As chemistry unit one point 3

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As chemistry unit one point 3
2013-12-23 04:07:19

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  1. Metallic bonding involves 
    oppositely charged ions in a lattice
  2. A covalent bond involves a
    Shared pair of electrons
  3. Coordinate bonding is a
    Dative covalency
  4. Metallic bonding involves
    involves a lattice of positive ions surrounded by delocalised electrons 
  5. electro negativity is 
    the power of an atom to withdraw electRon density from a covalent Bond 
  6. the electron distribution of in a covalent bond may not be 
  7. covalent bonds between different elements will be polar to 
    different extents 
  8. molecules interact by 
    permanent dipole-dipole , induced dipole-dipole forces and hydrogen bonding 
  9. hydrogen bonding is important in the determination of 
    boiling points of compounds and the structure of some solids such as ice 
  10. there are .... changes associated with changes of state 
  11. there are four types of crystal
    ionic , metallic , giant covalent (macromolecular) and molecular 
  12. the spec says we must know the structure of the following crystals 
    • sodium chloride
    • magnesium
    • diamond
    • graphite
    • iodine
    • ice  
  13. the physical properties of materials are related to
    the type of structure and bonding present 
  14. the spec says we must understand the concept of bonding and lone pairs of electrons as 
    charge clouds 
  15. the spec says that we should be able to , in terms of electron repulsion , predict the shapes of , and bond angles in , simple 
    molecules and ions , limited to 2 , 3 , 4, 5 and 6 coordination 
  16. lone pair/lone pair repulsion is greater than 
    lone/pair boding repulsion , which is greater than bonding pair/boding pair repulsion 
  17. all chemical bonds are
    forces of attraction 
  18. ionic bonds occur when
    when atoms lose or gain electrons which which is when a metal and non metal react to form a compound
  19. during ionic bonding metal atoms
    • lose electrons forming cations
    • for example a lithium atom loses 1 electron to form a lithium ion 
    • Li ---> Li+ e
    • the lithium ion has electronic configuration 1s2
  20. the lithium ion is
    isolectronic (has the same electronic configuration) with helium
  21. duirng ionic bonding non metal atoms
    • gain electrons forming anions . 
    • for example a fluorine atom gains an electron to form a fluoride ion . 
    • F + e---> F
  22. the fluoride ion has electronic configuration 
  23. the fluoride ion is
    isoelectronic with neon
  24. an ionic bond is
    the electrostatic attraction between cations and anions
  25. the ions in an ionic compound form a repeating 3D structure called a
  26. magnesium chloride ionic bonding diagram
  27. aluminium fluoride ionic bonding diagram
  28. ionic compounds are formed of
    lattice structures
  29. in an ionic compound each ion is surrounded by
    ions of opposite charge , this structure is repeated throughout the ionic compound , as a result ionic compounds are said to have giant structures
  30. ionic compounds are not
  31. you cannot write a ..... .... for an ionic compound .
    molecular formula
  32. however you can write an empirical formula for an ionic compound . this shows
    the ratio of cations to anions present in the ionic lattice
  33. a formula for an ionic compound is constructed from
    the formulae of the cations and anions present in the ionic lattice . in most cases you work out the charges of cation and anion using the periodic table
  34. for metal atoms the charge on the ion is the
    same as the group number of the metal
  35. for non metal atoms the charge on the ion is the same as
    the group number of the atom minus 8
  36. hydrogen ion
  37. silver ion
  38. ammonium ion
    • NH4+
  39. copper ion
    Cu2+  or Cu3+
  40. Zinc ion
  41. lead ion
  42. iron ion
    Fe2+ or Fe3+
  43. hydroxide ion
  44. nitrate ion
  45. Oxide ion
  46. sulphide ion
  47. sulphate ion
  48. Carbonate ion
  49. phosphate ion
  50. Chromate ions
  51. Ionic compounds have a .... charge overall so the number of positive charges must match the number of negative charges
  52. aluminium sulphate formula
    • Al3+ ions and SO4 2- ions 
    • Al2(SO4)3
  53. Property of ionic compounds - Do not conduct
    electricity when solid

    explanation -
    ions are fixed in position by strong ionic bonds so are not able to move and carry charge however when molten or dissolved in water ions are free to move
  54. property of ionic compounds - high melting and boiling points 

    explanation -
    giant lattice structure held together by strong forces of electrostatic attraction between oppositely charged ions which requires a lot of energy to overcome
  55. property of ionic compounds - brittle and shatter easily 

    explanation -
    small displacement causes contact between ions with the same charge , the ions therefore repel and the structure shatters
  56. property of ionic compounds - dissolve in water 

    explanation -
    polar water molecules pull ions away from lattice and cause it to dissolve
  57. does LiCl or KCl have the highest melting point
