Bonding Flashcards
(117 cards)
1
Q
What is ionic bonding?
A
The force of attraction between two oppositely charged ions
2
Q
Metals loose electrons to form..?
Non- metals gain electrons to form?
A
- positive ions (cations)
- negative ions (anions)
3
Q
Explain ionic bonding in sodium chloride?
A
Sodium transfers an electron to chlorine. Sodium becomes a 1+ ion. Chlorine gain an electron forming a 1- ion. There is an attraction between Na+ and Cl- ions as one sodium atom can only satisfy one chlorine atom this is called the ionic bond
4
Q
What is the overall charge of an ionic compound?
A
No overall charge because the positive and negative ions balance each other out
5
Q
Explain why ionic compounds are solids at room temperature/ have high melting and boiling points?
A
- strong electrostatic force of attraction between oppositely charged ions
- a lot of energy is needed to overcome these strong forces
- so at room temperature all ionic compounds are solids with high melting and boiling points
6
Q
The electrostatic forces increase in strength as ..?
A
- the charge on the ion increases
- the size of the ion decreases
7
Q
Why is the melting point of CaF2 is higher than CaCl2?
A
- fluoride ion is smaller than chloride ion
- so there is a greater force of attraction between the calcium ion and the fluoride ion so more energy is needed to overcome these forces between the ions
8
Q
What is a crystal?
A
A solid with a regular shape which contains particles organised in a regular structure
9
Q
Ionic bonds in solids form..?
A
Giant ionic lattice
10
Q
How are ions arranged in a giant ionic lattice?
A
in a regular repeating pattern
11
Q
Describe the giant ionic lattice of sodium chloride crystal?
Draw the sodium chloride structures/ diagrams
A
- each ions is surrounded by 6 of the oppositely charged ions to form a face- cantered cubic crystal structure
- coordination number = 6
- the chloride ions is much larger than the sodium ion the sodium ion fits into the spaces between the chloride ions.
12
Q
the type of ionic lattice formed by an ionic compound depends on what..?
A
the relative sizes of the ions present
13
Q
What does Caesium chloride crystal form?
A
forms a body-centered cubic crystal instead of a face centered cubic crystal
14
Q
What are the properties of ionic crystals?
A
- high melting points because of many strong ionic bonds
- hard and brittle
- soluble in water
- do not conduct electricity - no free ions
- when melted can conduct
15
Q
Why can ionic crystals only conduct electricity when molten?
A
when an ionic compound is molten or dissolved, the ions are free to move and carry current throughout the structure but a solid has no free ions.
16
Q
What is covalent bonding?
How is it formed?
A
- Involved the sharing of a pair of electrons between 2 atoms
- two atomic orbitals overlap
17
Q
What determines the strength of a covalent bond?
A
The amount of overlap of the atomic orbitals
greater overlap = stronger bond
18
Q
What are lone pairs?
A
Non-bonded pairs of electrons in a compound
19
Q
How many electrons does phosphorus have in its outer shell?
How many electrons does aluminium have in its outer shell?
A
- 10
- 6
20
Q
What are multiple bonds?
A
Sometimes atoms can share 4 electrons to form double covalent bond or 6 electrons to form a triple covalent bond e.g nitrogen co2
21
Q
What is coordinate bonding/ dative covalency?
In order for coordination bonding to occur what must we have?
Show the formation of NH4+ and NH3BF3
A
Coordinate and dative covalent bonds are formed when one atom contributes both of its electrons needed for the covalent bond
One atom must have a lone pair and one atom must have a vacant orbital
22
Q
why do the arrows show in the Co-ordinate/ dative covalent bond?
A
shows the origin of the lone pair of electrons
23
Q
Two types of covalent structures?
A
Simple molecular and giant macromolecular
24
Q
Properties of simple covalent compounds?
A
Strong covalent bonds within the molecules but weak forces of attraction - little energy required
- low melting and boiling points
- gases or volatile liquids at room temp
- do not conduct electricity - no free ions or electrons
25
Describe and draw the structure of an iodine crystal?
