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Chemistry Y12 > Bonding > Flashcards

Bonding Flashcards

1
Q

What is ionic bonding?

A

The force of attraction between two oppositely charged ions

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2
Q

Metals loose electrons to form..?
Non- metals gain electrons to form?

A
  • positive ions (cations)
  • negative ions (anions)
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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

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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

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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
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6
Q

The electrostatic forces increase in strength as ..?

A
  • the charge on the ion increases
  • the size of the ion decreases
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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
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8
Q

What is a crystal?

A

A solid with a regular shape which contains particles organised in a regular structure

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9
Q

Ionic bonds in solids form..?

A

Giant ionic lattice

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10
Q

How are ions arranged in a giant ionic lattice?

A

in a regular repeating pattern

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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.
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12
Q

the type of ionic lattice formed by an ionic compound depends on what..?

A

the relative sizes of the ions present

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13
Q

What does Caesium chloride crystal form?

A

forms a body-centered cubic crystal instead of a face centered cubic crystal

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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
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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.

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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
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17
Q

What determines the strength of a covalent bond?

A

The amount of overlap of the atomic orbitals
greater overlap = stronger bond

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18
Q

What are lone pairs?

A

Non-bonded pairs of electrons in a compound

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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
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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

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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

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22
Q

why do the arrows show in the Co-ordinate/ dative covalent bond?

A

shows the origin of the lone pair of electrons

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23
Q

Two types of covalent structures?

A

Simple molecular and giant macromolecular

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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

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25
Q

Describe and draw the structure of an iodine crystal?

A
  • 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
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26
Q

What happens when iodine is heated?

A

Changes from a black shiny solid into a purple vapour (sublimation) which separated the I2 molecules

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27
Q

Two examples that form giant covalent crystals?

A

Diamond and graphite which are allotropes of carbon

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28
Q

Sketch and Describe the structure and properties of diamond

A
  • giant covalent structure - strongly covalently bonded to 4 other carbon atoms
  • tetrahedral arrangement
  • does not conduct - no free electrons
  • hard
  • strong melting and boiling
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29
Q

Describe the structure and properties of graphite?

A
  • 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
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30
Q

what is electronegativity?

A

-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

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31
Q

Why are there no electronegativity values for group 0?

A

Because they do not form covalent bonds

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32
Q

What are the trends in electronegativity on the periodic table?

A
  • electronegativity decreases down a group
  • electronegativity increases across a period
  • generally, increases from bottom left to the top right of the periodic table
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33
Q

What can be used to explain the electronegativity trends?

A
  • size of the atom
  • size of the nuclear charge
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34
Q

Why does electronegativity decrease down a group?

A
  • 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
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35
Q

Explain why does electronegativity increase across a period?

A
  • 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
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36
Q

Which elements have the highest electronegativity values?

A

Nof elements
Nitrogen, oxygen, fluorine

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37
Q

What elements have the lowest electronegativity values?

A

Bottom left of periodic table (reactive metals)
Caesium, rubidium

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38
Q

How are non-polar molecules formed in terms of electronegativity ?
Forced between them are..?

A

When atoms of the same electronegativity are bonded together

They have weak Van Der Waals forces between the molecules
E.g h2 cl2 i2

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39
Q

Explain how covalent bonding occurs in terms of electronegativity?

A

When atoms of similar electronegativity are bonded together covalent bonding occurs and covalent substances are formed

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40
Q

Explain how ionic bonding occurs in terms of electronegativity?

A

When atoms with a large difference in electronegativity are bonded together ionic bonding occurs and ionic substances are formed

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41
Q

what is electronegativity?

A

-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

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42
Q

Non polar covalent bonds share the electron pair…?

A
  • Equally
    The electron cloud is symmetrical
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43
Q

Explain how polar covalent bonds are formed?

A

When atoms of different electronegativity are bonded together

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44
Q

What happens in polar covalent bonding?

A

The more electronegative atom withdraws the electron pair more strongly

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45
Q

The more electronegative atom is..?
The less electronegative atom has a..?

A
  • slightly negatively charged
  • slight positive charge
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46
Q

What creates the polar bond?