    • LiCl because both compounds have a chloride ion 
    • however Liis a smaller ion (same charge) and so has a higher charge density . This means that there is a greater force of attraction between cations and anions meaning LiCl requires more energy to overcome its forces of attraction 
  58. does Na2O or MgO have the highest melting point
    Both have O2- ions however MgO has Mg2+ ions whereas Na2O has Naions

    Mg2+ evidently has a higher charge and is also a smaller ion and has a higher charge density .  This means that there is a greater force of attraction between cations and anions meaning MgO requires more energy to overcome its forces of attraction
  59. atoms of metals are held together by
    metallic bonds 
  60. a metallic bond is
    the electrostatic force of attraction between metal ions and the sea delocalised electrons in a metallic lattice . the delocalised electrons come from the highest energy level
  61. why are metals good conductors of electricity
    • in a metallic bond 
    • each metal ion forms a positive ion
    • the positive ions are arranged in a lattice structure 
    • the ions in the structure are very close together so the electrons that are lost when the metal forms ions are delocalised 
    • delocalised electrons aren't attracted to any particular ion 
    • as a reasult these electrons are free to move when a voltage is applied . Pushing electrons in one end of the piece of metal causes electrons to come out of the other end
  62. the ionic lattice is a
    • regular arrangement of ions of opposite charge
  63. the metallic lattice is a
    • regular arrangement of cations surrounded by a sea of delocalised electrons
  64. metallic bonds don't have the same strength . if they did , then all metals would melt and boil at the same temperature . the strength of the metallic bond is dependant on
    • the charge (density) of the ions in the lattice -the larger the charge (density) the stronger the bond 
    • the number of electrons in the sea of delocalised electrons - the more electrons there are the stronger the bond
  65. Metallic bond strength:
    • Increases across a period as more electrons become delocalized 
    • Decreases down a group as the atomic radius increases - (the charge density decreases)
  66. magnesium has a higher boiling point than sodium because
    the attraction between the Mg2+ ions in the lattice and the sea of delcoalised electrons is greater than between the Naions in the lattice and the sea of delocalised electrons . this is because Mg2+ has a higher charge and more delocalised electrons
  67. lithium has a higher melting point than sodium
    because there is a stronger metallic bond in lithium because lithium has a higher charge density so there are stronger forces of electrostatic attraction between the metal ions and sea of delocalised electrons
  68. aluminium has a higher boiling point than sodium because
    alumnium has a higher charge density and more delocalised electrons . so there are stronger forces of electrostatic attraction between the metal ions and sea of delocalised electrons in aluminium so it has stronger metallic bonds which require more energy to overcome
  69. the cations formed in metallic bonding arrange themselves
    in a close packed structure
  70. there are two forms of close packed structure
    hexagonal close packed (HCP) and simple cubic
  71. in the simple cubic structure the ions are
    stacked directly on top of each other
  72. the hexagonal close packed structure is more
    efficient with 74% of the volume fixed with particles
  73. explain the HCP structure
    • the ions in the first layer are arranged so that they are touching each other
    • in the second layer the ions sit in the hollows between the ions in the first layer 
    • in the third layer the ions are in hollows in the second layer , directly over ions in the first layer
  74. the HCP structure can be shown as either a
    space filling model or unit cell

  75. properties of metal - good conductor of heat 
    explanation - 
    the metal ions in the lattice are very close to each other . heating one end of a piece of metal makes the ions at that end vibrate more , and these vibrations are passed along the piece of metal 
  76. property of metals - high melting points and boiling points 
    explanation - 
    metals have a giant structure and there are strong forces of electrostatic attraction between the cations and sea of delocalised electrons 
  77. properties of metals - malleable and ductile 
    explanation - 
    No bonds holding ions together and ions can slide over each other if a large enough force is applied . the strength of metallic bond stops the attraction being broken completely 
  78. why are metals strong 
    because of their metallic bonding , there are strong forces of electrostatic attraction between cations and the sea of delocalised electrons 
  79. why are metals insoluble (except in liquid metals)
    strength of metallic bonds
  80. non metal atoms can achieve full outer shells either by
    accepting electrons from metal atoms or sharing pairs of electrons 
  81. covalent bonds can form between 
    identical atoms or different atoms 
