Bonds and bonding

1. Ionic

2. Covalent

3. Metallic

• Electrons are transferred from one atom to another.

• Compounds are made.• Involves a metal and a nonmetal

– Metal : donates electron(s) so it becomes a positive ion (cation)

– Nonmetal receives electron and becomes a negative ion (anion)

– So...must have a cation and anion

Why do ions exist?

• Atoms want a full outer shell of valence electrons (8).

• Will gain or lose electrons to achieve this = Octet rule

• Then, positive and negative ion attract (bond)

• Alternating positive and negative ions form crystalline lattice structures

• The number and size of ion determines the shape of the crystal

Sodium chloride: NaCl

Calcium fluoride: CaF2

Quartz: SiO2

Alum

Properties of Ionic Compounds

All properties are a result of the arrangement of cation and anions.

1. Solid, crystalline

2. Hard and brittle– When stuck, the crystal will break along a

line of symmetry (fault line)

3. Extremely high melting and boiling points.

– Each molecule is strongly attracted to its neighboring molecule (strong intermolecular attraction)

Remember…

• The bonds are not broken during melting and boiling. The molecules are separated from one another.

4. Readily dissolves in water– The polar water molecule attracts and pulls

the cations and anions apart.

5. Non-conductive as a solid, but is conductive when dissolved in solution.

– Solid is neutral (cations and anions equal)– When dissolved exists as cations and

anions, which carry a charge.

What ions exist in our water supply?

• Iron– Causes rust spots, dingy clothes

• Copper from ground and pipes– Causes green residue

• Chlorine and fluorine added to city water• Bad heavy metals: mercury, arsenic,

lead

• Calcium—contributes to “hard water”– Scaling in showers, on cookware– Bathtub ring– Binds to soap and renders in inactive

• Won’t lather• Flat hair

How do we “soften” hard water?

• Ion-exchange resin filter: sodium ions on resin beads are exchanged for calcium ions in hard water. Calcium “sticks” to bead.

• Electrons are shared between atoms

• Involves 2 or more nonmetals

• Want to fulfill the octet rule.

• Molecules are made.

• Properties are a result of bond type

• (Polar and nonpolar will be discussed later)

• When the repulsive forces are overcome, the atoms can bond.

• Energy is released when bonds are made, and when bonds are broken.

• This energy can be measured.– How? Heat released, light, etc.

• The same amount of energy that is required to make a bond is needed to break the bond.

–Energy in = Energy out

Covalent bonds:

• Can be measured: Bond length is the distance between 2 nuclei.

• Are flexible: vibrate like a slinky, so bond length is variable.

• The longer the bond length, the weaker the bond (lower energy).

• The shorter the bond, the stronger the bond (higher energy).

Properties of covalent molecules.

1. Can be a solid, liquid, or gas.– The intermolecular attraction between

molecules is much weaker than ionic.• Gases: very weak intermolecular forces, so

molecules are spread apart

2. When struck, pulverizes or turns into a powder. Can also be soft.

3. Low melting and boiling points– Due to weak intermolecular forces.– Liquids are already melted at room temp.– Gases are already boiled at room temp.

4. Solubility in water is variable.

5. Generally not conductive as a solid, and not conductive when dissolved.

– Distilled water is NOT conductive

• Involve only metal atoms

• “Sea of electrons” hold adjacent atoms together, and contribute to the properties of metals.

Properties of Metallic bonds

1. Shiny solids

2. Malleable (flattens out) and ductile (can be stretched into wires)

3. Pretty high melting points

4. Insoluble in water.

5. Highly conductive

Polar covalent bonds:

• Have an uneven distribution of electrons/shared unevenly.

• Polar ends have a partial “charge”. More like a pull.

• Structure is bent.

Nonpolar covalent bond:

• Even distribution of electrons between atoms.

• No partial charge.

• Structure is usually linear (straight).

Using electronegativity values to predict bond type:

• Electronegativity: a measure of an atom’s tendency to attract electrons.

• Higher value = stronger pull.

• We use the difference between the electronegativity values of the atoms in a molecule.

• If the electronegativity difference is:• less than 0.5, then the bond is nonpolar covalent.• between 0.5 and 1.6, the bond is considered polar

covalent• greater than 2.0, then the bond is ionic.• between 1.6 and 2.0:

– if a metal is involved, then the bond is ionic. – If only nonmetals are involved, the bond is polar covalent.

Examples:• Sodium chloride: NaCl

Na = .9 Fluorine = 3.1

• 3.1 - .9 = 2.2 ionic

• Nitrogen dioxide: NO2

N = 3.0 O = 3.4 difference = .2 nonpolar covalent

• Carbon dioxide: CO2

C = 2.5 O = 3.4 difference = .9

But….structure is overriding factor

So….

• The type of elements involved is the key determiner between ionic and covalent, and…

• The structure of the molecule is the overriding factor in determining polarity.

How do we know the shape?• Lewis dot structures

– Dots represent valence electrons.– Allow us to determine polarity.– Satisfies octet rule.

Do some practice:

• H2S

• CH3OH

Lewis dots with double bonds:

• Any unshared electrons will form a double or triple bond.

• Example: carbon dioxide

• Practice: C2H4

Naming Covalent Compounds:

• Must specify how many of each element there is (unlike ionic compounds)

• Ionic: MgCl2 is magnesium chloride

• Covalent: CCl2 is carbon dichloride.

• Why?

How do we know whether it is an ionic or covalent

compound?

• Ionic is a metal and nonmetal

• Covalent is only nonmetals

Naming covalent compounds:

• Say the first element, then the second, indicating how many of each element

• Second elements may have different wording

• i.e. sulfur is sulfide, oxygen is oxide

• Do not say “mono” for first element.

Examples:• SCl4 :

• Silicon tetrachloride

• CO :

• Carbon monoxide

• N2F5 :

• Dinitrogen pentafluoride

• CCl2F2

• Carbon dichloride difluoride