- arranged in a regular repeating pattern
- face cantered cubic
- weak Van Der Waals forces so small amount of energy is needed to break
- diatomic iodine molecule
-does not conduct electricity - no free ions
- soft and brake easily
26
What happens when iodine is heated?
Changes from a black shiny solid into a purple vapour (sublimation) which separated the I2 molecules
27
Two examples that form giant covalent crystals?
Diamond and graphite which are allotropes of carbon
28
Sketch and Describe the structure and properties of diamond
- giant covalent structure - strongly covalently bonded to 4 other carbon atoms
- tetrahedral arrangement
- does not conduct - no free electrons
- hard
- strong melting and boiling
29
Describe the structure and properties of graphite?
- covalent bonded to 3 other carbon atoms
- strong covalent bonds between the carbon atoms in each layer
- electrons free to move between layers as they’re held by weak forces
- trigonal planar arrangement
- graphite is soft as layer can slide over each other
High melting and boiling points
30
what is electronegativity?
-electron negativity is the power of an atom to withdraw electron density from a covalent bond
in other words
-the ability of an atom to attract the bonding pair of electrons towards itself
31
Why are there no electronegativity values for group 0?
Because they do not form covalent bonds
32
What are the trends in electronegativity on the periodic table?
- electronegativity decreases down a group
- electronegativity increases across a period
- generally, increases from bottom left to the top right of the periodic table
33
What can be used to explain the electronegativity trends?
- size of the atom
- size of the nuclear charge
34
Why does electronegativity decrease down a group?
- Due to an increase in atomic radii.
- The covalent bond electron pair is further away from the nuclear charge and is shielded by the inner electrons.
- Therefore the atom is less able to withdraw electron density from the covalent bond
35
Explain why does electronegativity increase across a period?
- due to an increase in nuclear charge
- as we move across a period we are filling the same energy level and there is similar shielding
- therefore the atom is more able to withdraw electronegativity from the covalent bond
36
Which elements have the highest electronegativity values?
Nof elements
Nitrogen, oxygen, fluorine
37
What elements have the lowest electronegativity values?
Bottom left of periodic table (reactive metals)
Caesium, rubidium
38
How are non-polar molecules formed in terms of electronegativity ?
Forced between them are..?
When atoms of the same electronegativity are bonded together
They have weak Van Der Waals forces between the molecules
E.g h2 cl2 i2
39
Explain how covalent bonding occurs in terms of electronegativity?
When atoms of similar electronegativity are bonded together covalent bonding occurs and covalent substances are formed
40
Explain how ionic bonding occurs in terms of electronegativity?
When atoms with a large difference in electronegativity are bonded together ionic bonding occurs and ionic substances are formed
41
what is electronegativity?
-electron negativity is the power of an atom to withdraw electron density from a covalent bond
in other words
-the ability of an atom to attract the bonding pair of electrons towards itself
42
Non polar covalent bonds share the electron pair…?
- Equally
The electron cloud is symmetrical
43
Explain how polar covalent bonds are formed?
When atoms of different electronegativity are bonded together
44
What happens in polar covalent bonding?
The more electronegative atom withdraws the electron pair more strongly
45
The more electronegative atom is..?
The less electronegative atom has a..?
- slightly negatively charged
- slight positive charge
46
What creates the polar bond?
the charge separation creates and electric dipole and the bond is said to be polar
47
How can the polarity of a bond be measured?
I’m a unit called debye
48
The larger the electronegativity difference between the atoms the more…?
Polar the bond
49
Explain why hydrogen chloride is a polar molecule?
- the more electronegative chlorine has a greater share of the electron cloud so is slightly negative
- the less electronegative hydrogen has slightly less electron density so is slight positive
- the charge separation created an electric dipole so the bond is said to be polar
- the hydrogen-chlorine bond is polar so hydrogen chloride is a polar molecule
50
Explain why water is a polar molecule?
- The more electronegative atom attracts the electron pair more strongly than hydrogen
- so oxygen will be slightly negative and hydrogen will be slightly positive
51
What are symmetrical non-polar molecules?
These bonds cancel each other out and the molecules have no overall polarity
E.g ch4
52
What are the tree types of intermolecular forces in order of increasing strength
Van Der Waals forces
Permanent Dipole- dipole forces
Hydrogen bonding
53
What can Van Der waals forces also be known as…?
