Chemistry Topic 4

Define electronegativity

Define electronegativity

The ability for an atom to attract electrons in a covalent bond

Define permanent dipole

The permanent seperation of partial charges between two different atoms

Define polar molecule

A molecule with a net diople due to bonds of different polarity or bonds that aren’t symmetrically arranged

Define covalent bond

A bond formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.

Define metallic bond

A bond formed by the electrostatic attraction between a metal cation lattice and a sea of delocalized electrons.

Define ionic bond

A bond formed by the electrostatic attraction between two oppositely charged ions.

Define non-polar molecules

a molexule with no net dipole because the bonds are of equal polarity (involve the same elements) and are arranged symetrically with respect to eachother, cancelligng each other out.

What is the typical state and structure of ionic compounds under normal conditions?

Solid with lattice structures

What happens to bond length and bond strength when the number of shared electrons increases in a covalent bond?

Bond length decreases, bond strength increases.

What electronegativity difference is needed for a covalent bond to be considered non-polar?

<0.4

What is the octet rule?

The tendency of atoms to gain a valence shell with a total of 8 electrons.

When do resonance structures occur?

When there is more than one possible position for a double bond in a molecule

why do resonance hyrbids occur?

delocalized electrons from Pi bonds that spread themselves out between bonds

bond angle for tetrahedrals

109.5

bond angle for trigonal planars

120

bond angle for linear shapes

180

bond angle for trigonal pyramidal

107

bond angle for bent or V-shaped

104.5

Why do lone pairs of electrons in the valence shell around the central atom decrease the bond angle (VSEPR theory)

electron pairs in the same valence shell carry the same charge, they repel each other and so spread themselves out as far as possible (reducing the bond angle by 2.5 degrees per lone pair in a tetrahedral)

Lone-pairs have a higher concentration of charge than a bonding pair because they are not shared between two atoms, and thus have slightly more repulsion than bonding pairs.

Carbon and silicon form…

giant covalent network/layer lattice structures

When is something conductive

When there are charged particles that can move freely

Define london dispersion force

A weak force that occurs between opposite partially charged ends of two temporary/instantaneous dipoles in non-polar molecules

Define dipole-dipole forces

The attractive forces between the positive end of one polar molecule and the negative end of another polar molecule.

Define hydrogen bonding and explain the criteria for it

The strongest type of intermolecular force that only occurs between molecules in which:

• a hydrogen atom is covalently bonded to a highly electronegative oxygen, nitrogen or fluorine atom which creates a significant partially positive charge on the hydrogen atom.

• The significant partial positively charged hydrogen atom is then attracted to the lone elctron pair on a oxygen, nitrogen, or fluorine atom of a neighbouring molecule.

Intermolecular forces in descending order of strength

1. Hydrogen bonds

2. Dipole-dipole forces

3. London (dispersion) forces

How to determine strength of metallic bond

Cation radius and charge strength

greater positive charge and smaller cation radius = stronger metallic bond = higher melting point

What are the properties of alloys

• usually contains more than one metal held together by metallic bonding

• Harder and stronger (less malleable)

• melting point is a mix of the metals melting points

Why are alloys harder and stronger (less malleable)?

Different sized cations or atoms are introduced into the lattice, disrupting the regular lattice layers so they cannot slide over each other easily.

Outline why metals are conductive

they have a sea of delocalized electrons that are highly mobile

Outline why metals have good thermal conductivity

delocalized electrons and close packed ions enable efficient transfer of heat energy

Explain why metals are malleable (can be shaped under pressure) and ductile (can be drawn into threads)

movement of delocalized electrons is non-directional and random through the cation lattice, so metallic bond remains intact while the conformation changes under pressure

Explain why metals have high melting points

metallic bonds are strong and a lot of energy is required to seperate the atoms

Outline why metals are shiny/lustrous

delocalized electrons in metal crystal structure refelct light

NH4^+

Ammonium ion

NH3

Ammonia

nitrate ion

NO3^-

sulfate ion

SO4^2-

carbonate ion

CO3^2-

phosphate ion

PO4^3-

SO3^2-

sulfite ion

Nitrite ion

NO2^-

ClO^-

Hypochlorite ion

ClO2^-

Chlorite ion

Chlorate ion

ClO3^-

ClO4^-

Perchlorate ion

Permanganate ion

MnO4^-

ethanoate ion

CH3COO^-

why are ionic compounds soluble in water

ions attracted to polar water molecules

polar solutes generally dissolve in polar solvents like water

explain why ionic compounds such as NaCl are bad conducters when solid, but are excellent conducters when dissolved or in their molten state.

When ionic compounds are in their solid state they form a lattice like structure, and therefore the ions are not free to move and can’t conduct electricity. Where as in aqueous solution or when molten the ions can move freely and conduct electricity.

Why are ionic compounds relatively non-volatile (how readily a substance vaporizes)

ionic bonds are strong and hard to break so ionic compounds have a high melting and boiling point.

When is a covalent bond polar

electronegativity difference > 0.4

When comparing boiling points…

identify and compare the strengths of intermolecular forces

Explain why the strength of london (dispersion) forces increase with molecular size (Mr) and thus increase the boiling point of the substance

There is a greater number of electrons within the molecule, increasing the probability for a temporary/instantaneous dipole to form.

Note: individual strength of the disperion force stays the same but overall dispersion force strength increases as there is more.

HCO3

bicarbonate

Describe the structure of graphite and explain why it is a good electrical conductor

Each C atom is bonded to 3 others, forming hexagons in parrallel layers with bond angles of 120 degrees. The layers are held together by weak London dispersion forces.

It is a good electrical conductor because it contains one delocalized electron per carbon atom that gives the electron mobility.

Describe the structure of diamond and explain its properties

each C atom is covalently bonded to 4 others tetrahedrally arranged in a repetitive pattern with bond angles of 109.5 degrees.

Properties:

• Tetrahedrally arranged lattice structure and smaller molecule size meaning lots of shorter and stronger covalent bonds = hardedst known substance +very high melting point + extremely good thermal conductor

• No delocalized electrons in solid state = non-conductor of electricity

Buckminsterfullerene C60 structure a summary of properties

a spherical cage in which each C atom is bonded to three others, forming hexagons & pentagons.

• semiconductor at normal temp/pressure due to some electron mobility

• very bad thermal conductor

• very light and strong

• soluble in benzene

Describe the structure of graphene and summarize its properties

A 2D single layer material where each C atom is covalently bonded to 3 others forming hexagons with bond angles of 120 degrees (like graphite but only one layer)

• One delocalized electron per atom = very good electrical conductor

• Best thermal conductor ever known

• almost completely transparent

• very flexible

• very high melting point

What is a coordinate covalent bond?

A covalent bond where one atom donates a pair ofelectrons to make that bond (lewis base to lewis acid or metal ligand)

Van der Waals forces

An inclusive term, which includes dipole–dipole, dipole-induced dipole and London \n (dispersion) forces.

Example of an alloy

Bronze - mixture of mostly copper and a smaller amount of tin

• Melting point inbetween copper and tin but closer to copper

• hard (not malleable)