  82. as with ionic bonding only the ..... ..... .... .... are involved in covalent bonding 
    outer energy level electrons 
  83. in covalent bonding a 
    pair of electrons is shared between two atoms 
  84. a covalent bond is held together by 
    the attraction between each nucleus involved in the bond and the pair of electrons 
  85. covalent bonds can be represented by
    dot and cross diagrams . they can also be shown as a straight line between the atoms , this is called displayed formula 
  86. non metal can form double or triple covalent bonds by
    sharing more than one pair of electrons . these are represented in molecular formulae by multiple lines between atoms . double bonds are stronger than single bonds and triple bonds are stronger than single bonds 
  87. single covalent bonds 
    a shared pair of electrons
  88. double covalent bond
    two shared pair of electrons 
  89. triple covalent bond 
    three shared pairs of electrons 
  90. examples of double covalent bonds 
    • oxygen 
    • carbon monoxide 
  91. examples of triple covalent bonds 
  92. pairs of electrons that are not involved in bonding are called
    lone pairs of electrons 
  93. why are lone pairs important 
    • they are important in determining the shape of molecules 
    • they can also influence the chemical properties of molecules 
  94. why does ammonia have a lone pair of electrons 
    • nitrogen is in group 5 and form 3 single bonds with hydrogen 
  95. why does water have two lone pairs 
    • since oxygen is in group 6 and forms two covalent bonds with hydrogen 
  96. lone pairs of electrons are able to form ...... (or ....) bonds with atoms that have vacant orbitals 
    • co-ordinate 
    • dative 
  97. a coordinate bond is a 
    shared pair of electrons in which both electrons are contributed by one of the atoms in the bond 
  98. coordinate bonds are shown in displayed formula by an 
  99. co-ordinate bonds form between an 
    atom with an empty orbital and an atom with a lone pair of electrons . these bonds can only occur in simple molecular substances 
  100. ammonium ion coordinate covalent bond diagram 
    • the coordinate bond is is between the nitrogen atom and one of the hydrogen atoms . it has formed because nitrogen has a lone pair of electrons and hydrogen is able to gain a vacant orbital 
  101. water forms dative coordinate bonds
    • a coordinate bond forms because oxygen has a lone pair of electrons and hydrogen has a vacant orbital  
  102. in carbon monoxide the oxygen atom forms a double covalent bond and also a coordinate bond 
    • a coordinate bond forms because oxygen has a lone pair of electrons and carbon has a vacant orbital  
  103. molecules such as O2 , NH, CO2 and H2O are described as 
    simple molecules . these are discrete molecules made from a small number of atoms . 
  104. why do simple covalent substances not conduct electricity
    there are no ions or delocalised electrons involved 
  105. why do simple covalent substances have low melting and boiling points 
    there are weak intermolecular forces 
  106. Iodine is an example of a molecular crystal The iodine atoms pair up to form I2 molecules, held together by ------------ . ....................................  hold the crystal
    • strong covalent bonds
    • Intermolecular forces between the I2 molecules
  107. covalent bonding can form 
    simple covalent (molecular) or giant covalent structures (macromolecular) note we have just been looking at simple covalent
  108. carbon forms a number of structures called allotropes these are 
    different structures of the same element 
  109. each carbon atom can form .... covalent bonds 
  110. this ability to form a number of bonds enables the carbon to 
    bond to itself 
  111. there are two main macromolecular substances 
    diamond and graphite both of which are allotropes of carbon 
  112. a giant macromolecule is one in which the same 
    arrangement of atom is repeated many times 
  113. explain the structure of diamond 
    • each carbon atom forms 4 covalent bonds 
    • the shape around each carbon atom is tetrahedral 
    • the C-C bond angle about each atom is 109.5O
    • the tetrahedral structure is repeated around each carbon atom making diamond extremely hard 
  114. explain the structure of graphite 
    • each carbon atom forms three covalent bonds 
    • the shape around each atom is trigonal planar 
    • the C-C bond angle about each atom is 120o
    • the carbon atoms form a flat lattice structure made from hexagons 
    • the extra electron from each carbon is contributed to a delocalised sea of electrons between the layers 
  115. why is graphite soft and slippery and used in pencils 
    van der waals forces exist between the layers . these weak forces allow the layers to slide over each other 
  116. why is graphite strong and lightweight
    it's strong because there are strong covalent bonds between the carbon atoms in each layer and its lightweight because the layers are quite far apart which gives it a low density 
  117. why does graphite conduct electricity 
    • Graphite has 1 delocalized electron per carbon atom as it only forms three bonds so can
    • conduct electricity along the hexagonal sheets. 