Temporary induced dipole-dipole attractions
54
what compounds in Van Der Waals forces always in?
Present between molecules of all simple covalent substances and non-polar molecules
55
What is a temporary dipole?
Draw what a temporary dipole looks like.
So electrons are constantly moving, and at one instant there may be an unequal distribution of the electron cloud this results in a temporary dipole
56
What causes a temporary induced dipole dipole molecule
The temporary dipole can induce a dipole on another neighbouring molecule
57
Explain what are temporary induced dipole-dipole attractions?
- there is an electrostatic force of attraction between the slightly negatively charged and slightly positively charged on a neighbouring molecule
- These electrostatic forces of attraction between the temporary dipole and the induced dipole are Van Der waals forces
58
Why do Van Der waals forces only operate for a short amount of time?
Because the electron cloud is constantly changing
These forces are constantly being switched on and off
59
Why do Van Der waals forces have low melting and boiling points
Weak Abba Der callas forces require only a small amount of energy to overcome the forces
60
The size of van Der waals forces increases as..?
- the size of the molecule increases (larger electron cloud)
- the points of contact between the molecules increases
61
What does the boiling point increase down the halogens?
- Going down a group the size of the halogen increase,
- the size of the the electron cloud increases
- amount of temporary electron cloud distortion increases
- stronger Van Der waals forces so more energy needed to overcome these force s
62
Is Van Der waals forces stronger or weaker in straight or benched chain isomers. Explain why?
Straight chain isomers have stronger forces than. Branched chain isomers have a lower boiling point than straight chain because Van Der waals forces are weaker molecules are more spherical so fewer points of contact
63
permanent dipole-dipole forces
64
Where does this type of bonding occur in? And why?
Polar bonding - pair of electrons in the bond are unequally shared/ both atoms have different electronegativity values
65
how are polar molecules attracted by permanent dipole dipole forces
The permanent negative end of one molecule is attracted to the permanent positive end of another molecule (electrostatic attraction)
66
Why do polar molecules have higher melting and boiling points than non - polar molecules?
Polar molecules have stronger permanent pole dipole intermolecular forces, greater amount of energy is needed to overcome these forces
67
What is hydrogen bonding?
Hydrogen bonding is the attraction between a lone pair of electrons on a very electronegative atom (O,N,F) and a nearby slightly positive hydrogen which is directly bonded to O,N, or F on a neighbouring molecule
68
Why can N, O, F form a hydrogen bond but not molecules such as chlorine?
Because they are small enough to approach the slightly positive hydrogen atom closely enough
69
draw hydrogen bonding in hydrogen fluoride and water
70
the chlorine atom has the same electronegativity as nitrogen, why can it not undergo hydrogen bonding?
the chlorine atom is too large to get close enough to the slightly positive hydrogen
71
Explain why does ice have an open structure?
- hydrogen bond in ice cause the water molecules to be slightly further apart than In liquid water
72
Explain why ice floats on water?
Hydrogen bonds cause water molecules to be slightly further apart
- ice = openstructure
- density of solid water is less than liquid water hence why ice floats on water
73
Explain the bonding and structure in ice?
- each oxygen is linked to 4 hydrogens ( 2 covalent bond and 2 hydrogen)
- in a 3D tetrahedral arrangement
- 2 hydrogen bonds per water molecule
74
Draw hydrogen bonding in an ice crystal 2 ways
Use whiteboard
75
Draw Hydrogen bonding in proteins
Use whiteboard
76
When hydrogen bonding occurs it gives 2 types of secondary structure?
Alpha helix or beta pleated sheet - these structures are very stable
77
are covalent or hydrogen bonds weaker?
Hydrogen bonds are weaker but they are the strongest type of IMF
78
Why does boiling points of hydrides increase as you go go down the group in the periodic table?
As you go down size of molecule increases in size so there is an increase in IMF
79
why is there stronger bonding in H2O than H2S?
because hydrogen bonding occurs between water molecules and only dipole-dipole forces exist between H2S
80
Metalic bonding
81
What is metalic bonding
Metallic bonding is the force of attraction between the delocalised electrons and the positive metal ions in a lattice
82
What does the giant matallic lattice consist of?
