Chemistry - 3.1.3: Bonding

Ionic bond

A type of chemical bond that occurs between metal and non-metal atoms where the metal transfers electrons to the non-metal, forming charged ions. These oppositely charged ions are then attracted to each other by electrostatic forces of attraction.

Why do the metal atoms lose / non-metal atoms gain electrons?

To form ions with a full outer shell.

Formula of a:
1. Sulfate ion
2. Hydroxide ion
3. Nitrate ion
4. Carbonate ion
5. Ammonium ion

1. SO4 2-
2. OH -
3. NO3 -
4. CO3 2-
5. NH4 +

Covalent bond

A type of chemical bond that forms between two non-metal atoms when they share a pair(s) of electrons to have a full outer shell.

Dative / Co-ordinate covalent bond

A type of covalent bond that forms when both of the electrons in the shared pair are supplied from one atom.

How do you represent a dative bond?

An arrow

Metallic bonds

A chemical bond formed due to the electrostatic forces of attraction between the positively charged layers of metal ions and the negatively charged sea of delocalised electrons.

Effect of increasing charge on metallic bonds

Greater strength of the electrostatic forces of attraction (more electrons released in sea)

Effect of larger ions on metallic bonds

Weaker strength of the electrostatic forces of attraction due to greater atomic radius

4 types of crystal structure

1. Ionic
2. Metallic
3. Simple molecular
4. Macromolecular

3 features of giant ionic lattices

1. High MP/BP - strong electrostatic forces of attraction between oppositely charged ions require a lot of energy to break
2. Conduct electricity when molten/aqueous - ions are mobile (free to move) and can carry the charge [when solid, in a fixed position so can't]
3. Brittle - when broken and layers of alternating charges distort, like charges repel, leading to fragmentation

3 features of metallic structures

1. High MP/BP - strong electrostatic forces of attraction between positively charged ion layers and negatively charged delocalised electrons require a lot of energy to break
2. Conduct electricity - delocalised electrons are mobile (free to move) and can carry the charge
3. Malleable - layers of positive ions can slide over each other (delocalised electrons prevent fragmentation)

2 features of simple molecular substances

1. Low MP/BP - weak intermolecular forces do not require a lot of energy to break
2. Do not conduct electricity - no charged particles that are free to move so can't carry the charge

3 features of diamond

1. Rigid - each carbon atom forms a covalent bond with 4 other carbon atoms
2. High MP/BP - strong covalent bonds require a lot of energy to break
3. Do not conduct electricity - no charged particles that are free to move so can't carry the charge

3 features of graphite

1. Malleable - made up of flat sheets/layers which can slide over each other
2. High MP/BP (but less than diamond) - strong covalent bonds require a lot of energy to break
2. Conducts electricity - each carbon atom only forms 3 covalent bonds, so there are delocalised electrons which are free to move and carry the charge

Electron charge cloud

A region where there is a high chance of an electron pair being present (region of negative charge around the nucleus) - can be a bonding pair or a lone pair

Why are electron charge clouds as far away as possible from each other?

To minimise repulsion - they are like charges

Strength of repulsion / size of bond angle between pairs of electrons (from most to least)

Lone pair - Lone pair
Lone pair - Bond pair
Bond pair - Bond pair

Why is the bond angle between two lone pairs the greatest?

They have the greatest repulsion force (held closer to central atom than bond pairs + their clouds are more concentrated and wider)

VSEPR Theory

Valence-shell electron-pair repulsion theory: The number of electron pairs on the outermost shell of a central atom determines the shape of the molecule or ion. Electron pairs will be pushed apart by repelling forces till these are minimised, affecting bond angles and molecular shape.

8 Shapes of Molecules

Linear, Bent, Trigonal planar, Tetrahedral, Trigonal bipyramidal, Trigonal pyramidal, Octahedral, Square Planar

Linear

2 Bond pairs, 0 Lone pairs, 180°

Trigonal planar

3 Bond pairs, 0 Lone pairs, 120°

Tetrahedral

4 Bond pairs, 0 Lone pairs, 109.5°

Trigonal bi-pyramidal

5 Bond pairs, 0 Lone pairs, 90° between bonds in plane of sight, 120° between bonds not in it

Octahedral

6 Bond pairs, 0 Lone pairs, 90°

When drawing molecules, what does a dashed line and dark triangle mean?

Dashed line - going away from you
Dark triangle - going towards you

Trigonal pyramidal

3 Bond pairs, 1 Lone pair (4 total), 107°

Bent

2 Bond pairs, 2 Lone pairs (4 total), 104.5°

Electronegativity

The power of an atom to attract the pair of electrons in a covalent bond

Why do some molecules (e.g. Cl2) have no dipoles?

There is no difference in electronegativity between the two atoms.

Why do some molecules have no overall dipole moment?

