size of metal ions in lattice (due to charge density increasing)
3. Increasing number of mobile e- per atom (contribute to a bigger lattice = stronger electrostatic forces of attraction)
44
Most metals have high melting and boiling points why?
It takes a lot of energy ⛮ to weaken the strong attractive forces between the metal ions and the delocalised e-. These attractive forces can only be overcome at high temperatures
45
Exception *mercury(has low MP and BP)
Why?
Liquid at room temperature.
Bec some of the e-s in a mercury atom are bound more tightly than usual to the nucleus, weakening the metallic bonds between atoms
46
How do metals conduct electricity?
Current can flow bec the delocalised e- are free to move (and therefore can borne / carry / transfer the current)
47
Why can covalent solids *not* conduct electricity?
| Also name one *exception to this
Bec none of their e-s are free to move
*except : graphite
48
2 reasons why metals can conduct heat
1. Partly due to the movement of delocalised electrons
| 2. Partly due to the vibrations passed on from one metal ion to the next
49
Strength in order of the 3 intermolecular forces
Hydrogen bonding (strongest),
permanent dipole-dipole forces
and then van Der Waal's forces (weakest)
50
Strength in order of the 3 intermolecular forces
Hydrogen bonding (strongest),
permanent dipole-dipole forces
and then van Der Waal's forces (weakest)
51
Definition of electronegative
°It is the ability of a particular atom, which is covalently bonded to another atom, to attract the bond pair of electrons towards itself
52
Does electronegativity increase or decrease across a period from group 1 to 17
Increases
53
Does electronegativity increase or decrease down each group?
Decreases down each group
54
The 5 most electronegative elements in order from the most electronegative to the least
F(most)
O
N
Cl
Br (least)
55
Definition of non-polar
A molecule with no separation of charge ; it will not be attracted to a positive or negative charge.
(the pair of e-s in the covalent bond are equally shared)
E. G. Diatomic molecules
56
Definition of a polar bond (or that it has a dipole)
A separation of charge in a molecule. One end of the molecule is permanently positively charged and the other is negatively charged
E. G. HCl where H is slightly positive and Cl is slightly negative
57
When a covalent bond is formed between elements that have different electronegativity values, the more electronegative atom attracts the pair of e-s in the bond towards it
As a result 4 things happen :
1. The centre of + charge doesn't coincide with the centre of - charge
2. the e- distribution is asymmetric
3. The 2 atoms are partially charged :
One is slightly negative and one is slightly positive
4. The bond is polar (or has a dipole)
58
The degree of polarity is measured as a dipole moment. When does a bond become more polar?
As the difference in electronegativity values of atoms involved in the covalent bond increase
59
Why do some molecules contain polar bonds but have no overall polarity?
The polar bonds are arranged in such a fashion that the dipole moments cancel each other out
E. G tetrachloromethane CCl4 has 4 polar bonds but the dipole in each bond cancel each other out so CCl4 is NON-POLAR
60
The charge distribution in molecules and ions can be determined by a method called?
X-ray spectroscopy (firing X-rays @ molecules and measuring energy ⛮ given off by e-)
61
Bond polarity influences..
Chemical reactivity
62
Name 2 molecules that form a triple bond
1. Nitrogen (non polar)
| 2. Carbon monoxide (polar)
63
Polarity influences chemical reactivity. Many chemical reactions are started by a reagent.....
Attacking one of the electrically charged ends of a polar molecule.
64
Which molecule is more reactive chloroethane (C2H5Cl) or ethane (C2H6)? Plus an explanation
Chloroethane (C2H5Cl)
Bec reagents such a OH- ions(has one polar bond) can attack the slightly positive Carbon of the polarised C-Cl bond. It is the mechanism of nucleophilic substitution in halogenoalkanes.
Such an attack is not possible with ethane bec the C-H bond is virtually non-polar (this explains why alkanes are not very reactivity)
65
Definition of van Der Waal's forces (also called temporary dipole - induced dipole forces ) :
The weak forces of attraction between molecules caused by the formation of temporary dipoles
66
What are the intermolecular forces that keep bromine a liquid at room temperature?
Van Der Waal's forces
67
How is a temporary dipole formed? (3 points)
e- charge clouds in a non-polar molecule are constantly moving.
However, often more of the charge cloud is on one side of the molecule than the other.
This means for a short moment, one end of the molecule is more negative than the other and a temporary dipole is set up.
68
Why are temporary dipoles always temporary?
Bec electron clouds are constantly /always moving
69
A temporary dipole can induce a dipole on neighboring molecule and form a:
temporary dipole - induced dipole forces which are also called :
Van Der Waal's forces
70
Van Der Waal's forces increase with 2 things :
1. Increasing number of e- (and protons) in the molecule
2. Increasing number 🔢 of contact points between molecules - contact points are places where molecules come close together
71
Why does the boiling and enthalpy of vaporisation of the noble gases increase as you go down the group?
Bec the number of e-and p+ increase and therefore the van Der Waal's forces also increase in strength.
72
Definition of permanent dipole-dipole forces :
A type of intermolecular force between molecules that have permanent dipoles.
