IB Chem Structure Mix

The octet rule

Atoms bond together in order to achieve a full valence shell containing 8 electrons

Covalent bonding

Sharing of electrons

Bonding between non-metal elements

Ionic bonding

Transfer of electrons

Lewis structures

Show the bonding and the non-bonding electrons

Coordinate covalent bonds

One atom contributes both the bonding electrons to the bond

Intermolecular forces

Forces between molecules

Determines the physical properties of a substance

Intramolecular forces

Forces within molecules

London dispersion forces

Caused by the movement of electrons within an atom or molecule.

All atoms/molecules have London dispersion forces.

Can cause an instantaneous dipole. If a molecule with an instantaneous dipole gets close to another molecule it can induce a dipole.

Molar mass increases → London dispersion forces increase, leading to an increased BP.

Dipole - dipole forces

Between polar molecules.

The electrostatic attraction between the partial positive charge on one molecule and the partial negative charge on another.

Hydrogen bonding

When a hydrogen atom is bonded to a nitrogen, oxygen of fluorine atom.

A stronger dipole-dipole force

Polar solvent

A liquid composed of polar molecules

I.e. Water, methanol, ethanol, propane, ethanoic acid

Polar substances are soluble in polar solvents

Non-polar solvents

A liquid composed of non-polar molecules

I.e. hexane, octane, benzene, methylbenzene, carbon tetrachloride

Non-polar substances are soluble in non-polar solvents

Ion-dipole forces

Occur between ions and oppositely charged dipoles of water molecules in aqueous solutions.

H₂O surrounds the ion forming a hydration shell.

Retention factor (R_f) value

R_f=distance travelled by solute / distance travelled by solvent front

Properties of metals

Good conductors of heat and electricity

Malleable (can be bent into shape)

Ductile (can be drawn into wires)

Metallic bonding

The electrostatic attraction between the lattice of cations and the sea of delocalised electrons.

The strength of the bond is determined by:

• Charge on the ion

• Ionic radius of the ion

• Number of delocalised electrons

Alloys

Materials that are a mixture of a metal and other metals or non-metals.

Different properties; usually harder (less malleable) and have greater tensile strength (stronger)

Can distort the lattice struncture, making it more difficult for the layers to slide over eachother.

Graphite

Layered structure held together by weak intermolecular forces.

The layers can slide over one another.

Each C atom is bonded to 3 other C atoms.

Bond angle is 120^o

Shape: Trigonal planar

Good conductor of electricity (delocalised electrons).

Diamond

Giant covalent structure

High MP and BP, very hard.

Each C atom is bonded to 4 other C atoms

Bond angle is 109.5^o

Shape: Tetrahedral

Doesn’t conduct electricity.

Fullerene C₆₀

Structure consist of 12 pentagons and 20 hexagons.

Each C atom is bonded to 3 other C atoms.

Bond angle is 120^o

Shape: Trigonal planar

Shows some electrical conductivity

Graphene

One layer thick, very strong

Each C atom is bonded to 3 other C atoms.

Bond angle is 120^o

Shape: Trigonal planar

Very good electrical and thermal conductivity.

The periodic table: Group

Vertical column

Numbered 1-18

The number of valence electrons

The periodic table: Period

Horizontal row

Number = The outer energy level that is occupied by electrons

Metalloids

Have properties intermediate between those of metals and non-metals

B, Si, Ge, As, Sb, Te

Po and At are sometimes recognised as metalloids

Electron shielding

The inner electrons shield that outer valence electrons from the full attraction of the nucleus.

The valence electron(s) within an atom require less energy to remove than the inner electrons.

Increases down a group. Remains constant across a period.

Effective nuclear charge (Z_eff)

The net positive charge experienced by valence electrons.

Z_eff=Z-S (approximate value)

Z is the atomic nr

S is the nr of shielding electrons

Trends in atomic radii

Atomic radius decreases across a period

Reasons:

• Nuclear charge increases

• Electron shielding remains constant

Atomic radius increases down a group

Reasons:

• The number of occupied energy levels increases and the outer electrons are further from the attraction of the nucleus.

Trends in ionic radii

Positive ions lose electrons to obtain a full outer shell.

Positive ions are smaller than their parent atoms.

The ions has more protons than electrons → stronger attraction between nucleus and electrons.

Negative ions gain electrons to obtain a full outer shell.

Negative ions are bigger than their parent atoms.

The ions has more electrons than protons → weaker attraction between nucleus and electrons.

The first ionisation energy

The energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous 1⁺ ions.

Electron affinity

The energy released when one mole of electrons is added to one mole of gaseous atoms/ions to form one mole of ions.

The first electron affinity is exothermic (- value)

• In general, the greater the atomic radius and the greater the electron shielding, the less energy is released when an electron is added.

The second electron affinity is endothermic (+ value)

• Positive due to the extra repulsion when adding negative electrons to an already negative ion.

Electronegativity

A measure of the attraction of an atom for a bonding pair of electrons.

Increases across a period because of increasing nuclear charge.

Decreases down a group because of increasing atomic radius.

Metallic character

How easily an atom can lose electrons.

Increasing down a group. Increasing atomic radius results in a weaker attraction between nucleus and valence electrons.

Decreasing across a period. Increasing nuclear charge and decreasing atomic radius results in a stronger attraction between the nucleus and valence electrons.

Group 1 metals (alkali metals)

Li, Na, K, Rb, Cs, Fr

Li, Na, K float on water due to their low densities.

Stored in oil to prevent the reaction with oxygen in the air.

React vigorously with water to produce an alkaline solution and hydrogen gas.

React with group 17 elements to produce salts.

Group 17 elements (the halogens)

F, Cl, Br, I, At

Coloured: F is a pale yellow gas, Cl is a greenish-yellow gas, Br is a reddish-brown liquid and iodine a purple solid.

Reactivity decreases down the group.

When reacted, the more reactive halogen displaces the ions of the less reactive halogen from the solution.

Homologous series

A series of organic compounds of the same family which differ by a common structural unit.

E.g. CH3OH, CH3CH2OH, CH3CH2CH2OH

Show a gradation in physical properties, e.g. increasing BP

Factors which affect the BP of organic compounds

• Molar mass

• Structure (straight-chain vs branched-chain isomers)

• Type of functional group (hydrogen bonding, dipole-dipole, London dispersion)

Structural isomers

Compounds with the same molecular formula but different structural formulas