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1 3 Bonding chemrevise
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Bonding and Structure,Bonding Structure Examples,Ionic electrostatic force of Sodium chloride. attraction between Magnesium oxide,Giant Ionic Lattice. oppositely charged ions,Covalent shared pair of Simple molecular Iodine. electrons With intermolecular forces van der Ice,Waals permanent dipoles hydrogen Carbon dioxide. bonds between molecules Water,Covalent shared pair of Macromolecular Diamond.
electrons giant molecular structures Graphite,Silicon dioxide. Metallic electrostatic Magnesium Sodium, force of attraction between Giant metallic all metals. the metal positive ions and lattice,the delocalised electrons. Only use the words molecules and intermolecular forces when talking about simple molecular substances. Property Ionic Molecular simple Macromolecular Metallic. boiling and high because low because of high because of high strong electrostatic forces. melting of giant lattice weak intermolecular many strong covalent between positive ions and sea of. points of ions with forces between bonds in delocalised electrons. strong molecules specify macromolecular, electrostatic type e g van der structure Take a lot. forces waals hydrogen of energy to break the,between bond many strong bonds.
oppositely,charged ions, Solubility in Generally generally poor insoluble insoluble. water good, conductivity poor ions poor no ions to diamond and sand good delocalised electrons can. when solid can t move conduct and poor because move through structure. fixed in lattice electrons are electrons can t move. localised fixed in localised,place graphite good as free. delocalised electrons,between layers,conductivity good ions can poor no ions poor good. when molten move, general crystalline mostly gases and solids shiny metal.
description solids liquids Malleable as the positive ions in. the lattice are all identical So the,planes of ions can slide easily. over one another,attractive forces in the lattice are. the same whichever ions are,N Goalby chemrevise org 2. Shape of molecules,Name No No lone Diagram Bond angle Examples. bonding pairs,linear 2 0 Cl Be Cl 180 CO2 CS2 HCN,Trigonal 3 0 Cl Cl 120 BF3 AlCl3 SO3.
planar NO3 CO32,Tetrahedral 4 0 H 109 5 SiCl4 SO42 ClO4. Trigonal 3 1 107 NCl3 PF3 ClO3,pyramidal H3O,Bent 2 2 104 5 OCl2 H2S OF2. Trigonal 5 0 F 120 and 90 PCl5,Bipyramidal F,Octahedral 6 0 90 SF6. How to explain shape 1 State number of bonding pairs and lone pairs of electrons. 2 State that electron pairs repel and try to get as far apart as possible or to a. position of minimum repulsion, 3 If there are no lone pairs state that the electron pairs repel equally. 4 If there are lone pairs of electrons then state that lone pairs repel more than. bonding pairs,5 State actual shape and bond angle, Remember lone pairs repel more than bonding pairs and so reduce bond angles by about 2 5o.
per lone pair in above examples,N Goalby chemrevise org 3. Occasionally more complex shapes are seen that are variations of octahedral and trigonal. bipyramidal where some of the bonds are replaced with lone pairs You do not need to learn the. names of these but ought to be able to work out these shapes using the method below. Square planar Bond angle 89O Bond angle 89O Bond angles 119 89O. Bond angle 180O, Bond angle 90O Reduced by lone pair Reduced by lone pairs Reduced by lone pair. e g XeF4 e g BrF5 e g I3 e g ClF3 e g SF4 IF4, Xe has 8 electrons in its outer Cl has 7 electrons in its outer I has 7 electrons in its outer. shell 4 F s add 4 more shell 3 F s add 3 more shell 4 F s add 4 more. electrons This makes a total of electrons This makes a total of electrons Remove one electron. 12 electrons made up of 4 10 electrons made up of 3 bond as positively charged This. bond pairs and 2 lone pairs pairs and 2 lone pairs The makes a total of 10 electrons. The means it is a variation of means it is a variation of the 5 made up of 4 bond pairs and 1. the 6 bond pair shape bond pair shape trigonal lone pair The means it is a. octahedral bipyramidal variation of the 5 bond pair. shape trigonal bipyramidal,N Goalby chemrevise org 4. Electronegativity and intermediate bonding,Definition F O N and Cl are the.
