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![Ionic Bonding
• Always formed between metal cations
and non-metals anions
• The oppositely charged ions stick like
magnets
[METALS ]+
[NON-METALS ]-
Lost e-
Gained e-](https://image.slidesharecdn.com/chapter8covalentbondingwvideos2017-240811073833-fa95a0c8/75/Chapter-8-Covalent-Bonding-w-videos-2017-pptx-5-2048.jpg)
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Chapter 8 Covalent Bonding w videos 2017.pptx
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- 2. Chemical Bond A QuickReview…. • A bond results from the attraction of nuclei for electrons – All atoms are trying to achieve a stable octet • IN OTHER WORDS – the protons (+) in one nucleus are attracted to the electrons (-) of another atom • This is Electronegativity !! Lewis Diagrams Made Easy (7min)
- 4. Three Major Typesof Bonding • Ionic Bonding –forms ionic compounds –transfer of valence e- • Metallic Bonding • Covalent Bonding –forms molecules –sharing of valence e- – This is our focus this chapter
- 5. Ionic Bonding • Alwaysformed between metal cations and non-metals anions • The oppositely charged ions stick like magnets [METALS ]+ [NON-METALS ]- Lost e- Gained e-
- 6. Metallic Bonding • Alwaysformed between 2 metals (pure metals) –Solid gold, silver, lead, etc…
- 7. Covalent Bonding • Pairsof e- are shared between 2 non- metal atoms to acquire the electron configuration of a noble gas. molecules
- 8. Covalent Bonding • Occursbetween nonmetal atoms which need to gain electrons to get a stable octet of electrons or a filled outer shell. n o n m e t a l s
- 9. How to DrawLewi s Structures: Five E asy Steps Drawing molecules (covalent) using Lewis Dot Structures • Symbol represents the KERNEL of the atom (nucleus and inner electrons) • dots represent valence electrons • The ones place of the group number indicates the number of valence electrons on an atom. • Draw a valence electron on each side (top, right, bottom, left) before pairing them.
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- 11. Always remember atomsare trying to complete their valence shell! “2 will do but 8 is great!” The number of electrons the atoms needs is the total number of bonds they can make. Ex. … H? O? F? N? Cl? C? one two one three one four
- 12. Draw Lewis DotStructures You may represent valence electrons from different atoms with the following symbols x, , H or H or H x
- 13. Covalent bonding • Theatoms form a covalent bond by sharing their valence electrons to get a stable octet of electrons.(filled valence shell of 8 electrons) • There are two electrons per bond, each atom donates one electron to the bond. • Electron-Dot Diagrams of the atoms are combined to show the covalent bonds • Covalently bonded atoms form MOLECULES
- 14. Methane CH4 • Thisis the finished Lewis dot structure • Every atom has a filled valence shell How did we get here? OR
- 15. General Rules forDrawing Lewis Structures • All valence electrons of the atoms in Lewis structures must be shown. • Generally each atom needs eight electrons in its valence shell (except Hydrogen needs only two electrons and Boron needs only 6). • Multiple bonds (double and triple bonds) can be formed by C, N, O, P, and S. • Central atoms have the most unpaired electrons. • Terminal atoms have the fewest unpaired electrons.
- 16. • When carbonis one of you atoms, it will always be in the center • Sometimes you only have two atoms, so there is no central atom Cl2 HBr H2 O2 N2 HCl • We will use a method called ANS (Available, Needed, Shared) to help us draw our Lewis dot structures for molecules
- 17. EXAMPLE 1: Writethe Lewis structure for H2O where oxygen is the central atom. Step 1: Determine the total number of electrons available for bonding. Because only valence electrons are involved in bonding we need to determine the total number of valence electrons. AVAILABLE valence electrons: Electrons available 2 H Group 1 2(1) = 2 O Group 6 6 8 There are 8 electrons available for bonding. Step 2: Determine the number of electrons needed by each atom to fill its valence shell. NEEDED valence electrons Electrons needed 2 H each H needs 2 2(2) = 4 O needs 8 8 12 There are 12 electrons needed.
