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1 IB Topic 4: Bonding 4.1: Ionic bonding 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged ions. 4.1.2 Describe how ions can be formed as a result of electron transfer. 4.1.3 Deduce which ions will be formed when elements in groups 1,2 and 3 lose electrons. 4.1.4 Deduce which ions will be formed when elements in groups 5,6 and 7 gain electrons. 4.1.5 State that transition elements can form more than one ion. 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values. 4.1.7 State the formula of common polyatomic ions formed by non- metals in periods 2 and 3. 4.1.8 Describe the lattice structure of ionic compounds. Slide 2 2 IB Topic 4: Bonding Only one group of elements are stable (nonreactive). What is unique about their electron structure? Filled s & p sublevels All other elements react in order to achieve this stable electron configuration. Ionic Bond: Transfer of electrons; metal + nonmetal Covalent Bond: Sharing of electrons; nonmetal + nonmetal Slide 3 3 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged particles. Ionic bonding occurs when one or more electrons are transferred from the outer shell of another atom. When one atom gives up its electron(s), it becomes a positively charged cation. When the other atom gains the electron(s), it becomes a negatively charged anion. OPPOSITES ATTRACT The ionic bond formed between the two elements is caused by the attractive force (electrostatic attraction) between the positive and negative ions. Slide 4 4 4.1.1 Describe the ionic bond as the electrostatic attraction between oppositely charged particles. The lithium ion has a positive charge The fluorine ion (fluoride) has a negative charge The bond is formed due to the electrostatic attraction between the opposite charges Slide 5 5 4.1.2 Describe how ions can be formed as a result of electron transfer. Ions are formed when atoms gain or lose electrons Fluorine gains one electron to become the fluoride ion having a charge of -1 Lithium loses one electron to become the lithium ion having a charge of +1 Slide 6 6 4.1.2 Describe how ions can be formed as a result of electron transfer. Ions are formed when atoms gain or lose electrons Reaction between Sodium and Chlorine Sodium configuration: 1s 2 2s 2 2p 6 3s 1 or __ __ __ 1s 2s 2p 3s If Na loses one electron then it would end in 2s 2 2p 6 and be stable. Then sodium has 11 protons (11+), but only 10 electrons (10-) so it acquires a charge of 1+ and becomes the sodium ion, Na +. In order for it to lose an electron, something has to gain an electron Slide 7 7 4.1.2 Describe how ions can be formed as a result of electron transfer. Ions are formed when atoms gain or lose electrons Reaction between Sodium and Chlorine Chlorine configuration: 1s 2 2s 2 2p 6 3s 2 3p 5 or __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p If Cl gains one electron then it would end in 3s 2 3p 6 and be stable. Then chlorine has 17 protons (17 + ), and 18 electrons (18 - ) so it acquires a charge of 1 - and becomes the chloride ion, Cl -. It will gain the electron from the sodium. Slide 8 8 4.1.2 Making ionic bonds Sodium gives up an electron to chlorine. Sodium now has a full 2 nd shell, & chlorine has a full 3 rd shell. Slide 9 9 Li + F Li + F - The Ionic Bond 1s 2 2s 1 1s 2 2s 2 2p 5 1s 2 1s 2 2s 2 2p 6 [He][Ne] Li Li + + e - e - + FF - F - Li + + Li + F - Slide 10 10 4.1.2 Describe how ions can be formed as a result of electron transfer. Reaction between Sodium and Chlorine Since opposite charges attract, the Na + and Cl - ions form an ionic bond. Name: Sodium chloride Formula NaCl Slide 11 11 4.1.2 Describe how ions can be formed as a result of electron transfer. Ions are formed when atoms gain or lose electrons Reaction between Magnesium and Chlorine Magnesium configuration: 1s 2 2s 2 2p 6 3s 2 or __ __ __ 1s 2s 2p 3s If Mg loses two electrons then it would end in 2s 2 2p 6 and be stable. Then magnesium has 12 protons (12+), but only 10 electrons (10-) so it acquires a charge of 2+ and becomes the magnesium ion, Mg 2+. In order for it to lose two electrons, something