  118. why does graphite have high melting and boiling points
    because there are strong covalent bonds in hexagonal sheets , this requires a large amount of energy to overcome 
  119. why is graphite insoluble 
    there are covalent bonds which are difficult to break 
  120. why does diamond have a high melting point 
    strong covalent bonds in a crystal lattice structure  
  121. why is diamond hard
    rigid crystal structure 
  122. why doesn't diamond conduct electricity
    no delocalised electrons or ions 
  123. why is diamond insoluble 
    covalent bonds are too difficult to break 
  124. covalent bonding is the 
    sharing of one or more pairs of electrons 
  125. the covalent bond is held together by the attraction between 
    the nuclei of the two atoms involved in the bond and the pairs of electrons  
  126. if the covalent bond is between identical atoms the sharing of electrons is 
    equal e.g. H, Cl2
  127. if the covalent bond is between two different atoms then the sharing of the electrons is
    • unequal e.g. H-F 
  128. in H-F why is the fluorine atom much better at attracting the pair of electrons than hydrogen 
    fluorine has more protons 
  129. electro-negativity is  
    the ability of an atom to withdraw electron density from a covalent bond 
  130. if the two atoms in a bond have different electro-negativities then  
    the more electronegative element has a greater share of electrons 
  131. where are the most electronegative elements 
    at the tops of groups 5,6,7 
  132. describe and explain the trends in electro-negativity 
    • electro-negativity increases across the periodic table as the number of protons in the nucleus increases . this increases the ability of an atom to attract electrons
    • electro-negativity decreases down the periodic table as the amount of electrons in complete energy levels increases . As a result the nucleus is shielded and has less ability to attract electrons 
    • fluorine is the most electronegative element owing to its small size  
  133. polar bonds are those in which 
    a pair of electrons is not shared equally 
  134. polar bond 
    a covalent bond between atoms with different electro-negativities 
  135. the more electronegative element has a partial charge shown by 
  136. the less electronegative element has a partialo charge shown  by 
  137. in the hydrogen chloride molecule :
    • the chlorine atom is more electronegative so has a partial negative charge 
    • the hydrogen atom therefore has a partial positive charge 
    • as a result the H-Cl molecule can be described as polar Hδ+ - Clδ-
  138. in the carbon dioxide molecule :
    • the oxygen atom is more electronegative so has a partial negative charge 
    • the carbon atom has two partial charges positive charges as it is bonded to two oxygen atom 
    • as a result , the C=O bond can be described as polar  Oδ-=C2δ+=O δ-
  139. a molecule that contains polar bonds may not  be a polar molecule . to find out if a molecule is polar 
    • draw the molecule (in 3D if necessary remembering about the influence of lone pairs)
    • label any polar bonds using the δ+ , δ- convention 
    • then examine the shape of the molecule 
    • if molecule has a positive end and a negative end then the molecule is polar but if it doesn't the molecule is non polar 
  140. define dipole 
    opposite charges separated by a short distance in a molecule or ion 
  141. the greater the electro negativity difference between two atoms 
    the bigger the dipole (difference in charge between atoms) and the more polar the bond 
  142. if there is a large difference in electro-negativity , the covalent bond can be described as  
    having ionic character 
  143. if there is a large difference in electro-negativity then the bonding is  
    ionic rather than covalent 
  144. pure ionic characteristics  
    • large difference in electro-negativity
    • electrons completely transferred
    • electrons not shared   
  145. pure covalent characteristics 
    • no difference in electro-negativity
    • electrons not transferred
    • electrons shared equally 

  146. why is the H-F bond described as having an ionic character 
    there is a larger difference in electro-negativity 
  147. ionic bonds can also show 
  148. the electron cloud around a large negative ion can be distorted by a
    • small highly charged positive ion (these have high charge density) 
    • the positive ion is said to be polarizing 
    • the negative ion is said to be polarized 
    • if this happens to a large enough extent the ionic bond takes on a covalent character 
  149. do sodium ions have a high or low charge density 
    low charge density which means they are not able to polarise other ions that easily so its compounds have an ionic character 
  150. do aluminium ions have a low or high charge density 
    a high charge density so they are able to polarise large negative ions so many of its compounds have a covalent character 
  151. an intermolecular force is a
    weak attractive force between molecules . 