The individual dipoles cancel out due to the 3D shape of the molecule

3 types of intermolecular forces (in order of increasing strength)

Induced dipole-dipole forces
Permanent dipole-dipole forces
Hydrogen bonding

How do you draw hydrogen bonds?

Dashed lines

When do hydrogen bonds form?

When fluorine / oxygen / nitrogen atoms are bonded covalently to hydrogen atoms - the very electronegative atom MUST have a lone pair of electrons to attract the hydrogen atom

How do permanent dipole-dipole forces form? (3)

Difference in electronegativity leads to bond polarity. Positive and negative dipoles form (name them) - positive dipole on one molecule attracts negative dipole on another molecule.

How do induced dipole-dipole forces form?

(Electrons in all molecules are in constant motion at high speeds so) electrons are unevenly distributed - temporary dipoles.
This dipole induces another temporary dipole in a neighbouring molecule.
Attraction between these dipoles = induced dipole-dipole force

What does the size of the induced dipole-dipole forces depend on?

Number of electrons in a molecule + the size of the area of contact of one molecule with another

Where are induced dipole-dipole forces found?

All molecules

Where are permanent dipole-dipole forces found?

In polar molecules

Square Planar

4 Bond pairs, 2 Lone pairs (6 total), 90°

2 factors that affect the strength of metallic bonds

The charge of the ion (greater charge, greater number of delocalised electrons, stronger bonds) and the atomic radius [size] of the ion (smaller ions, electrons closer to nucleus, stronger bonds)

3 factors that affect electronegativity

Nuclear charge, atomic radius and shielding

Most electronegative elements

Fluorine, oxygen, nitrogen (chlorine is somewhat electronegative but not enough to form hydrogen bonds)

Effect of increasing the number of electrons on induced dipole-dipole forces

Increasing the number of electrons increases the size of the induced dipole-dipole forces

Why is ice less dense than water?

As a liquid, hydrogen bonds between water molecules easily break and reform. As a solid, hydrogen bonds hold molecules in place, slightly further away than in water. Fewer molecules in a given volume = less dense.

3 properties of solids + evidence for it

1. Regular arrangement: crystal shapes have straight edges + solids have definite shapes
2. Particles are close together: not easily compressed
3. Vibrate around a fixed position: diffusion is very slow + solids expand when heated

3 properties of liquids + evidence for it

1. Random arrangement: liquids change shape to fill the bottom of their container
2. Particles are close together: not easily compressed
3. Rapid 'jostling': diffusion is very slow + liquids evaporate

3 properties of gases + evidence for it

1. Random arrangement: gases fill their container
2. Far apart: gases are easily compressed
3. Fast moving: diffusion is rapid + gases exert pressure

What happens when you heat a substance?

The particles have more kinetic energy, if solid: vibrate more about a fixed position; if liquid or gas: move faster - the average distance between each particle increases, causing the substance to expand

What happens when a solid melts into a liquid?

The energy from the heat is used to weaken the intermolecular forces that act between the particles - no change in temperature

What happens when a liquid evaporates into a gas?

The energy is used to break the intermolecular forces that act between the particles

What is the bond in the hydrogen bonds between?

The lone pair of electrons and the hydrogen atom

In a molecule, how do the lone pair of electrons on the central atom affect the shape of the molecule?

Lone pairs of electrons repel more than bond pairs - leading to bond angle

How to answer why a certain substance has a high/low MP/BP? (3)

1. State type of structure
2. State bonds present and what particles they are between
3. State whether they need a lot / a little energy to be overcome

Give 3 tests for group 2 ions

Adding ammonia solution, adding excess NaOH and adding excess H2SO4

Do double and triple bonds repel more or less than single bonds?

They repel equally as single bonds

How is ammonia solution used to test for group 2 ions?

Mg2+ forms white precipitate of Mg(OH)2, Ca2+ doesn't form a precipitate

How is NaOH used to test for group 2 ions?

Mg2+ forms white precipitate of Mg(OH)2, Ca2+ forms white precipitate of Ca(OH)2, Sr2+ forms slight white precipitate of Sr(OH)2

How is sulfuric acid used to test for group 2 ions?

Mg2+ forms colourless solution of MgSO4, Ca2+ forms slight white precipitate of CaSO3, Sr2+ and Ba2+ forms white precipitate of SrSO4 and BaSO4 respectively

What happens if all the bonds in a molecule are bond pairs?

The electrons repel each other equally

Why is XeF4 square planar? (3)

1. It has 4 bonding pairs and 2 lone pairs of electrons
2. The lone pairs repel each other more than bonding pairs, so move to be as far apart as possible / opposite each other
3. This results in a bond angle of 90°

Why might a molecule with hydrogen bonds be easy to liquify?

The hydrogen bonds are a strong enough force to hold the molecules together in a liquid