Topic 2 Chemical bonding and structure Flashcards
State the physical properties of metals
- high melting temperatures
- good electrical conductivity
- good thermal conductivity
- malleability
- ductility
What is mainly responsibly for the characteristic properties of metals?
delocalised electrons
Define delocalised electrons
electrons that aren’t associated with any single atom or any single covalent bond
What is metallic bonding?
the electrostatic force of attraction between the nuclei of metal cations and delocalised electrons
Why do metals have high melting temperatures?
metals have a giant lattice structure so lots of electrostatic attractions need to be overcome
What determines the melting temperatures of a metal?
- the number of delocalised electrons per cation e.g. Group 1 metals have a lower melting temp than Group 2
- size of the cation e.g. smaller cations are closer to the delocalised electrons and greater force of attraction
Why can metals conduct electricity?
when a potential difference is applied across the ends of the metal, the delocalised electrons are attracted to and move towards the positive terminal
Why are metals good thermal conductors?
- Free-moving delocalised electrons pass kinetic energy along the metal
- cations are closely packed together and pass kinetic energy from one cation to the other
Why are metals ductile/malleable?
delocalised electrons can move throughout the metal structure
Define ionic bonding
the strong electrostatic attraction between oppositely charge ions
How do you determine the strength of ionic bonding?
by calculating the amount of energy required in one mole of solid to separate the ions to infinity so they can no longer interact
What are the factors that determine the strength of ionic bonding
- the smaller the ions, the more energy required (for ions of the same charge)
- higher ionic charge, increased strength
Name the physical properties of ionic compounds
- high melting temperatures
- brittleness
- poor electrical conductivity when solid but good when molten
- often soluble in water
Why do ionic compounds have high melting temperatures?
-consist of a giant lattice with lots of oppositely charged ions so combined electrostatic forces is large and requires large amounts of energy
Why do ionic compounds have poor electrical conductivity (when solid)?
- no delocalised electrons
- ions aren’t free to move when potential difference is applied
- molten will conduct as ions are mobile
- solid lithium nitride will conduct electricity
- aqueous ionic solutions conduct electricity
What is the most convincing evidence for the existence of ions?
the ability of an ionic compund to conduct electricity and undergo electricity when moleten or in aqueous solution
Define a covalent bond
a covalent bond is formed when 2 atoms share 1 or more pairs of electrons
What are the 2 types of overlap in a covalent bond?
sigma and pi bonds
end on end overlap and sideways overlap
Define bond length
the distance between nuclei of the 2 atoms that are covalently bonded together
What is electronegativity?
the ability of an atom to attract a bonding pair of electrons
Describe the trends of electronegatvity
- decreases down a group
- increases from left to right across a period
What is a polar covalent bond?
a type of covalent bond between 2 atoms where the bonding electrons are unequally distributed, because of this, one atom carries a slightly positive charge and the other a slightly negative charge
What is a discrete simple molecule?
an electrically neutral group of 2 or more atoms held togther by chemical bonds
What is a displayed formula?
shows each bonding pair as a line drawn between the two atoms involved
Whta is a dative covalent bond?
formed when an empty orbital of one atoms overlaps with an orbital containing a non-bonding/lone pair of electrons of another atom
Give examples where dative covalent bonds occur
- hydroxonium ion, H3O+
- aluminium chloride, Al2Cl6
What is the electron pair repulsion theory?
- the shape of a molecule or ion is caused by repulsion between the pairs of electrons surrounding the central atom (both bond pairs and lone pairs)
- the elctron pairs arrange themselves around the central atom so that repulsion between them is minimal
- lone pair-lone pair repulsion > lone pair-bond pair repulsion > bond pair-bond pair repulsion
What shape consists of 2 bond pairs?
linear e.g. BeCl2, CO2(each double bond treated as an electron pair)
What shape consists of 3 bond pairs?
trigonal planar e.g. BCl3
What shape consists of 4 bond pairs?
tetrahedral e.g. CH4
What shape consists of 5 bond pairs?
trigonal bipyramidal e.g. PCl5
What shape consists of 6 bond pairs?
octahedral e.g. SF6
What shape consists of 3 bond pairs and 1 lone pair?
trigonal pyramidal e.g. NH3
What shapes consists of 2 bond pairs and 2 lone pairs?
v-shaped e.g. H2O
Bond angle for linear shape e.g. BeCl2
180 degrees
Bond angle for trigonal planar shape e.g. BCl3
120 degrees
Bond angle for tetrahedral shape e.g. CH4
109.5 degrees
Bond angle for trigonal pyramidal NH3
107 degrees
Bond angle for v-shaped e.g. H2O
104.5 degrees
Bond angle for octahedral e.g. SF6
90 and 180 degrees
Why are ionic compounds brittle?
- layers of ions slide over one another when stress is applied
- ions of same charge are next to each other and repel one another which breaks crystal apart
Why are ionic compounds soluble in water?
-positive and negative ions both attracted to water molecules due to polarity that water possesses
What’s a dipole?
exists when 2 charges of equal magnitude but opposite signs are separated by a small distance
Define a hydrogen bond
an intermolecular interaction between a hydrogen atom of a molecule bonded to an atom which is more electronegative than hydrogen and another atom in the same or a different molecule
How do London forces occur?
- the electron density of a non-polar molecule fluctuates due to kinetic energy
- this causes an induced dipole in the next molecule which causes a London force
When does the force of a London force increase?
- when there are more electrons in the molecule, it causes there to be greater fluctuations in electron density so the induced dipoles and instantaneous dipoles are larger
- more points of contact, the larger the London forces e.g. shape and size of molecules
When are London forces present?
in every molecule- permanent dipole or not
When are permanent dipoles present?
if the molecule is polar- the electronegativities -differ
When does hydrogen bonding occur?
when the atom bonded to hydrogen is more electronegative than hydrogen, mainly with oxygen, nitrogen or fluorine