has to gain two electrons Slide 12 12 4.1.2 Describe how ions can be formed as a result of electron transfer. Ions are formed when atoms gain or lose electrons Reaction between Magnesium and Chlorine Chlorine configuration: 1s 2 2s 2 2p 6 3s 2 3p 5 or __ __ __ __ __ __ __ __ __ 1s 2s 2p 3s 3p If Cl gains one electron then it would end in 3s 2 3p 6 and be stable. Then chlorine has 17 protons (17 + ), and 18 electrons (18 - ) so it acquires a charge of 1 - and becomes the chloride ion, Cl -. Since chlorine can only gain one electron and magnesium gives up two electrons, magnesium requires two chlorine atoms. Slide 13 13 4.1.2 Describe how ions can be formed as a result of electron transfer. Reaction between Magnesium and Chlorine Since opposite charges attract, the Mg 2+ and the 2 Cl - ions form an ionic bond. Name: Magnesium chloride Formula MgCl 2 Slide 14 14 4.1.2 Describe how ions can be formed as a result of electron transfer. Magnesium gives away 2 electrons; one to each chlorine atom Now Magnesium has a full 2 nd shell and both Chlorines have a full 3 rd shell Slide 15 15 4.1.2 Describe how ions can be formed as a result of electron transfer. Reaction between Potassium and Oxygen Potassium configuration: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 Potassium will lose 1 electron and become the potassium ion K +. Oxygen configuration: 1s 2 2s 2 2p 4 Oxygen will gain 2 electrons and become the oxide ion O 2-. Two potassiums are needed to combine with one oxygen. Formula:K 2 O Name:Potassium oxide Slide 16 16 Reaction between Aluminum and Bromine Diagram the bonding between Al and Br, write the formula, and give the name. 4.1.2 Describe how ions can be formed as a result of electron transfer. Slide 17 17 Reaction between Aluminum and Bromine Diagram the bonding between Al and Br, write the formula, and give the name. Formula: AlBr 3 Name: Aluminum bromide 4.1.2 Describe how ions can be formed as a result of electron transfer. Slide 18 18 4.1.3 Deduce which ions will be formed when elements in groups 1, 2 and 3 lose electrons. 4.1.4 Deduce which ions will be formed when elements in groups 5, 6 and 7 gain electrons Positive Ions (cations) Group 1: Lose 1 valence electron Charge of +1: Li +, Na +, K + Group 2 Lose 2 valence electrons Charge of +2: Mg 2+, Ca 2+ Group 3 Lose 3 valence electrons Charge of +3: Al 3+ Negative Ions (anions) Group 5: Gain 3 electrons Charge of -3: N 3-, P 3- Group 6: Gain 2 electrons Charge of -2: O 2-, S 2- Group 7 Gain 1 electron Charge of -1: F -, Cl -, Br -, I - Slide 19 19 4.1.5 State that transition elements can form more than one ion. In most transition elements, d electrons can become involved in the reaction Iron can lose 2 electrons (the 2 in the 4s) (Fe 2+ ) or 3 electrons (the 2 in the 4s and 1 in the 3d) (Fe 3+ ) The name of the Fe 2+ ion is iron(II) or ferrous The name of the Fe 3+ ion is iron(III) or ferric Chromium can lose 2 electrons (the 2 in the 4s) (Cr 2+ ) or 3 electrons (the 2 in the 4s and 1 in the 3d) (Cr 3+ ) The name of the Cr 2+ ion is chromium(II) or chromous The name of the Cr 3+ ion is chromium(III) or chromic Slide 20 20 4.1.5 State that transition elements can form more than one ion. MOST transition metals REQUIRE Roman Numerals because they can have more than one possible charge. BUT t here are ALWAYS exceptions to rules (in chemistry) These transition metals only form ONE ion: Ag +1, Zn +2 and Cd +2 Label these on your periodic table!! Slide 21 21 Metal + Nonmetal Ionic compounds are generally composed of a metal combined with a nonmetal(s) Electronegativity Difference Ionic bonds have electronegativity differences greater than or equal to 1.8 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values. Slide 22 22 Practice 4.1.6 Predict whether a compound of two elements would be ionic from the position of the elements in the periodic table or from their electronegativity values. Slide 23 23 Polyatomic ions are ions consisting of two or more atoms bonded together NitrateNO 3 - HydroxideOH - SulfateSO 4 2- CarbonateCO 3 2- PhosphatePO 4 3- AmmoniumNH 4 + Bicarbonate (hydrogen