- 18. Step 3: Moreelectrons are needed then there are available. Atoms therefore make bonds by sharing electrons. Two electrons are shared per bond. SHARED (two electrons per bond) # of bonds = (# of electrons needed – # of electrons available) = (N-A) = (12 – 8) = 2 bonds. 2 2 2 Draw Oxygen as the central atom. Draw the Hydrogen atoms on either side of the oxygen atom. Draw the 2 bonds that can be formed to connect the atoms. OR Step 4: Use remaining available electrons to fill valence shells for each atom. All atoms need 8 electrons to fill their valence shell (except hydrogen needs only 2 electrons to fill its valence shell, and boron only needs 6). For H2O there are 2 bonds, and 2 electrons per bond. # available electrons remaining = # electrons available – # electrons shared = A-S = 8 – 2(2) = 4 extra e- s
- 19. Sometimes multiple bondsmust be formed to get the numbers of electrons to work out • DOUBLE bond – atoms that share two e- pairs (4 e-) O O • TRIPLE bond – atoms that share three e- pairs (6 e-) N N
- 21. Step 3: SHARED(two electrons per bond) # of bonds = (N – A) = (20 – 12) = 4 bonds. 2 2 Draw carbon as the central atom (Hint: carbon is always the center when it is present!). Draw the Hydrogen atoms and oxygen atom around the carbon atom. Draw 2 bonds of the 4 bonds that can be formed to connect the H atoms. Draw the remaining 2 bonds to connect the O atom (oxygen can form double bonds) Step 4: Use remaining available electrons to fill valence shell for each atom. # electrons remaining = Available – Shared = A – S = 12 – 4(2) = 4 extra e- s
- 22. Let’s Practice H2 Available =2 valence e- Needed = H:2, H:2 = 4 e- Shared = 4 - 2= 2 e- shared ÷ 2 = 1 bond DRAW
- 23. Let’s Practice CH4 Available =8 Needed = C: 8 H:2 , H:2, H:2, H:2 = 16 Shared = (N-A)16 – 8 = 8 e- shared ÷2 = 4 bonds DRAW Carbon is the central atom because it has the most unpaired (single) valence electrons
- 24. “ANS” Method (available ,needed, shared) • Used to determine the number of bonds in a covalent molecule • “A” is the total number of valence electrons from all atoms in the chemical formula • “N” is the total number of valence electrons needed for an octet for all of the atoms in the chemical formula • “S” is the number of valence electrons that must be shared (N > A). • S ÷2 = number of bonds in molecule
- 25. Let’s Practice NH3 A =total of 8 valence e- N = N:8 H:2, H:2, H:2 = 14 valence e- S = 14-8 = 6 shared e- ÷2 = 3 bonds DRAW
- 26. Let’s Practice CO2 Available =C:4, O:6, O:6, = 16 Needed = C:8 O:8, O:8 total = 24 Shared = 24-16 = 8 valence e- must be shared = 4 bonds- put 2 valence e- in each bond. Remaining = (A-S) 16 – 8 = 8 not bonding DRAW – carbon is the central atom
- 27. Let’s Practice BCl3 borononly needs 6 valence electrons, it is an exception. A = = 24 N = B-6, Cl-8, Cl-8 , Cl-8 30 S = 30-24 = 6 ÷ 2 = 3 bonds Remaining = 24 – 6 = 18 e- not bonding DRAW Lewis Diagrams Made Easy (7min)
- 28. Naming Molecular Compounds (Covalent) TypeIII Nonmetal + nonmetal
- 29. The Covalent Bond Sharingof electrons
- 30. Properties of Molecularor Covalent Compounds • Made from 2 or more non metals • Consist of molecules not ions
- 31. Molecular Formulas Show thekinds and numbers of atoms present in a molecule of a compound. Molecular Formula = H2O
- 32. H N H H NH3 Structuralformula Molecular formula
- 33. Molecular Formulas • Examples •CO2 • SO3 • N2O5
- 34. Rules for Naming Molecularcompounds • The most “metallic” nonmetal element is written first (the one that is furthest left) • The most non metallic of the two nonmetals is written last in the formula • NO2 not O2N • All binary molecular compounds end in - ide
- 35. • Ionic compoundsuse charges to determine the chemical formula • The molecular compound‘s name tells you the number of each element in the chemical formula. • Uses prefixes to tell you the quantity of each element. • You need to memorize the prefixes ! Molecular compounds
- 36. Prefixes • 1 mono •2 di • 3 tri • 4 tetra • 5 penta • 6 hexa • 7 hepta • 8 octa • 9 nona • 10 deca Memorize!
