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3.1 Physical Chemistry > 3.1.3 Bonding > Flashcards

3.1.3 Bonding Flashcards

1
Q

Ionic Bonding

A

The electrostatic force of attraction between two oppositely

charged ions formed by electron transfer

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2
Q

Covalent Bonding

A

A shared pair of electrons

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3
Q

Dative Covalent Bonding (AKA Co-ordinate bonding)

A

Formed when the shared pair of electrons in the covalent

bond come from only one of the bonding atoms.

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4
Q

Metallic Bonding

A

The electrostatic force of attraction between the positive

metal cations and the sea of delocalised electrons

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5
Q

Factors affecting the strength of metallic bonding: The

number of protons

A

The more protons in the cations, the stronger the electrostatic
force of attraction between the cations and the sea of
delocalised electrons

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6
Q

Factors affecting the strength of metallic bonding: Number

of delocalised electrons per atom

A

The more delocalised electrons, the stronger the electrostatic
force of attraction

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7
Q

Factors affecting the strength of metallic bonding: Size of ion

A

The smaller the ion, the stronger the electrostatic force of

attraction

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8
Q

Electronegativity

A

The relative tendency of an atom in a covalent bond in a

molecule to attract electrons in a covalent bond towards itself

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9
Q

Why does electronegativity increase as you go across a

period?

A
  • The number of protons increased

- The atomic radius decreases because the electrons in the same shell are pulled in more

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10
Q

Why does electronegativity decrease as you go down a

group?

A

-Distance between the nucleus and the outer electrons
increases
-Shielding increases

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11
Q

Why aren’t the noble gases electronegative?

A

Because they don’t form bonds

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12
Q

Using electronegativity to predict bonding: Covalent

A

If both atoms have a similar electronegativity, the pull on the electrons from them will be of a similar strength, making a non-polar covalent bond.
If one atom has a stronger electronegativity than the other, the electrons will be pulled more towards one atom, making the bond polar-covalent

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13
Q

Using electronegativity to predict bonding: Ionic

A

If the electronegativity difference is really large, the sharing of electrons is so uneven that the more electronegative atom has full possession of the 2 electrons, creating an ionic bond

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14
Q

Using electronegativity to predict bonding: Metallic

A

If both atoms have a low electronegativity, neither can attract
electrons, so the electrons don’t remain localised to the bond
at all, causing a sea of delocalised electrons and a metallic
bond

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15
Q

Orbitals & Covalent Bonds

A

When a covalent bond is formed, the 2 outer orbitals overlap, forming a normal covalent bond.
Some atoms promote electrons to give more unpaired
electrons and to allow more covalent bonding. For example, carbon promotes one of the electrons in the 2s orbital to the 2p orbital, meaning there are 4 unpaired electrons, so it can form 4 covalent bonds

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16
Q

Orbitals and Dative Covalent Bonds

A

Any atom with filled valence shell (outer shell) orbitals can donate their electrons for the covalent bond. This includes group 5,6,7 and 0
Any atom which has an empty orbital in their valence shell can accept a pair of electrons.

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17
Q

Sigma Bonds

A

Where the atomic orbitals overlap directly along the internuclear axis. All single bonds are sigma bonds.

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18
Q

Pi Bonds

A

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

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19
Q

Strength of covalent bonds is affected when..

A

The atoms are smaller because the closer the electrons are to
the nuclei, the stronger the bond

