Chemical structure and bonding Flashcards
Properties of metal
- high melting point
- good conductors of electricity and heat
- malleable and ductile
Metallic bonding
and what increases its strength
Strong electrostatic attraction between metal cations and delocalised electrons
Strength increased by:
- more delocalised electrons
- smaller ion size
Metallic structure
Giant lattice of metal cations with delocalised electrons
Why metal conducts electricity
and what makes it better at conducting
Delocalised electrons can carry negative charge through the metal
Conductivity increased by:
- more delocalised electrons
- smaller ion size (for higher electron density)
Why metal is a good thermal conductor
Delocalised electrons quickly transfer kinetic energy around the metal
Why metal is malleable and ductile
- metallic bonding relies on delocalised electrons so it is non-directional
- if layers of the lattice slide over each other, bonding stays intact
Ionic bonding
and what increases its strength
Strong electrostatic attraction between oppositely charged ions
Strength increased by:
- larger ion charges
- smaller ion size
Properties of ionic compounds
- high melting point
- conducts electricity when molten or dissolved
- brittle
- soluble in water
Why ionic compounds conduct electricity only when molten or dissolved
- solid ionic compounds have charged particles, but they are not free to move
- as part of a liquid, the ions can move freely and carry charge
Why ionic compounds are brittle
- ionic bonding is between oppositely charged ions
- if layers of the lattice slide over each other, ions with the same charge repel, breaking the bonding
Why ionic compounds are soluble in water
- water is a polar molecule
- poles of water molecules surround oppositely charged ions, pulling them out of the lattice
Demonstration of the existence of ions
Electrolysis of green copper chromate solution separates blue copper ions from yellow chromate ions
Covalent bonding
and what increases its strength
Strong electrostatic attraction between two nuclei and a bonding pair of electrons
Strength increased by:
- shorter bond length (smaller atomic radii)
Sigma bond
- covalent bond formed by the overlap of two orbitals directly between the nuclei
- the first type of covalent bond to form
- stronger than a pi bond
Pi bond
- covalent bond formed by a sideways overlap of two p orbitals
- only forms after a sigma bond is present
- present in double or triple covalent bonds
- weaker than a sigma bond
Electronegativity
and what increases it
An element’s ability to attract a bonding pair of electrons
Increased by: - higher nuclear charge - smaller atomic radius - less electron shielding (highest in top right of table)
Polar covalent bond
- covalent bond between atoms of different electronegativity (different elements)
- EN diff causes a dipole to form
Dative covalent bond
- covalent bond where both bonding electrons are from the same atom
- functionally the same as any other covalent bond
Octet rule
- atoms usually want 8 electrons in their outer shell
- C, N, O, and F always follow this
- H, Be, B, and elements past period 3 do not follow
Polarising power of a cation
Ability of a cation to distort an anion’s electrons (increases as EN increases)
Polarisability of an anion
Tendency of an anion to be distorted by a cation (increases as EN decreases)
London forces
- occurs in all molecules
- electron density constantly moves around in molecules
- when it is higher on one side, a dipole is formed
- this induces more dipoles on neighbouring molecules, which are then attracted
What makes london forces stronger
- more electrons (larger molecule)
- more points of contact (flatter molecule / less branches)
Permanent dipole forces
- occurs only in polar molecules
- attraction between the dipoles of polar molecules
Hydrogen bonding
- H is bonded to something very electronegative (N, O, or F)
- a lone pair of electrons is present
- strong attraction between the very positive H pole and the negative electrons
- the strongest intermolecular force
Hydrogen bonding in water
- HF has 3 lone pairs per molecule but only 1 H
- ammonia has 3 Hs but only 1 lone pair
- water has 2 lone pairs and 2 Hs so can form 2 hydrogen bonds per molecule
What is required for two liquids to mix?
They must have the same intermolecular forces
Intermolecular forces in alcohols
- all alcohols can hydrogen bond due to the OH group
- as the carbon chain gets longer, london forces increase and hydrogen bonding decreases
- larger alcohols are less soluble in water
Why is ice less dense than water?
Hydrogen bonds cause the solid molecules to form a crystal structure that is less dense than water
Structure of diamond
- giant covalent lattice
- 4 sigma covalent bonds per carbon
- very hard with high melting point
Structure of graphite
- giant covalent lattice
- 3 sigma covalent bonds per carbon with delocalised 4th electrons
- conducts electricity
- layers easily slide over each other
- high melting point
Structure of fullerene
- large ball shaped molecule
- 3 bonds per carbon with delocalised 4th electrons
- cannot conduct electricity because molecules
- low melting point
Repulsion between electron pairs
- bonding regions cause less repulsion
- lone pairs cause more repulsion
Shape of a carbon dioxide molecule
- linear
- 2 bond regions, 0 lone pairs
- bond angle of 180
Shape of a boron trifluoride molecule
- trigonal planar
- 3 bond regions, 0 lone pairs
- bond angles of 120
Shape of a sulfur dioxide molecule
- v shaped
- 2 bonding regions, 1 lone pair
- bond angle of 119
Shape of a water molecule
- v shaped
- 2 bonding regions, 2 lone pairs
- bond angle of 104.5
Shape of a methane molecule
- tetrahedral
- 4 bonding regions, 0 lone pairs
- bond angles of 109.5
Shape of an ammonia molecule
- trigonal pyramidal
- 3 bonding regions, 1 lone pair
- bond angles of 107.5
Shape of a phosphorus pentachloride molecule
- trigonal bipyramidal
- 5 bonding regions, 0 lone pairs
- bond angles of 120, 180, 90
Shape of a sulfur hexafluoride molecule
- octahedral
- 6 bonding regions, 0 lone pairs
- bond angles of 90, 180
What makes a molecule polar?
- polar covalent bonds
- shape of the molecule (so that dipoles don’t cancel each other out)
Example of a non-polar molecule with polar bonds
Methane, tetrachloromethane, carbon dioxide
Example of a polar molecule
Water, chloromethane, ammonia
Types of bonds present in single, double, and triple covalent bonding
Single - 1 sigma
Double - 1 sigma, 1 pi
Triple - 1 sigma, 2 pi
Formation of ions
Electrons lost»_space; cation formed
Electrons gained»_space; anion formed
Why ionic radius decreases across a period (for isoelectronic ions)
- shielding stays the same
- nuclear charge increases
- attraction between nucleus and outer shell increases