Close packed metal ions surrounded by delocalised electrons
83
What is meant by delocalised?
The metals outer shell electrons are mobile and free to move throughout the metal structure
84
Because of the delocalised electrons the metal atoms become..?
Ions or positive centres
85
Draw a sketch of metallic bonding
86
Why do metalic bonds have high melting and boiling points
Metalic bonds between the ions are strong.
So a large amount of energy is required to remove a metal atom from the attraction of the delocalised electrons
87
The strength of a metalic bond increases as..?
- the charge on the positive ion increases
- the size of the metal ion decreases
- the number of mobile electrons per atom increases
88
Why is sodium relatively weak?
Because each atom donated 1 electron to the sea
89
Why is the strength of potassium weaker than sodium?
Because the ions are larger ( electron sea has a bigger volume to cover so is less effective at holding the ions together)
90
Why is metalic bonding stronger in magnesium?
Because each atom has donated 2 electrons to the sea and each ion has a 2+ charge
91
Draw a magnesium sketch
92
Why are metals good conductors of electricity?
Because they have delocalised mobile electrons
93
When does the electrical conductivity of a metal increase?
As the number of mobile electrons per atom increases
94
What else is delocalised electrons responsible for?
A metals ability to reflect light
95
Metals are shiny they have a m…?
Metallic lustre
96
What are metals malleable?
Metals can be hammered into shape (malleable) and drawn out into wires (ductile) because the planes of the ions can slide over each other
97
How are ions packed together in a metalic crystal
The spherical ions are usually packed together as closely as possible
98
What are metal crystals usually made from?
Repeating units which are usually hexagonal
99
Magnesium is said to have a what crystal structure?
Hexagonal close-packed crystal
Each ion has 6 nearest neighbours in the same plane, 3 above plane, 3 below plane. Therefore each ion has 12 nearest neighbours = coordination number
100
Shapes of molecules
101
If the electron pair has 2 electrons in its outer shell what shape does it form?
A linear shape - 180
102
If the number of electrons pairs is 3 what shape does it form
Trigonal planar - 120
103
If the number of electron pairs on the outer shell is 4 what shape is formed?
Tetrahedral - 109.5
104
If the number of electron pairs is 5 what shape is formed?
Trigonal by-pyramid - 90 and 120
105
If the number of electron pairs is 6 what shape is formed?
Octahedral - 90
106
What is the order of repulsion?
- Lone pair - lone pair = most
- Lone air- bonding pair = second
- bonding - pair - bonding pair = least
107
why do lone pairs repel more strongly then bonding pairs?
lone pairs are held more closely to the nucleus and therefore more compact
108
How do we calculate the number of electrons in the outer shell by the molecular formula?
- add the group of the central atom and the number of atoms bonded to the central atom
109
When do bonded pairs not repel each other equally?
When all the bond angles are not the same
E.g PBr5 has bond angles 120 and 90
110
Shapes of molecules with lone pairs
111
What shape do we get when there is one lone pair on a Trigonal planar based shape?
Bent/ angular shape - 120
Lone pair at the top
112
Tetrahedral shape when:
1 lone pair?
2 lone pairs?
1 lone pair - trigonal pyramid - 109
2 lone pairs - bent / angular
113
Trigonal bupyramidal shape when:
1 lone pair?
2 lone pairs?
3 lone pairs?
1 lone pair = seesaw- 120 and 90
2 lone pairs - T-shape 90
3 lone pairs = linear 180
114
Octahedral shape when there is:
One lone pair?
Two lone pairs?
Three lone pairs?
4 lone pairs?
1 lone pair = square pyramid - 90
2 lone pairs = square planar - 90
3 lone pairs = T-shape 90
4 lone pairs = linear 180
115
Each lone pair decreases the bond angle by…?
2.5 degrees
116
If an ion has a 1+ charge it has..?
If an ion has a 1- charge it has..?
- lost 1 electron
- gained one electron
117
DONEE!!1
