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3.1 Physical Chemistry > 3.1.3 Bonding > Flashcards

3.1.3 Bonding Flashcards

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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
Q

Molecular Shapes: 2 bonding pairs, 2 lone pairs

A

Bent, 104°

26
Q

Molecular Shapes: 5 electron pairs

A

Trigonal Bipyramidal, 120°&90°

27
Q

Molecular Shapes: 3 bonding pairs, 2 lone pairs

A

Trigonal Planar, 120° OR T-Shape 89°

28
Q

Molecular Shapes: 6 electron pairs

A

Octahedral, 90°

29
Q

Molecular Shapes: 5 bonding pairs, 1 lone pair

A

Distorted square pyramid, 89°

30
Q

Molecular Shapes: 4 bonding pairs, 2 lone pairs

A

Square planar, 90°

31
Q

How to work out shapes

A

-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
Q

Intermolecular Forces: Van Der Waals Forces

A

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
Q

Factors which affect the strength of VDW forces: Number of

electrons in the molecule

A

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
Q

Factors which affect the strength of VDW forces: Surface

area of the molecules

A

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
Q

Intermolecular Forces: Permanent Dipole-Dipole

A

Occur between polar molecules

Stronger than VDW forces, so boiling points are higher

36
Q

Intermolecular Forces: Hydrogen Bonding

A

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
Q

Low density of ice

A

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
Q

Helical nature of DNA

A

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
Q

Molecular Structures (Covalent)

A

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
Q

Macromolecular Structures (Covalent)

A

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
Q

Giant Covalent Layered (Covalent)

A

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
Q

Ionic Structures

A

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
Q

Metallic Structures

A

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
Q

Why are metals malleable?

A

Because the layers of cations can easily slip over each other

45
Q

Solids

A

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
Q

Liquids

A

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
Q

Gases

A

All particles are in rapid and random motion and behave

3.1 Physical Chemistry (5 decks)

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