20
Q

Molecular Shapes: 2 electron pairs

A

Linear, 180°

21
Q

Molecular Shapes: 3 electron pairs

A

Trigonal Planar, 120°

22
Q

Molecular Shapes: 2 bonding pairs, 1 lone pair

A

Bent, 118°

23
Q

Molecular Shapes: 4 electron pairs

A

Tetrahedral, 109.5°

24
Q

Molecular Shapes: 3 bonding pairs, 1 lone pair

A

Trigonal Pyramidal, 107°

25
Molecular Shapes: 2 bonding pairs, 2 lone pairs
Bent, 104°
26
Molecular Shapes: 5 electron pairs
Trigonal Bipyramidal, 120°&90°
27
Molecular Shapes: 3 bonding pairs, 2 lone pairs
Trigonal Planar, 120° OR T-Shape 89°
28
Molecular Shapes: 6 electron pairs
Octahedral, 90°
29
Molecular Shapes: 5 bonding pairs, 1 lone pair
Distorted square pyramid, 89°
30
Molecular Shapes: 4 bonding pairs, 2 lone pairs
Square planar, 90°
31
How to work out shapes
-Write group number of central atom -Add number of atoms around the central atom -Add/Subtract charge (If charge is negative, add it & if charge is positive, subtract it) -To find bonding pairs, divide step 3 by 2 -To find lone pair, subtract step 4 by the number of atoms around the central atoms
32
Intermolecular Forces: Van Der Waals Forces
Occur between all simple covalent molecules & the separate atoms in noble gases As the electrons move, parts of the molecule can become more or less electronegative. These dipoles can induce dipoles in adjacent molecules
33
Factors which affect the strength of VDW forces: Number of | electrons in the molecule
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.
34
Factors which affect the strength of VDW forces: Surface | area of the molecules
The larger the surface area of a molecule, the more contact it will have with adjacent molecules. This means it has a greater ability to induce a dipole in an adjacent molecule and the VDW forces are stronger, and melting points and boiling points will be higher. Straight chain isomers have a larger S.A than branched isomers
35
Intermolecular Forces: Permanent Dipole-Dipole
Occur between polar molecules | Stronger than VDW forces, so boiling points are higher
36
Intermolecular Forces: Hydrogen Bonding
Occurs in compounds which have H attached to F, O, N which must have an available lone pair of electrons. There is a large difference between electronegativities They have the highest boiling points as they're the strongest intermolecular force
37
Low density of ice
Due to hydrogen bonding, the water molecules arrange themselves to maximise the hydrogen bonding between the molecules, resulting in an open hexagonal structure with large spaces within the crystal. This means it has a low density. When ice melts, the structure collapses into the open spaces and the resulting liquid occupies less space and is denser
38
Helical nature of DNA
Due to hydrogen bonding. Molecules of DNA contain N-H bonds and C=O bonds, meaning hydrogen bonding is possible between N-H and H-O, resulting in the shape spiralling
39
Molecular Structures (Covalent)
Substance is made up of molecules with no strong bonding between them Molecules held together by intermolecular forces Melting & boiling points are low because intermolecular forces are weak. Most substances are soft & crumbly. There's little electrical conductivity in any state as there are no ions and no delocalised electrons
40
Macromolecular Structures (Covalent)
Where the bonding capacity isn't satisfied, so a lattice is formed C, B, Si & SiO2 are examples There are covalent bonds between adjacent atoms, so melting points and boiling points are high. There's little electrical conductivity because there are no ions or delocalised electrons. Most substances are hard, stong and brittle
41
Giant Covalent Layered (Covalent)
Graphite is an example Lattices form layers, held together by intermolecular forces There is electrical conductivity as there are delocalised electrons in each plane. Graphite has a lower density as there's large distances between each plane. It' softer than diamond because the planes can slip over each other
42
Ionic Structures
After ions are formed, they come together to form a lattice. All anions are surrounded by cations and cations are surrounded by anions. Melting and boiling points are high because the attraction between opposite charges are very strong, the higher the charge, the harder it will be to break the attraction. Ionic structures can't conduct electricity in solid state as the ions aren't free to move. In liquid state, however, they're free to move, so electrical conductivity is possible. All ionic compounds are hard and brittle
43
Metallic Structures
Cations are arranged to form a lattice and the electrons are free-flowing. Electricity conductivity is possible in both solid and liquid states as the electrons are delocalised and free to move. The bonding in metals is relatively high so the melting and boiling points are high. Metals are malleable and ductile
44
Why are metals malleable?
Because the layers of cations can easily slip over each other
45
Solids
Particles are tightly packed into a lattice. A solid has a fixed volume and a fixed state. The kinetic energy that the particles have isn't enough to fling them apart, so they're restricted to rotational and vibrational motion
46
Liquids
The particles are packed into a lattice, however the structure breaks down so there's enough space for particles to move from one cluster to another. They have a fixed volume but no fixed shape. The kinetic energy they have allows translational motions
47
Gases
All particles are in rapid and random motion and behave

3.1 Physical Chemistry (5 decks)

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