  152. the strength of intermolecular forces determines 
    the melting points and boiling points of substances and can influence some of their other properties to 
  153. intermolecular forces hold 
    substances together in solid or liquid form .
  154. a permanent dipole-dipole attraction is an 
    attractive force that exists between polar molecules 
  155. polar molecule 
    a molecule in which the charge isn't symmetrically distributed so that one area is slightly positively charged and another negatively charged 
  156. molecules with a permanent dipole have regions of 
    different electron density within them . these molecules are described as polar . so permanent dipole dipole forces are intermolecular forces (forces between molecules) in polar molecules . they only exist between any two molecules that have permanent dipoles.  
  157. for the hydrogen chloride molecule 
    • the electronegative chlorine of one HCl molecule will attract the electro positive hydrogen of another HCl molecule 
    • the doted line shows the permanent dipole dipole force 
  158. there are two other types of intermolecular forces (forces between molecules) which are 
    • hydrogen bonding 
    • van der waals 
  159. name the intermolecular forces and order them from strongest to weakest 
    • hydrogen bonding 
    • permanent dipole - dipole forces
    • van der waals forces 
  160. hydrogen bonds are an especially 
    especially strong permanent dipole-dipole force which exists between very electronegative elements 
  161. a hydrogen bond is an 
    • intermolecular force between a lone pair of electrons of an N, O , or F atom in one molecule , and a H atom joined to an N , O or F atom in another molecule 
  162. O , N and F are the most electronegative atoms so as a result
    the dipole-dipole force of hydrogen bonding between the molecules is especially strong 
  163. the most common examples of hydrogen bonding asked about in AS exams are 
    • water 
    • ammonia
    • hydrogen fluoride 
  164. when drawing hydrogen bonding you must include 
    • labelled dipoles on each molecule 
    • lone pairs on the , O , N or F molecule 
    • a dotted line to represent the hydrogen bond 
  165. ammonia hydrogen bonding
  166. water hydrogen bonding
    • with a lone pair of e- on every O
  167. hydrogen fluoride hydrogen bonding 
    • with a lone pair of e- on every F 
  168. what type of structure does ice have 
    ice has a very regular structure held together by hydrogen bonding between molecules 
  169. hydrogen bonding causes ice to have some unusual properties 
    • ice floats on water . It is the only substance for which this is the case 
    • water has a large surface tension , caused by a network of hydrogen bonds on the surface 
    • water has a higher boiling point than would be expected . this is because hydrogen bonds must be broken when water is boiled 
  170. a van der waals force is a
    force of attraction between a temporary dipole one one molecule and an induced dipole on another molecule 
  171. van der waals forces form between 
    non polar and polar molecules but are the dominant force between non-polar molecules 
  172. what is a temporary dipole 
    the asymmetrical distribution of the electron pair in a covalent bond  
  173. electrons in a molecule are 
    constantly moving 
  174. the fact that electrons in a molecule are constantly moving means that 
    • the electron cloud around an atom or within a non polar molecule is not static 
    • at any instant in time the distribution of the electrons may be uneven although on average they are distributed evenly . 
    • as a result a non polar molecule may have a temporary dipole 
  175. the presence of a temporary dipole in one atom or molecule can cause 
    a dipole to form in a nearby atom or molecule . this dipole is called an induced dipole . The induced dipole can then induce a dipole in a neighbouring atom or molecule . The net effect of this is a force of attraction between the particles called a temporary dipole-induced dipole force 
  176. temporary dipole forces and induced dipole forces and temporary dipole -induced dipole forces are examples of 
    van der waals forces 
  177. all ................................... have van der waals forces between the particles when they are in the liquid or solid state . For example iodine is a group seven element that exists as ................................................................................................................................iodine is a ....... at room temperature . so the molecules of iodine are in a ............ with van der waals forces ............... these forces are relatively ................................................................................... . so as a ..............................................................................................................................