carbonate) HCO 3 - More in chart of ions: nitrite, sulfite, perchlorate, chlorate, chlorite, hypochlorite, silicate 4.1.7 State the formula of common polyatomic ions formed by non-metals in periods 2 and 3. Slide 24 24 4.1.8 Describe the lattice structure of ionic compounds. Ionic compounds form crystalline solids that have a particular structure (lattice) of alternating positive and negative ions Sodium Chloride: Each sodium ion is surrounded by up to 6 chloride ions and each chloride ion is surrounded by up to 6 sodium ions. Slide 25 25 IB Topic 4: Bonding 4.2: Covalent bonding Part 1 4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei. 4.2.2 Describe how the covalent bond is formed as a result of electron sharing. 4.2.3 Deduce the Lewis (electron dot) structures for up to four electron pairs on each atom. 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. 4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or from their electronegativity values. 4.2.6 Predict the relative polarity of bonds from electronegativity values. Slide 26 26 4.2.1 Describe the covalent bond as the electrostatic attraction between a pair of electrons and positively charged nuclei Covalent bonds occur between nonmetals since both want to gain electrons. They share electrons to achieve a stable configuration. The electrons in the bond are electrostatically attracted by both nuclei, so that it forms a directional bond between the two atoms. Slide 27 27 4.2.2 Describe how the covalent bond is formed as a result of electron sharing. Covalent bonds occur between nonmetals since both want to gain electrons. They share electrons to achieve a stable configuration Electron Dot Diagrams Need 1 electron:HFClBrI Need 2 electrons:OSSeTe Need 3 electrons:NPAs Need 4 electrons:CSi Slide 28 28 4.2.2 A covalent bond is a chemical bond in which two or more electrons are shared by two atoms. Why should two atoms share electrons? FF + 7e - FF 8e - F F F F Lewis structure of F 2 lone pairs single covalent bond Slide 29 29 4.2.3 Deduce the Lewis (electron dot) structures for up to four electron pairs on each atom. Double covalent bonds Sharing 2 pair of electrons O 2 H 2 CO CO 2 Single covalent bonds Sharing 1 pair of electrons H 2 HCl CCl 4 Slide 30 30 4.2.3 Deduce the Lewis (electron dot) structures for up to four electron pairs on each atom. Triple covalent bonds Sharing 3 pair of electrons N 2 WS 7,9 Dative (Coordinate) Covalent Bonds Both electrons in the bond come from the same atom CO NH 4 + H 3 O + Slide 31 31 8e - H H O ++ O HH O HHor 2e - 4.2.3 Lewis structure of water Double bond two atoms share two pairs of electrons single covalent bonds O C O or O C O double bonds Triple bond two atoms share three pairs of electrons N N N N triple bond or Slide 32 32 Modeling Covalent Bonds Molecular model sets: modeling covalent bonding. Create the 8 molecules listed on the next slide and complete a data table similar to the following: Formula (given) Line DiagramDot Diagram You will be making 8 molecules!! Slide 33 33 Modeling Covalent Bonds 1 bond = 2 shared electrons Red = oxygen Black = carbon White (& green) = hydrogen Nitrogen is a carbon atom with 1 plug (only 3 bonding sites) H 2 O H 2 C 2 H 4 O 2 CO 2 C 6 H 12 NH 3 N 2 Slide 34 34 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. The more pairs of Electrons that are shared between two atoms (bonds) in a molecule will make the attraction between the atoms stronger and shorter Length nm Strength (kj mol-1) C-O0.143356 C-C0.154348 C=O0.121736 C=C0.134657 C 0.120908 Slide 35 35 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. Length nm Strength (kj mol-1) C-O0.143356 C-C0.154348 C=O0.121736 C=C0.134657 C 0.120908 Bond Strength Triple bond > Double Bond > Single Bond Bond Lengths Triple bond < Double Bond < Single Bond So, if you compare a single bond to a triple bond, Which one is stronger? Which one is shorter? Slide 36 36 4.2.4 State and explain the relationship between the number of bonds, bond length and bond strength. Lets practice! Draw C 3 H 6 (or CH 3 CHCH 2 ) Identify the longest carbon-to-carbon bond Identify