3.1.3 Bonding Flashcards
Ionic Bonding
The electrostatic force of attraction between two oppositely
charged ions formed by electron transfer
Covalent Bonding
A shared pair of electrons
Dative Covalent Bonding (AKA Co-ordinate bonding)
Formed when the shared pair of electrons in the covalent
bond come from only one of the bonding atoms.
Metallic Bonding
The electrostatic force of attraction between the positive
metal cations and the sea of delocalised electrons
Factors affecting the strength of metallic bonding: The
number of protons
The more protons in the cations, the stronger the electrostatic
force of attraction between the cations and the sea of
delocalised electrons
Factors affecting the strength of metallic bonding: Number
of delocalised electrons per atom
The more delocalised electrons, the stronger the electrostatic
force of attraction
Factors affecting the strength of metallic bonding: Size of ion
The smaller the ion, the stronger the electrostatic force of
attraction
Electronegativity
The relative tendency of an atom in a covalent bond in a
molecule to attract electrons in a covalent bond towards itself
Why does electronegativity increase as you go across a
period?
- The number of protons increased
- The atomic radius decreases because the electrons in the same shell are pulled in more
Why does electronegativity decrease as you go down a
group?
-Distance between the nucleus and the outer electrons
increases
-Shielding increases
Why aren’t the noble gases electronegative?
Because they don’t form bonds
Using electronegativity to predict bonding: Covalent
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
Using electronegativity to predict bonding: Ionic
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
Using electronegativity to predict bonding: Metallic
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
Orbitals & Covalent Bonds
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
Orbitals and Dative Covalent Bonds
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.
Sigma Bonds
Where the atomic orbitals overlap directly along the internuclear axis. All single bonds are sigma bonds.
Pi Bonds
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
Strength of covalent bonds is affected when..
The atoms are smaller because the closer the electrons are to
the nuclei, the stronger the bond
Molecular Shapes: 2 electron pairs
Linear, 180°
Molecular Shapes: 3 electron pairs
Trigonal Planar, 120°
Molecular Shapes: 2 bonding pairs, 1 lone pair
Bent, 118°
Molecular Shapes: 4 electron pairs
Tetrahedral, 109.5°
Molecular Shapes: 3 bonding pairs, 1 lone pair
Trigonal Pyramidal, 107°
Molecular Shapes: 2 bonding pairs, 2 lone pairs
Bent, 104°
Molecular Shapes: 5 electron pairs
Trigonal Bipyramidal, 120°&90°
Molecular Shapes: 3 bonding pairs, 2 lone pairs
Trigonal Planar, 120° OR T-Shape 89°
Molecular Shapes: 6 electron pairs
Octahedral, 90°
Molecular Shapes: 5 bonding pairs, 1 lone pair
Distorted square pyramid, 89°
Molecular Shapes: 4 bonding pairs, 2 lone pairs
Square planar, 90°
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
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
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.
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
Intermolecular Forces: Permanent Dipole-Dipole
Occur between polar molecules
Stronger than VDW forces, so boiling points are higher
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
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
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
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
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
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
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
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
Why are metals malleable?
Because the layers of cations can easily slip over each other
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
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
Gases
All particles are in rapid and random motion and behave
