Chapter 4.1 Notes Ionic compounds consist of ions held together in lattice structures by ionic bonds

Important terms for this section Chemical bond Valence electrons Ions Cations Anions Ionization Ionic bond Ionic compound Lattice Formula unit Lattice energy Volatility Hydrated Solvated Binary Electronegativity values Bonding continuum Note: only outer shell (valence) electrons take part in bonding

Ionic bonding Atoms bonded chemically have different properties than the individual atoms Ionic bonds are chemical bonds between oppositely charged particles Ions are created by the gain or loss of valence electrons Cations: metals – loss of electrons so positively charged (groups 1, 2, 13) Anions: non-metals – gain of electrons so negatively charged (groups 15, 16, 17) Stable transition metal ions to know: Pb2+, Sn4+, Sn2+, Ag+, H+, H-, Cu+, Cu2+, Fe2+, Fe3+ Polyatomic ions to know: see page 142

Ionic compounds Ionic bonding: transfer of electrons from one atom to another Ionic formula writing and naming: Mg2+ + F-  NH4+ + PO4 3-  Al3+ + S2-  Ionic formula writing: Sodium sulfate Magnesium cyanide Mercury hydrogencarbonate

Ionic compound structure Ionic lattice is formed: Electrostatic attraction Coordination number: Number of ions around the given ion in lattice Lattice can grow indef. Formula unit: Ratio of ions present Lattice energy: Strength of attraction between ions in same lattice (smaller = ↑ attraction)

Physical props. reflect lattice structure Melting and boiling points are usually very high (strong electrostatic attraction means stronger bonds) Industry tends to not extract metals from molten ionics due to the very high temps needed Low tendency to vaporize = low volatility Therefore, low odor (think of table salt)

Physical props. reflect lattice structure Solubility: how easy will it dissolve into a solvent (become dispersed through a liquid) to form a solution Like dissolves like! Salt is ionic, so it will likely dissolve into….? Polar solvent Non-polar solvent Ionic solvent Ions dissociate in water and are surrounded by water molecules = hydrated Ions dissociate in other liquid = solvated

Physical props. reflect lattice structure Electrical conductivity: Ionic compounds that are molten have moveable charges Ionic compounds dissolved in solution have moveable charges Brittleness: Ionic compounds tend to shatter on impact (not malleable or ductile) This occurs because the ions moving from impact will put like charges next to each other, so they repulse and split

Ionic character (how we know it’s ionic) How do you know if it’s ionic or covalent? Sometimes the metalloids (for example) show some grey areas… Si can have 4+ or 4- and H can have 1- or 1+ Remember:

Ionic character (how we know it’s ionic) Determining if a compound is ionic or covalent depends on the difference of electronegativity (section 8 of data booklet)

Ionic character The bonding continuum is the range of ionic to covalent Based on difference of electronegativity values Which will likely form ionic bonds? Covalent? Polar covalent? Be and F Si and O N and Cl K and S

Covalent compounds form by the sharing of electrons Chapter 4.2 Notes Covalent compounds form by the sharing of electrons

Important terms in this section Molecule Diatomic Triatomic Octet rule Non-bonding pairs Double bond Triple bond Bond length Bond strength Bond enthalpy Polar Directionality Dipole Partial charge Pure covalent

Covalent bonding Atoms in covalent molecules share electrons Covalent bond is an electrostatic attraction between a pair of e- and the positive nuclei Atoms with covalent bonds are called molecules (as opposed to the ionic called formula units) Hₓ •H  H – H Two hydrogen atoms bond covalently to form a hydrogen molecule (these are diatomic molecules) ex. of triatomic? Covalent bonding is favorable (stabilizes the atoms) so forming the bonds releases energy

Bonding releases energy

Covalent bonding Octet rule: tendency for atoms to form stable arrangements of 8 e- in outer shell (like Ar) This is not really a “rule” but a common occurrence as this “rule” has exceptions Non-bonding pairs or lone pairs: pairs of electrons on an atom that do not participate in the covalent bond

Multiple bonds & bond length/strength Double and triple bonds can be formed when there are not enough electrons to give all atoms a desired octet Double bond: sharing of 4 e- Triple bond: sharing of 6 e- Bonds: use sections 10 & 11 in data booklet Bond length: measure of distance of bonded nuclei Bond strength: measure of E needed to break the bond (bond enthalpy)

Bond length and strength Increase of number of bonds increase strength and decreases length

Polar covalent bonds These bonds do not share electrons evenly There is DIRECTIONALITY to the bond and/or the molecule Dipole: having one end being slightly more positive and one end being slightly more negative

Polar covalent bonds Again, the difference in electronegativity will tell you the extent of how polar the bond

Pure covalent bond When the electronegativity difference is zero H – H F – F are examples C – H is typically referred to as non-polar, but there is a slight polarity as C is more electronegative (0.4) Also referred to as non-polar covalent

The continuum from ionic to pure covalent demonstrates why it is better to use terms such as “predominately” covalent or ionic or “strongly polar” to describe bonds or substances.