The attractive force between the slightly positive on one molecule and the slightly negative on a neighbouring molecule causes a weak attractive force between the molecules
73
For small molecules with the same number of e-s, permanent dipole-dipole forces are often stronger than
Plus an example : why butane has a lower boiling point than propanone
Van Der Waal's forces
Ans ex: butane is a gas at room temperature bec it has weak van Der Waal's forces
Propanone is a liquid bec the permanent dipole-dipole forces between the molecules are strong enough to make this substance a liquid at room temperature
74
Definition of hydrogen bonding
°the strongest type of intermolecular force - it is formed between molecules having a hydrogen atom bonded to one of the most electronegative elements (F, O or N)
75
What 2 things do we need for hydrogen bonding to occur?
1. One molecule having a hydrogen atom covalently bonded to F, O, N (3 most electronegative atoms)
2. A 2nd molecule having a F, O or N atom with an available lone pair of e-s
76
The hydrogen bonding is highly...
Polarised
77
Average number of hydrogen bonds formed per molecule depends on 2 thing:
1. #number of hydrogen atoms attached to F, O, N in molecule
| 2. The #number of lone pairs present on the F, O, N (e.g water 💦 has 2 lone pairs ( on O) so 2 hydrogen bonds can form)
78
Trend of BP of hydrogen halides namely :
HF, HCl, HBr and HI
Don't look @ H bec that's is common look at halides to compare
State the trend and then why
HF the highest BP >Hydrogen bonding
Then increases from HCl(smallest BP) , HBr to HI (2nd highest)> van Der Waal's forces
Explanation :
HF ~ hydrogen bonding bec F is highly electronegative and H-bonding is the strongest intermolecular force. Higher BP bec requires more energy to break
It drops to HCl and than increases from HBr to HI and this is bec as you go down the halide group the number of P+ and e- increases, increasing the size of the molecule which increases the strength of the van Der Waal's forces
79
Name the 3 peculiar properties of water
1. Enthalpy change of vaporisation & BP
2. Surface tension & viscosity (resistance to flow)
3. Ice is less dense than liquid water💦
80
Trend of enthalpy change of vaporisation and BP of group 16 hydrides namely:
H2O, H2S, H2Se and H2Te
And explanation
*H2O highest bec requires a lot of energy to break strong Hydrogen bonds (H-bond bec O)
* lowest is H2S
Then increases to H2Se and then H2Te (2nd highest)
Bec increasing number of e-s as you go down group 16. This leads to increased van Der Waal's forces
81
Why does water have a high viscosity?
H-bonding reduces ability of H2O molecules from sliding over each other
82
Why does water have a high surface tension?
H-bonding exerts a significant downward force at the surface of the liquid
83
Why is ice less dense than water 💦?
3D H-bonded network of water molecule produces a rigid lattice in which each O atom is surrounded by a tetrahedron of H atoms.
Due to the relatively long H-bonds, the water molecules are slightly further apart, therefore less dense
84
2 reasons why ionic compounds are solids @ room temperature & pressure
1. Strong electrostatic forces of attraction holding + & - ions together
2. The ions are regularly arranged in a lattice with oppositely charged ions close to each other.
85
Ionic compounds have high BP, MP & enthalpy changes of vaporisation because....
Requires a lot of energy to overcome strong electrostatic forces in the lattice
86
All metals = solids @ room temperature and pressure except ***
Why?
** mercury which is a liquid
Explanation :
Mercury's valence e- (full 6s shell) are bound tightly to the positive nucleus.
And therefore metallic bonding in mercury is weak bec e-s are not fully delocalised. So doesn't required a lot of energy to break
87
Most metals have high BP, MP & enthalpy changes of vaporisation because...
Requires a lot of energy to overcome the strong attractive forces between the positive ions and the 'sea' of delocalised e-s
88
Covalently bonded substances with a simple molecular structure such as H2O and NH3 are usually liquids or gases bec
Intermolecular forces =weak therefore doesn't require a lot of energy to break bonds
89
Giant molecular structure (covalent bonds) are solid @ room temperature and pressure bec
Van Der Waal's forces are considerable.
90
Are most ionic substances soluble in water and why?
Yes,
Bec water 💦 molecules are polar & they are attracted to ions on the surface of ionic solid.
These attractions are called ion-dipole attractions & these replace the electrostatic forces between ions and ions go into a solution
91
Are metals soluble in water?
No, but some react with water e.g. Sodium and calcium
92
Simple molecular structures (covalent) are divided into 2 groups concerning their solubility in water :
>insoluble bec non-polar so H2O molecules are NOT attracted to them e.g. Iodine
>soluble bec small molecules that can form H-bonds with water e.g. Ethanol
93
A hydrolysis reaction of covalent bonds is when
The covalent bond breaks e.g. In HCl it reacts with water to form soluble ions
94
Ionic compounds and their electrical conductivity :
Don't conduct electricity in solid state bec ions are fixed in the lattice and can only vibrate around a fixed point
Conduct electricity when molten bec ions are then mobile & can carry the current
95
Metals and their electrical conductivity :
Can conduct electricity in solid and molten bec has delocalised e-s that are mobile and carry the electric current
96
Covalent compounds with a simple molecular structure and their electrical conductivity
Do NOT conduct electricity because they have neither ions nor e-s that are mobile
