Ch7 Bonding and structure Flashcards
Ionic bond basic definition
Electrostatic attraction between oppositely charged ions(positive and negative)
Ionic bond description
metal lose electron(s) from valence shell and become positive anion
electron is transferred to non-metal which becomes negative cation
Oppositely charged ions are attracted to one another and form bond
4 Properties of giant ionic lattice
- High mpt+bpt- ionic bonds are strong and occur throughout, requiring large amount of energy to overcome
- poor conductor of electricity when solid- ions are in fixed positions and cannot move so no mobile carriers
- good conductors of electricity when molten- ions are mobile and free to move
- soluble in water and polar solvents
Hydration of ionic solids
some ionic solid are soluble in water
the atoms have charges which are attracted to polar water molecules and bind to them
energy released as bonds are formed is enough to compensate energy needed to overcome ionic bonds
Insoluble ionic solids
Energy needed to be released when ions are hydrated is not sufficient to balance energy required to separate ions
Covalent bond basic definition
electrostatic attraction between shared pair of electrons and nuclei of two bonded atoms
Covalent bond description
Adjacent non-metal atoms share electrons in valence shells to achieve full outer shells
Dative covalent bond
two atoms share two electrons but one atom provides both electrons from a lone pair
e.g. NH3 + H+ -> NH4+
Lone pair of electrons
- can form variations on covalent bond- dative covalent bond
- affect shape of molecules
- can behave as nucleophiles in organic reactions
3 Properties of covalent compounds
- low mpt + bpt - attraction between discrete, separate neighbouring covalent molecules are weak so small amount of energy required to overcome
- poor conductor electricity - no mobile charge carries, ions or electrons
- soluble in non-polar solvents
Electron pair repulsion theory
Electron pairs repel each other as far apart in space as possible around central atom
shape of molecule depends upon number and type of electron pairs around the central atom
Lone pairs repel more than bonded pairs(subtract 2.5° in bond angle)
Shape of molecules with 2 bonded pairs and no lone pairs
Linear
Bond angle 180°
Shape of molecules with 3 bonded pairs and no lone pairs
Trigonal planar
Bond angle 120°
Shape of molecules with 4 bonded pairs and no lone pairs
tetrahedral
bond angle 109.5°
Shape of molecules with 5 bonded pairs and no lone pairs
Trigonal bipyramidal
Bond angle 120° and 90°
Shape of molecules with 6 bonded pairs and no lone pairs
Octahedral
Bond angle 90°
Shape of molecules with 3 bonded pairs and 1 lone pair
Pyramidal
Bond angle 107°
Shape of molecule with 2 bonded pairs and 2 lone pairs
Angular(bent)
Bond angle 104.5°
Metallic bond basic definition
Electrostatic attraction between delocalised electrons and positive ions held within the lattice
Delocalised electrons
Bonding electrons that are not fixed between two atoms in a bond
Mobile and are shared by several atoms
3 properties of metallic structures
- high mpt + bpt- large amount of energy needed to overcome many strong metallic bonds
- good conductor of electricity- contains mobile charge carriers(sea of delocalised electrons) that are free to move
- malleable and ductile- nondirectional, independent of shape and exists in all directions
electronegativity
the tendency of an atom to attract the shared pair of electrons in a covalent bond
Bond polarity
when pair of electrons making up asymmetrical covalent bond aren’t shared evenly between the atoms
ionic character when there is a difference in electronegativity in two atoms forming covalent bond
creates a dipole bond
Two trends in electronegativity
- Increases across a period
- Decreases down a group
symmetrical shapes
linear, trigonal planar, tetrahedral and octahedral
asymmetrical shapes
pyramidal and angular(bent)
Intermolecular forces
forces of attraction that occur between molecules
Permanent dipole-dipole interactions
weak electrostatic forces between dipoles of neighbouring molecules
Hydrogen bonds
relatively strong electrostatic attractions between polar molecules (that contain hydrogen and has lone pair) covalently bonded to elements with higher electronegativity e.g. F, O, N
Examples of hydrogen bonds
water H2O
ammonia NH3
hydrogen fluoride HF
compounds with -OH group
compounds with -NH2 group
Water propertied due to hydrogen bonds
- solid ice is less dense than liquid water - h20 molecules hydrogen bond to form interlocking hexagonal rings, so are held further apart creating greater volume and lower density
- water has high mpt + bpt - additional energy required to overcome large amount of hydrogen bonds
Induced dipole-dipole interactions
Temporary movement of electrons generates an instantaneous dipole
This induces other dipoles in neighbouring molecules
Two dipoles generate a weak temporary force of attraction between molecules
- weakest intermolecular force
- larger amount of electrons, greater number of forces