A

the charge separation creates and electric dipole and the bond is said to be polar

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47
Q

How can the polarity of a bond be measured?

A

I’m a unit called debye

48
Q

The larger the electronegativity difference between the atoms the more…?

A

Polar the bond

49
Q

Explain why hydrogen chloride is a polar molecule?

A
  • 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
Q

Explain why water is a polar molecule?

A
  • 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
Q

What are symmetrical non-polar molecules?

A

These bonds cancel each other out and the molecules have no overall polarity
E.g ch4

52
Q

What are the tree types of intermolecular forces in order of increasing strength

A

Van Der Waals forces
Permanent Dipole- dipole forces
Hydrogen bonding

53
Q

What can Van Der waals forces also be known as…?

A

Temporary induced dipole-dipole attractions

54
Q

what compounds in Van Der Waals forces always in?

A

Present between molecules of all simple covalent substances and non-polar molecules

55
Q

What is a temporary dipole?
Draw what a temporary dipole looks like.

A

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
Q

What causes a temporary induced dipole dipole molecule

A

The temporary dipole can induce a dipole on another neighbouring molecule

57
Q

Explain what are temporary induced dipole-dipole attractions?

A
  • 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
Q

Why do Van Der waals forces only operate for a short amount of time?

A

Because the electron cloud is constantly changing
These forces are constantly being switched on and off

59
Q

Why do Van Der waals forces have low melting and boiling points

A

Weak Abba Der callas forces require only a small amount of energy to overcome the forces

60
Q

The size of van Der waals forces increases as..?

A
  • the size of the molecule increases (larger electron cloud)
  • the points of contact between the molecules increases
61
Q

What does the boiling point increase down the halogens?

A
  • 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
Q

Is Van Der waals forces stronger or weaker in straight or benched chain isomers. Explain why?

A

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
Q

permanent dipole-dipole forces

A
64
Q

Where does this type of bonding occur in? And why?

A

Polar bonding - pair of electrons in the bond are unequally shared/ both atoms have different electronegativity values

65
Q

how are polar molecules attracted by permanent dipole dipole forces

A

The permanent negative end of one molecule is attracted to the permanent positive end of another molecule (electrostatic attraction)

66
Q

Why do polar molecules have higher melting and boiling points than non - polar molecules?

A

Polar molecules have stronger permanent pole dipole intermolecular forces, greater amount of energy is needed to overcome these forces

67
Q

What is hydrogen bonding?

A

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
Q

Why can N, O, F form a hydrogen bond but not molecules such as chlorine?

A

Because they are small enough to approach the slightly positive hydrogen atom closely enough

69
Q

draw hydrogen bonding in hydrogen fluoride and water

A
70
Q

the chlorine atom has the same electronegativity as nitrogen, why can it not undergo hydrogen bonding?

A

the chlorine atom is too large to get close enough to the slightly positive hydrogen

71
Q

Explain why does ice have an open structure?

A
  • hydrogen bond in ice cause the water molecules to be slightly further apart than In liquid water
72
Q

Explain why ice floats on water?

A

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
Q

Explain the bonding and structure in ice?

A
  • 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
Q

Draw hydrogen bonding in an ice crystal 2 ways

A

Use whiteboard

75
Q

Draw Hydrogen bonding in proteins

A

Use whiteboard

76
Q

When hydrogen bonding occurs it gives 2 types of secondary structure?

A

Alpha helix or beta pleated sheet - these structures are very stable

77
Q

are covalent or hydrogen bonds weaker?

A

Hydrogen bonds are weaker but they are the strongest type of IMF

78
Q

Why does boiling points of hydrides increase as you go go down the group in the periodic table?

A

As you go down size of molecule increases in size so there is an increase in IMF

79
Q

why is there stronger bonding in H2O than H2S?

A

because hydrogen bonding occurs between water molecules and only dipole-dipole forces exist between H2S

80
Q

Metalic bonding

A
81
Q

What is metalic bonding

A

Metallic bonding is the force of attraction between the delocalised electrons and the positive metal ions in a lattice

82
Q

What does the giant matallic lattice consist of?