the strongest carbon-to-carbon bond Slide 37 37 4.2.5 Predict whether a compound of two elements would be covalent from the position of the elements in the periodic table or from their electronegativity values. Nonmetal + Nonmetal Covalent compounds generally occur between elements that are closer to each other on the periodic table (with the exception of hydrogen) Electronegativity Difference Covalent bonds have electronegativity differences from 0 to 1.8. A bond with electronegativity difference of 1.8 is classified as covalent, NOT ionic. Slide 38 38 Slide 39 39 4.2.6 Predict the relative polarity of bonds from electronegativity values. Covalent bonds can either be polar or nonpolar Polar Covalent Bond: Polar bonds form when electrons are shared unequally between two atoms due to a difference in electronegativity One atom will have a partial positive charge and the other atom will have a partial negative charge The greater the difference in electronegativity, the more polar the bond will be Slide 40 40 H F F H Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms electron rich region electron poor region e - riche - poor ++ -- 2.24.0 Electronegativity difference is 1.8, therefore it is polar covalent Slide 41 41 4.2.6 Predict the relative polarity of bonds from electronegativity values. Nonpolar Covalent Bond: Nonpolar bonds form when electrons are shared equally between two atoms with the same electronegativity values 2.22.2 2.2 = 0 Slide 42 42 4.2.6 Predict the relative polarity of bonds from electronegativity values. Electronegativity Difference Covalent bonds have electronegativity differences from 0 to 1.8. covalent bonds can either be polar or nonpolar: Nonpolar covalent bonds have electronegativity difference of 0 Polar covalent bonds have electronegativity differences above 0, and equal to or less than 1.8 http://animatedchemistry.org/?p=528 Slide 43 43 4.2.6 Predict the relative polarity of bonds from electronegativity values. Nonpolar vs Polar Covalent Bonds: Slide 44 44 Covalent share e - Polar Covalent partial transfer of e - Ionic transfer e - Increasing difference in electronegativity 4.2.6 Classification of bonds by difference in electronegativity DifferenceBond Type 0 Covalent > 1.8 Ionic Nonpolar Covalent 0 and 1.8 Polar Covalent0 < and 1.8 01.8 Slide 45 45 Slide 46 46 Naming & Formulas We need to learn the rules for naming ionic vs. covalent compounds Slide 47 47 IB Topic 4: Bonding 4.2: Covalent bonding Part 2 4.2.7 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell elctron pair repulsion theory (VSPER). 4.2.8 Predict whether or not a molecule is polar from its molecular shape and bond polarities. 4.2.9 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). 4.2.10 Describe the structure of and bonding in silicon and silicon dioxide. Slide 48 48 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSPER). Shapes and bond angles of simple molecules and ions are determined by the number of electron centers around the central atom. Valence Shell Electron Pair Repulsion (VSEPR): pairs of electrons are arranged so that they are far away from each other as possible 1)Multiple bonds count as if they are just one pair of electrons (negative charge centers) 2)Non-bonding pairs exert a greater repulsion than bonding pairs Slide 49 49 Predict the shape and bond angles for species with four, three and two negative charge centres on the central atom using the valence shell electron pair repulsion theory (VSPER). 1) Four negative charge centers Four bonding pairs (CH 4, NH 4 + ); Tetrahedral shape; 109.5 o angle Three bonding pairs, one non-bonding pair (NH 3, H 3 O+) Non- bonding pair exerts greater repulsion; Trigonal pyramid shape; 107 o angle Two bonding pairs, two non-bonding pairs (H 2 O); Bent shape; 105 o angle 2) Three negative charge centers (BF 3, C 2 H 4, SO 3 ) Trigonal planar shape, 120 o angle 3) Two negative charge centers (C 2 H 2, CO 2 ) Linear shape, 180 o angle Predict the shape and bond angles for H 2 S, H 2 CO, PH 3, POCl 3 For each of molecules on the worksheet, identify the shape. Slide 50 50 Exceptions to the