topic 2: bonding and structure Flashcards
what is the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs and some examples
shape: linear
bond angle: 180
examples: CO2, CS2, HCN, BeF2
what is the shape and bond angle of a molecule with 3 bonding pairs and 0 lone pairs and some examples
shape: trigonal planar
bond angle: 120
examples:BF3, AlCl3, SO3, NO3-, CO3 2-
what is the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs and some examples
shape: tetrahedral
bond angle: 109.5
examples: SiCl4, SO42-, ClO4-, NH4+
what is the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pairs and some examples
shape: trigonal pyramidal
bond angle: 107
examples: NCl3 ,PF3,ClO3
,H3O+
what is the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs and some examples
shape: bent
bond angle: 104.5
examples: OCl2, H2S, OF2, SCl2
what is the shape and bond angle of a molecule with 5 bonding pairs and 0 lone pairs and some examples
shape: trigonal bipyramidal
bond angle: 120 and 90
example: PCl5
what is the shape and bond angle of a molecule with 6 bonding pairs and 0 lone pairs and some examples
shape: octahedral
bond angle: 90
example: SF6
what is ionic bonding
strong electrostatic force of attraction
between oppositely charged ions formed by electron transfer
why are positive ions smaller in comparison to their atoms
because it has one less shell of electrons and the ratio of protons to electrons has increased so there is greater net force on remaining electrons holding them more closely
why are anions larger than their atoms
the negative ion has more electrons than the corresponding atom but the same number of protons. so the pull of the nucleus is shared over more electrons and the attraction per electron is less, making the ion bigger
does the ionic radius increase or decrease from N3- to Al3+ and why
decrease because there are increasing numbers of protons from N to F and then Na to Al but the same number of electrons. the effective nuclear attraction per electron therefore increases and ions get smaller
does the ionic radius increase or decrease down the group and why
increases down group
because as one goes down the group the ions have more shells of electrons
what are the physical properties of ionic compounds
*high melting points ( there are strong attractive forces between the ions)
*non conductor of electricity when solid (ions are held together tightly and can not move)
*conductor of electricity when in solution or molten. ( ions are free to move)
*brittle / easy to cleave apart
what is covalent bond
strong electrostatic attraction between 2 nuclei and the shared/ bonding pair of electrons
what is the strength of covalent bond demonstrated by
high melting points
why do giant covalent structures have high melting points
because they contain many
strong covalent bonds in a macromolecular structure. it takes a lot of energy to break the many
strong bonds
what is the effect of multiple bonds on bond strength and length
nuclei joined by multiple (i.e. double and triple) bonds have a greater electron density between them. this causes an greater force of attraction between the nuclei and the electrons between them, resulting in a
shorter bond length and greater bond strength
what is a dative bond and examples
when the shared pair of electrons in the covalent bond come from only one of the bonding atoms. a dative covalent bond is also called co-ordinate bonding
eg. NH4+, H30+, NH3BF3
what is electronegativity
the relative tendency of an atom in a covalent bond in a molecule to attract electrons in a covalent bond to itself is measured on the Pauling scale
why do electronegativity increases across a period
the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more
why do electronegativity decreases down the group
because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases
what does a small electronegative difference mean
purely covalent
what would the electronegative difference when ionic
very large electronegativity difference (> 1.7)
describe the formation of a permanent dipole
a polar covalent bond forms when the elements in the bond have different electronegativities . (Of around 0.3 to 1.7)
When a bond is a polar covalent bond it has an unequal distribution of electrons in the bond and produces a charge separation, (dipole) δ+ δ- ends
what would a symmetrical molecule do to the polar
a symmetric molecule (all bonds identical and no lone pairs) will not be polar even if individual bonds within the molecular are polar
the individual dipoles on the bonds ‘cancel out’ due to the symmetrical shape of the molecule
there is no net dipole moment: the molecule is non-polar
when do London forces occur
occur between all molecular substances and noble gases
they do not occur in ionic substances
what are induced dipoles
temporary dipoles can cause dipoles to form in neighbouring molecules
what is the main factor affecting size of London forces
the more electrons there are in the molecule the higher the chance that temporary dipoles will form
this makes the London forces stronger between the molecules and more energy is needed to break them so boiling points will be greater
why do the boiling points increase down group 7
increasing number of electrons in the bigger molecules causing an increase in the size of the London
forces between the molecules
why do the boiling points increase through the alkane homologous series
increasing number of electrons in the bigger molecules causing an increase in the size of the London forces between molecules
describe permanent dipole - dipole forces
- permanent dipole-dipole forces occurs between polar molecules
- it is stronger than London forces and so the compounds have higher boiling points
- polar molecules have a permanent dipole. (commonly compounds with C-Cl, C-F, C-Br H-Cl, C=O bonds)
- polar molecules are asymmetrical and have a bond where there is a significant difference in electronegativity between the atoms
describe hydrogen bonding
occurs in compounds that have a hydrogen atom attached to one of the three most electronegative atoms of nitrogen, oxygen and fluorine, which must have an available lone pair of electrons. e.g. a –O-H -N-H F- H bond. there is a large electronegativity difference between the H and the O,N,F
what is metallic bonding
the electrostatic force of attraction between the positive metal ions and the delocalised electrons
what are the 3 factors that affect metallic bonding
number of protons
number of delocalised electrons per atom
size of ion
why do metals have a high melting point
because the strong electrostatic forces between positive ions
and sea of delocalised electrons require a lot of energy to break
why can metals conduct electricity
because the delocalised electrons can move through the structure
why are metals malleable
because the positive ions in the lattice are all identical. so the planes of ions can slide easily over one another. the attractive forces in the lattice are the same whichever ions are adjacent
describe the structure of diamond
Tetrahedral arrangement of carbon atoms. 4 covalent
bonds per atom diamond cannot conduct electricity because all 4 electrons per carbon atoms are involved in covalent bonds. they are localised and cannot move
describe the structure of graphite
planar arrangement of carbon
atoms in layers. 3 covalent bonds
per atom in each layer. 4th outer
electron per atom is delocalised.
