topic 2: bonding and structure

Question 1

what is the shape and bond angle of a molecule with 2 bonding pairs and 0 lone pairs and some examples

Answer 1

shape: linear
bond angle: 180
examples: CO2, CS2, HCN, BeF2

Question 2

what is the shape and bond angle of a molecule with 3 bonding pairs and 0 lone pairs and some examples

Answer 2

shape: trigonal planar
bond angle: 120
examples:BF3, AlCl3, SO3, NO3-, CO3 2-

Question 3

what is the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs and some examples

Answer 3

shape: tetrahedral
bond angle: 109.5
examples: SiCl4, SO42-, ClO4-, NH4+

Question 4

what is the shape and bond angle of a molecule with 3 bonding pairs and 1 lone pairs and some examples

Answer 4

shape: trigonal pyramidal
bond angle: 107
examples: NCl3 ,PF3,ClO3
,H3O+

Question 5

what is the shape and bond angle of a molecule with 2 bonding pairs and 2 lone pairs and some examples

Answer 5

shape: bent
bond angle: 104.5
examples: OCl2, H2S, OF2, SCl2

Question 6

what is the shape and bond angle of a molecule with 5 bonding pairs and 0 lone pairs and some examples

Answer 6

shape: trigonal bipyramidal
bond angle: 120 and 90
example: PCl5

Question 7

what is the shape and bond angle of a molecule with 6 bonding pairs and 0 lone pairs and some examples

Answer 7

shape: octahedral
bond angle: 90
example: SF6

Question 8

what is ionic bonding

Answer 8

strong electrostatic force of attraction
between oppositely charged ions formed by electron transfer

Question 9

why are positive ions smaller in comparison to their atoms

Answer 9

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

Question 10

why are anions larger than their atoms

Answer 10

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

Question 11

does the ionic radius increase or decrease from N3- to Al3+ and why

Answer 11

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

Question 12

does the ionic radius increase or decrease down the group and why

Answer 12

increases down group
because as one goes down the group the ions have more shells of electrons

Question 13

what are the physical properties of ionic compounds

Answer 13

*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

Question 14

what is covalent bond

Answer 14

strong electrostatic attraction between 2 nuclei and the shared/ bonding pair of electrons

Question 15

what is the strength of covalent bond demonstrated by

Answer 15

high melting points

Question 16

why do giant covalent structures have high melting points

Answer 16

because they contain many
strong covalent bonds in a macromolecular structure. it takes a lot of energy to break the many
strong bonds

Question 17

what is the effect of multiple bonds on bond strength and length

Answer 17

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

Question 18

what is a dative bond and examples

Answer 18

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

Question 19

what is electronegativity

Answer 19

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

Question 20

why do electronegativity increases across a period

Answer 20

the number of protons increases and the atomic radius decreases because the electrons in the same shell are pulled in more

Question 21

why do electronegativity decreases down the group

Answer 21

because the distance between the nucleus and the outer electrons increases and the shielding of inner shell electrons increases

Question 22

what does a small electronegative difference mean

Answer 22

purely covalent

Question 23

what would the electronegative difference when ionic

Answer 23

very large electronegativity difference (> 1.7)

Question 24

describe the formation of a permanent dipole

Answer 24

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

Question 25

what would a symmetrical molecule do to the polar

Answer 25

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

Question 26

when do London forces occur

Answer 26

occur between all molecular substances and noble gases they do not occur in ionic substances

Question 27

what are induced dipoles

Answer 27

temporary dipoles can cause dipoles to form in neighbouring molecules

Question 28

what is the main factor affecting size of London forces

Answer 28

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

Question 29

why do the boiling points increase down group 7

Answer 29

increasing number of electrons in the bigger molecules causing an increase in the size of the London forces between the molecules

Question 30

why do the boiling points increase through the alkane homologous series

Answer 30

increasing number of electrons in the bigger molecules causing an increase in the size of the London forces between molecules

Question 31

describe permanent dipole - dipole forces

Answer 31

* 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

Question 32

describe hydrogen bonding

Answer 32

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

Question 33

what is metallic bonding

Answer 33

the electrostatic force of attraction between the positive metal ions and the delocalised electrons

Question 34

what are the 3 factors that affect metallic bonding

Answer 34

number of protons number of delocalised electrons per atom size of ion

Question 35

why do metals have a high melting point

Answer 35

because the strong electrostatic forces between positive ions and sea of delocalised electrons require a lot of energy to break

Question 36

why can metals conduct electricity

Answer 36

because the delocalised electrons can move through the structure

Question 37

why are metals malleable

Answer 37

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

Question 38

describe the structure of diamond

Answer 38

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

Question 39

describe the structure of graphite

Answer 39

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

Question 40

describe the structure of graphene

Answer 40

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

Question 41

describe the structure of carbon nanotubes

Answer 41

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

Question 42

what are the properties of ionic

Answer 42

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

Question 43

what are properties of molecular simple molecule

Answer 43

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

Question 44

what are the properties of a macromolecular molecule

Answer 44

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

Question 45

what are the properties of a metallic molecule

Answer 45

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

Question 46

what are the possible reasons why there are 2 widely different values for compressive strength of graphite (both are valid)

Answer 46

lower values relate to the weak London forces high values relate to the strong covalent bonds

Question 47

what is a possible reason as to why the density of iron is much higher than the density of graphite

Answer 47

iron atoms have a greater mass than carbon atoms iron atoms pack closer together than carbon atoms

Question 48

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

Answer 48

donation of lone pair from nitrogen to boron which has only has 6 electrons in outer shell

Question 49

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

Answer 49

both 109.5 bc there is now 4 BP and 0LP

Question 50

what is electronegativity

Answer 50

the relative ability of an atom to attract the bonding electrons in a covalent bond

Question 51

explain why both water and carbon dioxide molecules have polar bonds but only water is a polar molecule

Answer 51

- 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

Question 52

explain, in terms of the structure and bonding of each element, the difference between these values - silicon and chlorine

Answer 52

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

Question 53

explain why diamond has a much higher melting temperature than iodine

Answer 53

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

Question 54

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

Answer 54

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