Octet Rule Satisfied with less than 8 electrons: BF 3 Boron is happy with 6 v.e-s KNOW THIS ONE!!! BCl 3 BeCl 2 Beryllium is happy with 4 v.e-s Dont forget hydrogen only wants 2! Satisfied with more than 8 electrons: PF 5 Phosphorus is fine with 10 v.e-s SF 6 Sulfur is fine with 12 v.e-s Slide 51 51 Valence shell electron pair repulsion (VSEPR) model: Predict the geometry of the molecule from the electrostatic repulsions between the electron (bonding and nonbonding) pairs. AB 2 20 Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry linear B B Slide 52 52 Cl Be 2 atoms bonded to central atom 0 lone pairs on central atom Slide 53 53 AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar Slide 54 54 Slide 55 55 AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 4 40 tetrahedral Slide 56 56 Slide 57 57 AB 2 20linear Class # of atoms bonded to central atom # lone pairs on central atom Arrangement of electron pairs Molecular Geometry VSEPR AB 3 30 trigonal planar AB 4 40 tetrahedral AB 5 50 trigonal bipyramidal trigonal bipyramidal Slide 58 58 Slide 59 59 bonding-pair vs. bonding pair repulsion lone-pair vs. lone pair repulsion lone-pair vs. bonding pair repulsion >> Slide 60 60 Slide 61 61 Predicting Molecular Geometry 1.Draw Lewis structure for molecule. 2.Count number of lone pairs on the central atom and number of atoms bonded to the central atom. 3.Use VSEPR to predict the geometry of the molecule. What are the molecular geometries of SO 2 and SF 4 ? SO O AB 2 E bent S F F F F AB 4 E distorted tetrahedron Slide 62 62 Predict whether or not a molecule is polar from its molecular shape and bond polarities. Polarity is a physical property that impacts other physical properties, such as solubility, melting points and boiling points. It has to do with the electronegativity of elements within a bond or molecule. Slide 63 63 Predict whether or not a molecule is polar from its molecular shape and bond polarities. Polar molecules are determined by bond polarity and the shape of the molecule. HCl Cl 2 H 2 O CH 4 For each of molecules on the worksheet, identify if the molecule is polar. Slide 64 64 Predict whether or not a molecule is polar from its molecular shape and bond polarities. Polar BONDS are different from polar MOLECULES! To determine whether a BOND is polar Only pay attention to the two atoms in the bond To determine whether a MOLECULE is polar Pay attention to ALL of the atoms in the compound! Slide 65 65 IB Topic 4: Bonding 4.3: Intermolecular Forces Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. Describe and explain how intermolecular forces affect the boiling points of substances. Slide 66 66 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. Hydrogen Bonds Permanent Dipoles When hydrogen is attached to a highly electronegative element (F,O,N) it forms a very strong polar molecule Hydrogen becomes positive and can weakly bond to the unshared electrons of another molecule Hydrogen bonds are stronger than other dipole-dipoles and give water its unusual properties such as liquid state and density H2OH2OH2SH2SH 2 Se Molar Mass 183481 Boiling Point ( o C) 100-61-41 HFHClHBr Molar Mass 2036.581 Boiling Point ( o C) 10-70-65 Slide 67 67 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. Slide 68 68 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. Dipole-Dipole Interactions Permanent Dipoles Polar Molecules Permanent electrostatic forces of attraction between the molecules. Stronger attractions called dipole moments caused by opposite charged regions of the dipoles. Stronger than dispersion forces resulting in higher melting/boiling points. Propane C 3 H 8 Ethanal C 2 H 4 O Molar Mass 44 Boiling Point ( o C) -42.220.8 BondingNonpolarPolar CH 3 ClCH 2 Cl 2 CHCl 3 Molar Mass 50.584105.5 Boiling Point ( o C) -23.740.261.2 Slide 69 69 Types of Intermolecular Forces Dipole-Dipole Forces Attractive forces between polar molecules Orientation of Polar Molecules in a Solid 11.2 Slide 70 70 Describe the types of intermolecular forces (attractions between molecules that have temporary dipoles, permanent dipoles or hydrogen bonding) and