- 37. • If thereis only one of the first element do not put (prefix) mono • Example: carbon monoxide (not monocarbon monoxide) • If the nonmetal starts with a vowel, drop the vowel ending from all prefixes except di and tri • monoxide not monooxide • tetroxide not tetraoxide More Molecular Compound Rules
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- 43. N2O5 dinitrogen Molecular compound NamingPractice pentaoxide
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- 57. Write molecular formulas forthese • diphosphorus pentoxide • P2O5 • trisulfur hexaflouride • S3F6 • nitrogen triiodide • NI3
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- 60. Bond Types 3 PossibleBond Types: • Ionic • Non-Polar Covalent • Polar Covalent
- 61. Use Electronegativity Valuesto Determine Bond Types • Ionic bonds – Electronegativity (EN) difference > 2.0 • Polar Covalent bonds – EN difference is between .21 and 1.99 • Non-Polar Covalent bonds – EN difference is < .20 – Electrons shared evenly in the bond
- 62. “Ionic Character” “Ionic Character”refers to a bond’s polarity –In a polar covalent bond, • the closer the EN difference is to 2.0, the more POLAR its character • The closer the EN difference is to .20, the more NON-POLAR its character
- 63. Place these moleculesin order of increasing bond polarity using the electronegativity values on your periodic table • HCl • CH4 • CO2 • NH3 • N2 • HF 1 EN difference = 0 2 EN difference = 0.3 4 EN difference = 0.9 4 EN difference = .9 3 EN difference = 0.8 5 EN difference = 1.9
- 64. Polar vs. Nonpolar MOLECULES •Sometimes the bonds within a molecule are polar and yet the molecule itself is non-polar
- 65. Nonpolar Molecules • Moleculeis Equal on all sides –Symmetrical shape of molecule (atoms surrounding central atom are the same on all sides) H H H H C Draw Lewis dot first and see if equal on all sides
- 66. Polar Molecules • Moleculeis Not Equal on all sides –Not a symmetrical shape of molecule (atoms surrounding central atom are not the same on all sides) Cl H H H C
- 67. H Cl - + PolarMolecule Unequal Sharing of Electrons
- 68. Cl Non-Polar Molecule Cl Equal Sharingof Electrons
- 69. H Cl Polar Molecule Notsymmetrical H B
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- 71. H H O Water is aPOLAR molecule ANY time there are unshared pairs of electrons on the central atom, the molecule is POLAR
- 72. Making sense ofthe polar non-polar thing BONDS Non-polar Polar EN difference EN difference 0 - .2 .21 – 1.99 MOLECULES Non-polar Polar Symmetrical Asymmetrical OR Unshared e- s on Central Atom
- 73. 5 Shapes ofMolecules you must know! (memorize)
- 74. Copy this slide •VSEPR – Valence Shell Electron Pair Repulsion Theory –Covalent molecules assume geometry that minimizes repulsion among valence electrons. –Shape of a molecule can be predicted from its Lewis Structure which shows all valence electrons
- 75. 1. Linear (straightline) Ball and stick model Molecule geometry X A X Unshared Pairs = 0 OR A X OR
- 76. 2. Bent Ball andstick model Lewis Diagram A X X .. Unshared Pairs = 1 or 2
- 77. 3. Trigonal Planar(flat) Ball and stick model Molecule geometry X A X X Unshared Pairs = 0
- 78. 4.Trigonal Pyramidal Ball andstick model Molecule geometry Shared Pairs = 3 Unshared Pairs = 1
- 79. 5.Tetrahedral Ball and stick model Moleculegeometry Unshared Pairs = 0
- 80. • I candescribe the 3 intermolecular forces of covalent compounds and explain the effects of each force.
- 81. Like Dissolves Like (memorizethis!) • Polar substances WILL only dissolve in polar solvents • Non-polar substances WILL only dissolve in nonpolar solvents • Non-polar substances WILL NOT dissolve in polar solvents and vice versa • Example: acetone and styrofaom Video – Styrofoam and acetone
- 82. • Attractions within orinside molecules, also known as bonds. –Ionic –Covalent –metallic Intramolecular attractions Roads within a state
- 83. • Attractions betweenmolecules – Hydrogen “bonding” • Strong attraction between special polar molecules (F, O, N, P) – Dipole-Dipole • Result of polar covalent Bonds – Induced Dipole (Dispersion Forces) • Result of non-polar covalent bonds Intermolecular attractions Interstate highways connect one state to another
- 84. More on intermolecularforces Hydrogen “Bonding” • STRONG attraction between highly polar molecules • HF, H2O, NH3 • Occurs between H of one molecule and N, O, F of another molecule Hydrogen “bond” - + + - + + + + - Hydrogen bonding 1 min
- 85. Why does Hydrogen “bonding”occur? • Nitrogen, Oxygen and Fluorine –are small atoms with strong nuclear charges • powerful atoms –N, O & F have very high electronegativities, these atoms strongly attract shared electrons in a bond –Create very POLAR molecules
- 86. Dipole-Dipole Interactions – WEAKintermolecular force – Bonds have high EN differences forming polar covalent molecules, but not as high as those that result in hydrogen bonding. .21<EN<1.99 – Partial negative and partial positive charges slightly attracted to each other. – Result of polar covalent bonds – Only occur between polar covalent molecules
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- 88. Induced Dipole Attractions –VERY WEAK intermolecular force – Result of nonpolar covalent bonds – Temporary partial poles are caused by a nearby polar covalent molecule disturbing the arrangement of electrons in the nonpolar molecule – Occur briefly between NONPOLAR & POLAR molecules Induced dipole video 30 sec
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- 90. Intermolecular Forces affect chemicalproperties • For example, strong intermolecular forces cause high boiling point –Water has a high boiling point compared to many other liquids due to the hydrogen bonding between water molecules.
- 91. Predict which substancehas the highest boiling point. • HF • NH3 • CO2 19.5 ˚C (67.1 F ; 292.6 K) −33.3 ˚C (−28.0 ˚ F ; 239.8 K) −56.6 ˚C (−69.6 ˚ F ; 216.6 K)
- 92. Why HF hasthe highest boiling point • HF • NH3 • CO2 • WHY? The H-F bond has the highest electronegativity difference SO HF is a highly polar molecule resulting in the Hydrogen bonding. Therefore HF needs the most energy to overcome the intermolecular forces and boil.
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Editor's Notes
- #66 PER 5 11/14/2107