most electronegative, Electronegativity is the relative tendency of an atom in a covalent bond atoms. in a molecule to attract electrons in a covalent bond to itself. The most electronegative, Electronegativity is measured on the Pauling scale ranges from 0 to 4 element is fluorine and it. is given a value of 4 0,Factors affecting electronegativity. Electronegativity increases across a period as the number of protons increases and the atomic radius. decreases because the electrons in the same shell are pulled in more. It decreases down a group because the distance between the nucleus and the outer electrons increases and. the shielding of inner shell electrons increases,Intermediate bonding. Ionic and covalent bonding are the extremes of a continuum of bonding type Differences in. electronegativity between elements can determine where a compound lies on this scale. A compound containing elements of similar electronegativity and hence a small. electronegativity difference will be purely covalent. A compound containing elements of very different electronegativity and hence a. very large electronegativity difference 1 7 will be ionic. Formation of a permanent dipole polar covalent bond The element with the. A polar covalent bond forms when the elements in the bond have different electronegativity in a. electronegativities Of around 0 3 to 1 7 polar compound will. be the end, When a bond is a polar covalent bond it has an unequal distribution of d d.
electrons in the bond and produces a charge separation dipole H Cl. Polar and Non Polar molecules,Symmetric molecules, A symmetric molecule all bonds identical and no lone. pairs will not be polar even if individual bonds within. the molecular ARE polar,CO2 is a symmetrical molecule and. The individual dipoles on the bonds cancel out,is a non polar molecule. due to the symmetrical shape of the molecule,There is no NET dipole moment the molecule is. e g CCl4 will be non polar whereas CH3Cl will be polar H. N Goalby chemrevise org 5,Intermolecular Forces, Van der Waals Forces Van der Waals forces occur between all molecular substances and noble gases.
They do not occur in ionic substances, These are also called transient induced dipole dipole interactions They occur between all simple. covalent molecules and the separate atoms in noble gases. In any molecule the electrons are moving constantly and randomly As this happens the electron density. can fluctuate and parts of the molecule become more or less negative i e small temporary or transient. dipoles form, These instantaneous dipoles can cause dipoles to form in neighbouring molecules These are called. induced dipoles The induced dipole is always the opposite sign to the original one. Main factor affecting size of Van der waals, The more electrons there are in the molecule the higher the chance that temporary dipoles will form This. makes the van der Waals stronger between the molecules and so boiling points will be greater. The increasing boiling points of the halogens down the group 7 series can be explained by the. increasing number of electrons in the bigger molecules causing an increase in the size of the van der. Waals between the molecules This is why I2 is a solid whereas Cl2 is a gas. The increasing boiling points of the alkane homologous series can be explained by the increasing. number of electrons in the bigger molecules causing an increase in the size of the van der Waals. between molecules, The shape of the molecule can also have an effect on the size of the van der Waals forces Long chain. alkanes have a larger surface area of contact between molecules for van der waals to form than compared to. spherical shaped branched alkanes and so have stronger van der waals. Permanent dipole dipole forces, Permanent dipole dipole forces occurs between polar molecules.
It is stronger than van der waals 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,Permanent dipole forces. occurs in addition to van,der waals forces,Hydrogen bonding. It 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,Always show,the lone pair of,electrons on the. O F N and the,dipoles and all, Hydrogen bonding occurs in addition to van der waals forces. N Goalby chemrevise org 6,Hydrogen bonding is stronger than the other two.
types of intermolecular bonding,Boiling point K,300 HF H2Te. The anomalously high boiling points of H2O H2Se SbH3. NH3 and HF are caused by the hydrogen NH3 H2S,bonding between the molecules AsH3. PH3 HCl SnH4, The general increase in boiling point from H2S GeH4. to H2Te is caused by increasing van der Waals, forces between molecules due to an increasing 100 CH4. number of electrons,Alcohols carboxylic acids proteins amides all.
can form hydrogen bonds 25 50 75 100 125,Molecular mass. Four types of crystal structure ionic metallic molecular and giant covalent macromolecular. You should be able to draw the following diagrams or describe the structure in words to. show the four different types of crystal You should also be able to explain the properties of. these solids The tables earlier in the revision guide explain these properties. Ionic sodium chloride,Giant Ionic lattice showing alternate. Na and Cl ions,Metallic magnesium or sodium,Giant metallic lattice showing close. packing magnesium ions,metal Molecular Ice This is a difficult diagram. Molecular Iodine H H The main point to show is,O a central water molecule.
with two ordinary covalent,bonds and two hydrogen, Regular arrangement of I2 H bonds in a tetrahedral. molecules held together by H arrangement,weak van der Waals forces H. H H The molecules are held,HO O further apart than in liquid. H H water and this explains the,lower density of ice. Macromolecular diamond Macromolecular Graphite,Planar arrangement of carbon.
Tetrahedral arrangement of atoms in layers 3 covalent bonds. per atom in each layer 4th outer,carbon atoms 4 covalent. electron per atom is delocalised,bonds per atom Delocalised electrons between. Both these macromolecular structures have very high melting points because of strong covalent. forces in the giant structure It takes a lot of energy to break the many strong covalent bonds.

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