explain how they arise from the structural features of molecules. van der Waals forces Temporary Dipoles Non-polar Molecules No permanent electrostatic forces of attraction between the molecules. Very weak attractions called dispersion forces (van der Waals forces) caused by motions of electrons. The strength generally increases as the number of electrons and the molar mass increases. Halogen melting & boiling points increase down the group. Halogen F2F2 Cl 2 Br 2 I2I2 # electrons 183470110 Molar Mass 3871160254 Boiling Point (oC) -188-54.758.884 State Gas LiquidSolid Slide 71 71 Describe and explain how intermolecular forces affect the boiling points of substances. Boiling points dependent on the type of intermolecular bonding Non-polar: weak interactions holding the molecules together so easily broken Polar: stronger interactions so higher boiling points than comparably sized non-polar compounds Hydrogen bonds: strongest covalent interactions so higher boiling points than comparable polar compounds. Read pg. 463-464: 23,24 Propane C 3 H 8 Dimethyl ether C 2 H 6 O Ethanal C 2 H 4 O Ethanol C 2 H 6 O Molar Mass 44464446 Boiling Point ( o C) -42.2-232178.4 BondingNon- polar Slightly polar PolarHydrogen Slide 72 72 IB Topic 4: Bonding 4.4: Metallic Bonding Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons Explain the electrical conductivity and malleability of metals Slide 73 73 Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalized electrons When metal atoms combine, one or more of their valence electrons become detached from each atom to become delocalized. These valence electrons are free to move throughout the metallic structure The bonding in metals consist of the attraction between the delocalized valence electrons and the cations. Slide 74 74 Explain the electrical conductivity and malleability of metals Metals are ductilethey can be drawn into wires A force can change the shape of a metal Cations can slide past each other to rearrange the shape of the solid. Because the electrons are delocalized, there is no significant change in the bonding forces, so they are malleable and ductile Slide 75 75 Explain the electrical conductivity and malleability of metals Electricity is the directional flow of electrons between atoms Electrons need a medium to flow through, such as Materials with free-floating electrons Metals or graphite Ions Molten ionic substances, like NaCl (l) Ionic substances dissolved in water, like NaCl (aq) Slide 76 76 Explain the electrical conductivity and malleability of metals Valence electrons are free to move throughout the structure so metals are excellent conductors of heat and electricity. Metallic melting points determined by the number of valence electrons, size of the cation, and how the cations are packed together. In general, melting point decreases down a group. Alkali Metal Ionic Radius (nm) Melting Point ( o C) Lithium0.060180 Sodium0.09597.8 Potassium0.13363.7 Rubidium0.14838.9 Cesium0.16928.7 Slide 77 77 IB Topic 4: Bonding 4.5: Physical Properties Compare and explain the properties of substances resulting from different types of bonding Slide 78 78 Melting points (mp) - solid to liquid Boiling points (bp) - liquid to gas Volatility - how easily it is converted to gas Conductivity (conducts electricity) depends on whether the substance contains electrically charged particles that are free to move through it Solubility - solutes ability to dissolve in solvent -Effect of impurities on the melting point of a substance impurities disrupt the regular lattice that its particles adopt in the solid state, weakening the bonding and lowering the boiling pt. -Compare/explain solubilities (solutes ability to dissolve in solvent) of compounds in non-polar and polar solvents solute will dissolve if forces between the solutes particles are as strong or stronger than the forces in the solvent Compare and explain the properties of substances resulting from different types of bonding Slide 79 79 Compare and explain the properties of substances resulting from different types of bonding Ionic bonding (Sodium Chloride, Calcium