A

Close packed metal ions surrounded by delocalised electrons

83
Q

What is meant by delocalised?

A

The metals outer shell electrons are mobile and free to move throughout the metal structure

84
Q

Because of the delocalised electrons the metal atoms become..?

A

Ions or positive centres

85
Q

Draw a sketch of metallic bonding

A
86
Q

Why do metalic bonds have high melting and boiling points

A

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
Q

The strength of a metalic bond increases as..?

A
  • the charge on the positive ion increases
  • the size of the metal ion decreases
  • the number of mobile electrons per atom increases
88
Q

Why is sodium relatively weak?

A

Because each atom donated 1 electron to the sea

89
Q

Why is the strength of potassium weaker than sodium?

A

Because the ions are larger ( electron sea has a bigger volume to cover so is less effective at holding the ions together)

90
Q

Why is metalic bonding stronger in magnesium?

A

Because each atom has donated 2 electrons to the sea and each ion has a 2+ charge

91
Q

Draw a magnesium sketch

A
92
Q

Why are metals good conductors of electricity?

A

Because they have delocalised mobile electrons

93
Q

When does the electrical conductivity of a metal increase?

A

As the number of mobile electrons per atom increases

94
Q

What else is delocalised electrons responsible for?

A

A metals ability to reflect light

95
Q

Metals are shiny they have a m…?

A

Metallic lustre

96
Q

What are metals malleable?

A

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
Q

How are ions packed together in a metalic crystal

A

The spherical ions are usually packed together as closely as possible

98
Q

What are metal crystals usually made from?

A

Repeating units which are usually hexagonal

99
Q

Magnesium is said to have a what crystal structure?

A

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
Q

Shapes of molecules

A
101
Q

If the electron pair has 2 electrons in its outer shell what shape does it form?

A

A linear shape - 180

102
Q

If the number of electrons pairs is 3 what shape does it form

A

Trigonal planar - 120

103
Q

If the number of electron pairs on the outer shell is 4 what shape is formed?

A

Tetrahedral - 109.5

104
Q

If the number of electron pairs is 5 what shape is formed?

A

Trigonal by-pyramid - 90 and 120

105
Q

If the number of electron pairs is 6 what shape is formed?

A

Octahedral - 90

106
Q

What is the order of repulsion?

A
  • Lone pair - lone pair = most
  • Lone air- bonding pair = second
  • bonding - pair - bonding pair = least
107
Q

why do lone pairs repel more strongly then bonding pairs?

A

lone pairs are held more closely to the nucleus and therefore more compact

108
Q

How do we calculate the number of electrons in the outer shell by the molecular formula?

A
  • add the group of the central atom and the number of atoms bonded to the central atom
109
Q

When do bonded pairs not repel each other equally?

A

When all the bond angles are not the same
E.g PBr5 has bond angles 120 and 90

110
Q

Shapes of molecules with lone pairs

A
111
Q

What shape do we get when there is one lone pair on a Trigonal planar based shape?

A

Bent/ angular shape - 120
Lone pair at the top

112
Q

Tetrahedral shape when:
1 lone pair?
2 lone pairs?

A

1 lone pair - trigonal pyramid - 109
2 lone pairs - bent / angular

113
Q

Trigonal bupyramidal shape when:
1 lone pair?
2 lone pairs?
3 lone pairs?

A

1 lone pair = seesaw- 120 and 90
2 lone pairs - T-shape 90
3 lone pairs = linear 180

114
Q

Octahedral shape when there is:
One lone pair?
Two lone pairs?
Three lone pairs?
4 lone pairs?

A

1 lone pair = square pyramid - 90
2 lone pairs = square planar - 90
3 lone pairs = T-shape 90
4 lone pairs = linear 180

115
Q

Each lone pair decreases the bond angle by…?

A

2.5 degrees

116
Q

If an ion has a 1+ charge it has..?
If an ion has a 1- charge it has..?

A
  • lost 1 electron
  • gained one electron
117
Q

DONEE!!1

A

Chemistry Y12 (11 decks)

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