    • non polar atoms or molecules 
    • diatomic molecules (two atoms joined together with a covalent bond)
    • solid crystalline 
    • regular arrangement 
    • between molecules 
    • weak so little energy is needed to overcome them 
    • iodine has a low melting and boiling point and can sublime 
  178. if a substance can sublime it means that 
    it can change directly from a solid to a gas 
  179. the strength of van der waals forces is dependant on 
    • the size of the atom or molecule 
    • the area of contact between the atoms or molecules  
  180. explain why the boiling point increase significantly down group 7 
    • the size of the halogen molecules increases down the group
    • the number of electrons in each molecule increases down the group
    • as a result temporary dipoles form more readily in the large halogen molecules
    • in addition dipoles are more readily induced in adjacent molecules
    • as a result the van der waals forces get stronger down the group .
    • as a result more energy is needed to overcome these forces  
    • the same trend occurs in the alkanes as the number of carbon atoms in the chain increases 
  181. molecules with the lowest boiling points have the most
    chain branching . this reduces the area of contact between other molecules and reduces the strength of the van der waals forces 
  182. when a solid is melted or a liquid frozen ,a liquid boiled or a gas condensed it is said to have 
    changed state 
  183. all changes of state involve 
    changes in energy 
  184. when a solid melts or a liquid boils , the energy is used to
    break the forces between the atoms molecules or ions involved 
  185. as a change of state occurs the temperature stays constant because 
    the energy provided to the system is to break the force 
  186. when melting a pure metal what attraction is broken
    the attraction between the lattice of positive charge ions and the de localised electrons . this is a strong force so requires a huge amount of energy , which means metals have a high melting point . the melting points increase as the charge on the metal ion and the number of electrons increases 
  187. when an ionic substance melts the energy provided is used to break the attraction between
    ions of opposite charge . this attraction is strong so ionic substances are hard to melt 
  188. melting simple molecular substances requires the breaking of 
    intermolecular force between the molecules . there are three types of force that can be broken . each of these is weaker than the covalent bond that exists between the atoms within the molecules 
  189. giant molecular substances have a large network of 
    covalent bonds . melting these substances involves a large amount of energy many strong covalent bonds must be broken in order to change state 

  190. explain the formation of hydrogen bonding between protein molecules 
    • both C=O and H-C are polar bonds since electro negativity N>H and O>C 
    • hydrogen bonding between H and =O or N in different molecules 
    • using a lone pair of e- on an O or N atom
  191. an example of an ionic crystal is 
    Sodium chloride
  192. what structure do ionic crystals have 
    giant lattice 
  193. give an example of a metallic crystal
    • magnesium
    • the positive metal ions represent Mg2+ ions 
  194. what structure do metallic crystals have 
    giant lattice 
  195. what structure do giant covalent crystals have 
    • giant covalent 
    • macromolecular 
  196. what structure do molecular crystals have 
    simple molecular 
  197. crystals must be 
  198. give examples of molecular crystals and draw them 
    • iodine 
    • ice 
  199. arrangement of particles in a solid , liquid , gas 
    • a solid has its particles very close together . this gives it a high density and makes it incompressible . the particles vibrate about a fixed point and can't move freely .
    • A typical liquid has similar density to a solid and is virtually incompressible . the particles move about freely and randomly within the liquid , allowing it to flow . However it's particles cannot travel too far except on the surface 
    • in gases , the particles have loads more energy and are much further apart . So the density is pretty low and and it's very compressible . The particles move about freely with not a lot of attraction between them , so they'll quickly diffuse to fill a container . 
  200. ........... ..... don't break during melting and boiling of simple covalent substances 
    covalent bonds 
  201. to melt or boil a simple covalent compound you have to overcome 
    the van der waals forces or the permanent dipole-dipole forces or the hydrogen bonding forces that hold the molecules together . you don't need to break the much stronger covalent bonds that hold atoms together in the molecules . That's why simple covalent compounds generally have relatively low melting and boiling points 
  202. explain what happens when a solid is heated 
    energy is supplied to the particles making them vibrate more about a fixed position 
  203. describe melting or fusion (how a solid is changed into a liquid)
    more energy is supplied to the solid so the forces between the particles that hold them together in a solid state are weakened , so the particles are held more loosely . The energy change is called enthalpy change of fusion . The temperature doesn't change because the heat energy is absorbed and used to weaken the forces 
  204. Describe what happens as a liquid is changed to a gas 
    enough energy must be supplied to break all of the forces between the particles . this is the enthalpy change of vaporisation . as with melting there is no energy change during boiling 
  205. explain what happens when a liquid is heated 
    energy is supplied to the particles . they gain kinetic energy and move more quickly . again , the average distance between particles increases causing the liquid to expand 
  206. when talking about intermolecular forces always say
    between molecules