delocalised electrons between
layers
graphite can conduct electricity well between layers because one electron per carbon is free and delocalised, so electrons can move easily along
layers
it does not conduct electricity from one layer to the next because the energy gap between layers is too
large for easy electron transfer
describe the structure of graphene
graphene is a new substance that is a one layer of graphite .i.e. 3 covalent bonds per atom and
the 4th outer electron per atom is delocalised
these have very high tensile strength because of the strong structure of many strong covalent
bonds
graphene can conduct electricity well along the structure because one electron per carbon is free
and delocalised, so electrons can move easily along the structure
describe the structure of carbon nanotubes
These have very high tensile strength because of the strong
structure of many strong covalent bonds.
Nanotubes can conduct electricity well along the tube
because one electron per carbon is free and delocalised,
so electrons can move easily along the tube
nanotubes have potentially many uses. one being the potential to use as vehicles to deliver drugs to cell
there are delocalised electrons in buckminsterfullerene
what are the properties of ionic
high melting and boiling points -because of giant lattice of ions with
strong electrostatic forces between
oppositely charged ions
solubility in water - generally good
poor conductivity when solid: ions
can’t move/ fixed in lattice
good conductivity when molten: ions can move
what are properties of molecular simple molecule
low melting and boiling point - because of weak intermolecular
forces between molecules (specify type e.g London forces/hydrogen
bond)
solubility in water - generally poor
poor conductivity when solid: no ions to conduct and electrons are
localised (fixed in place)
poor conductivity when molten: no ions
what are the properties of a macromolecular molecule
high melting and boiling point - because of many strong covalent bonds in macromolecular structure. take a lot of energy to break the many strong bonds
insoluble in water
conductivity when solid - diamond and sand: poor, because electrons can’t move (localised) graphite: good as free delocalised electrons between layers
poor conductivity when molten
what are the properties of a metallic molecule
high melting and boiling point - strong electrostatic forces between positive ions and sea of delocalised electrons
insoluble in water
good conductivity when solid: delocalised electrons can move through structure
good conductivity when molten
what are the possible reasons why there are 2 widely different values for compressive strength of graphite (both are valid)
lower values relate to the weak London forces
high values relate to the strong covalent bonds
what is a possible reason as to why the density of iron is much higher than the density of graphite
iron atoms have a greater mass than carbon atoms
iron atoms pack closer together than carbon atoms
ammonia and boron trifluoride react to form a compound NH3BF3 which contains a dative
covalent bond. Each of the molecules, NH3 and BF3, has a different feature of its electronic
structure that allows this to happen. Use these two different features to explain how a dative
covalent bond is formed
donation of lone pair from nitrogen to boron which has only has 6 electrons in outer shell
during this reaction, the bond angles about the nitrogen atom and the boron atom
change. state the new H—N—H and F—B—F bond angles
both 109.5 bc there is now 4 BP and 0LP
what is electronegativity
the relative ability of an atom to attract the bonding electrons in a covalent bond
explain why both water and carbon dioxide molecules have polar bonds but only
water is a polar molecule
- O is more electonegative than H and C
- which results in a polar bond with O dipole -ve so C and H dipole +ve
- CO2 is a symmetrical molecule and so the dipole moments cancel out
- the V shape of water molecule/ the lone pairs of electrons of O means that the dipole moments don’t cancel
explain, in terms of the structure and bonding of each element, the difference between these
values - silicon and chlorine
Si is a giant covalent structures and contains many strong covalent bonds
Cl2 is a simple molecular, diatomic molecule and contains weak London forces
covalent bonds in Si are stronger than London forces in chlorine so covalent bonds take more energy to break than London forces
explain why diamond has a much higher melting temperature than iodine
iodine is simple molecular
- molecules are held together by weak London forces
diamond is a giant covalent structure with each carbon covalently bonded to 4 other carbons
- carbon atoms in diamond are held together by strong covalent bonds
strong covalent bonds require more energy to break than intermolecular forces
Aluminium fluoride and aluminium chloride are both crystalline solids at room temperature.
Aluminium fluoride sublimes to form a gas at 1291°C (1564 K), whilst aluminium chloride
sublimes at 178°C (451 K).
Use the Pauling electronegativity values in the Data Booklet to explain these differences in
sublimation temperature
al and cl electronegative difference 1.5
al and f electronegative difference 2.5
aluminium chloride mostly covalent
aluminium fluorine bonds are more polar
aluminium fluoride has strong electrostatic forces of attraction between the ions
aluminium chloride is molecular so has weak London forces
more energy needed to break the strong bonds to cause sublimation in aluminium fluoride
