Bonds: forces that hold compounds together
- Strong: chemical bonds → ionic, covalent, metallic
- Weak: intermoleculare forces (only apply to covalent molecules)
Ionic bond: when electrons are lost and gained between two elements with very different
electronegativities (>1.8 electronegativity difference); metal + non metal
Covalent bond: when electrons are shared between the two elements (<1.8 electronegativity
difference - both gain); two non metals
Metallic: electron cloud created around the cations (<1.8 electronegativity difference - both
lose); two metals
Ionic bonds
- It occurs between metals and non metals
- Metals lose electrons (+1) and nonmetals gain (+1)
- Anions and cations attract due to electrostatic forces
- Ionic compounds DO NOT form molecules; they DO form giant structures called lattices
- Lattices: 3D nets of anions and cations alternated
- They are solid at room temperature; they have high melting and boiling points
**Difference between ions in gas and solid state energy = lattice energy
- They are brittle (break easily); this is due to a break in alternation (change in relative
position) which causes the anions to repel
- Lattice energy (U) is directly proportional to the product in ion charges |q1q2| and
inversely proportional to the distance between nuclei
- High U means high mp/bp; Low U means low mp/bp
- Soluble in water
- They DON’T conduct electricity in solid state but they DO conduct in liquid state
Formulation of ionic compounds
- Neutral: as many negative as positive charges
- Cations
- 1 possible charge: name of the element + ion (Na+ = sodium ion)
- 2 or more possible charges: state the charge in roman numerals (Cu2+ =
copper II ion)
, - Polyatomic cation: H3O+ = hydronium; NH4+ = ammonium
- Anions
- 1 possible charge: suffix -ide (Cl- = chloride ion)
- Polyatomic:
- O22- = peroxide
- OH- = hydroxide
- NO3- = nitrate (V) ion
- PO43- = phosphate (V) ion
- CO32- = carbonate (IV) ion
- HCO3- = hydrogen carbonate (IV)
** suffix -ate means there is oxygen in the compound
- Compounds
- Name the cation first and then the anion
- Deduce charge of the cation from the charge of the anion
Covalent bonding
- Non metals bond without losing electrons as they both have a high electronegativity
- They share electrons although not equally (no pure covalent compound)
- They are held together by forces between nuclei and electrons
- Lewis diagrams show how valence electrons are arranged in a covalent bond
Octate rule: atoms are more stable with 8 valence electrons
- Exception to the octet rule:
- Boron (B) is stable with 6 electrons
- Nonmetals in period 3 and below can expand their octates
Dative covalent bonds: it’s a covalent bond where both electrons are given by the same atom ;
both atoms come from the same element
- CO, OH3, NH4+, SO2
Resonance hybrids: it means a molecule has alternative arrangements of the electrons in the
molecule; this is due to the presence of delocalized electrons; both atoms come from the
- When a central atom is joined to 2 or more external atoms (same element) and some
bonds are double and some are single in reality the bond is neither.
- Strong: chemical bonds → ionic, covalent, metallic
- Weak: intermoleculare forces (only apply to covalent molecules)
Ionic bond: when electrons are lost and gained between two elements with very different
electronegativities (>1.8 electronegativity difference); metal + non metal
Covalent bond: when electrons are shared between the two elements (<1.8 electronegativity
difference - both gain); two non metals
Metallic: electron cloud created around the cations (<1.8 electronegativity difference - both
lose); two metals
Ionic bonds
- It occurs between metals and non metals
- Metals lose electrons (+1) and nonmetals gain (+1)
- Anions and cations attract due to electrostatic forces
- Ionic compounds DO NOT form molecules; they DO form giant structures called lattices
- Lattices: 3D nets of anions and cations alternated
- They are solid at room temperature; they have high melting and boiling points
**Difference between ions in gas and solid state energy = lattice energy
- They are brittle (break easily); this is due to a break in alternation (change in relative
position) which causes the anions to repel
- Lattice energy (U) is directly proportional to the product in ion charges |q1q2| and
inversely proportional to the distance between nuclei
- High U means high mp/bp; Low U means low mp/bp
- Soluble in water
- They DON’T conduct electricity in solid state but they DO conduct in liquid state
Formulation of ionic compounds
- Neutral: as many negative as positive charges
- Cations
- 1 possible charge: name of the element + ion (Na+ = sodium ion)
- 2 or more possible charges: state the charge in roman numerals (Cu2+ =
copper II ion)
, - Polyatomic cation: H3O+ = hydronium; NH4+ = ammonium
- Anions
- 1 possible charge: suffix -ide (Cl- = chloride ion)
- Polyatomic:
- O22- = peroxide
- OH- = hydroxide
- NO3- = nitrate (V) ion
- PO43- = phosphate (V) ion
- CO32- = carbonate (IV) ion
- HCO3- = hydrogen carbonate (IV)
** suffix -ate means there is oxygen in the compound
- Compounds
- Name the cation first and then the anion
- Deduce charge of the cation from the charge of the anion
Covalent bonding
- Non metals bond without losing electrons as they both have a high electronegativity
- They share electrons although not equally (no pure covalent compound)
- They are held together by forces between nuclei and electrons
- Lewis diagrams show how valence electrons are arranged in a covalent bond
Octate rule: atoms are more stable with 8 valence electrons
- Exception to the octet rule:
- Boron (B) is stable with 6 electrons
- Nonmetals in period 3 and below can expand their octates
Dative covalent bonds: it’s a covalent bond where both electrons are given by the same atom ;
both atoms come from the same element
- CO, OH3, NH4+, SO2
Resonance hybrids: it means a molecule has alternative arrangements of the electrons in the
molecule; this is due to the presence of delocalized electrons; both atoms come from the
- When a central atom is joined to 2 or more external atoms (same element) and some
bonds are double and some are single in reality the bond is neither.