Carbonate) Held tighter by strong electrostatic forces in 3D Non-volatile, high mp, high bp Solid at Rm temp. Hard and brittle b/c of lattice structure If it is moved slightly + and ions would be lined up and repel each other. Non-mobile e- as solid, but move freely as liquid or in solution so solids dont conduct electricity, aqueous ionic substances do Insoluble in most solvents (some are soluble in the very polar H2O molecule) b/c Slide 80 80 Compare and explain the properties of substances resulting from different types of bonding Giant Covalent Carbon (diamond) Silicon dioxide (sand) Atoms joined by strong covalent bonds, Insoluble in all solvents High mp and bp, e- firmly held in place therefore no conductivity Slide 81 81 Carbon exhibits allotropy-it can exist in more than one physical form. Diamond: Each carbon bonded equally to four other carbon atoms to form a giant tetrahedral structure. Strong covalent bonds so high melting point, hard, electrons involved in the bond so non-conductor Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). Slide 82 82 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). Carbon exhibits allotropy-it can exist in more than one physical form. Graphite: Each carbon bonded strongly to three other carbons in a trigonal planar structure forming layers of hexagonal rings so high melting point. However forces of attraction between layers are weak because they are formed by delocalized electrons that can move between layers so good conductor and the layers can slide over each other so good lubricant. Slide 83 83 Carbon exhibits allotropy-it can exist in more than one physical form. Fullerene (C 60 ): Carbon atoms joined in a combination of hexagonal and pentagonal rings to form a sphere. Basic unit of nanotubes It looks like a soccer ball! Theyre called Buckyballs Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). Slide 84 84 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). Slide 85 85 Carbon Nanotubes http://www.youtube.com/watch?v=zQAK4x xPGfMhttp://www.youtube.com/watch?v=zQAK4x xPGfM Nova: http://www.youtube.com/watch?v=19nzPt6 2UPg&NR=1&feature=endscreen http://www.youtube.com/watch?v=19nzPt6 2UPg&NR=1&feature=endscreen Slide 86 86 Describe and compare the structure and bonding in the three allotropes of carbon (diamond, graphite and C 60 fullerene). http://www.edinformatics.com/interactive_ molecules/graphite.htm Slide 87 87 Describe the structure of and bonding in silicon and silicon dioxide. Silicon displays the very similar bonding to diamond, only it is a bit weaker due to the larger size of silicon atoms. Silicon dioxide is also similar to the diamond structure, having a tetrahedral shape. Common forms of SiO 2 are quartz and sand (which is impure SiO 2, colored yellow by iron (III) oxide). Slide 88 88 Compare and explain the properties of substances resulting from different types of bonding Molecular covalent Carbon dioxide, ethanol, iodine Strong intramolecular forces, weak intermolecular forces, usually liquids or gases at Rm temp, soft solid, Often dissolve in nonpolar solvents but insoluble in very polar solvent b/c of the stronger intermolecular force in H2O H-bonding effects properties by increasing bp and mp and becoming soluble in H2O Consider also the solubilities of alcohols in water as the length of the carbon chain increases. As the hydrocarbon chain increases it interferes with the hydrogen bonding, therefore becoming less soluble Slide 89 89 Solubility of salt in water: http://www.youtube.com/watch?v=xdedxfh cpWo&feature=related http://www.youtube.com/watch?v=xdedxfh cpWo&feature=related Slide 90 90 Compare and explain the properties of substances resulting from different types of bonding If the type of bonding is known, physical properties can be predicted Type of Bonding Melting Point Boiling PointVolatilityElectrical ConductivitySolubility in Non-polar Solvent Solubility in Polar Solvent Non-polarLow HighNoYesNo PolarNo Yes Hydrogen bonding No Yes Ionic BondingYes (molten or aqueous) NoYes (most) Metallic Bonding YesNo CovalentNo Giant CovalentHigh LowNo (except graphite)No