Chemical Bonding What is a Bond? Force that holds atoms together Results from the simultaneous attraction of electrons (-) to the nucleus (+) Breaking/Forming Bonds When a bond is broken energy is absorbed
When a bond is formed energy is released Endothermic Exothermic The greater the energy released during the formation of the bond, the greater its stability
Stable bonds require a great deal of energy to break Lewis Dot Diagrams Use dots to represent the number of valence electrons How to write: Write the symbol.
Put one dot for each valence electron Electrons go on the 4 sides, no more than 2 per side Dot Diagram Examples: Draw dot-diagrams for the following 1. 2. 3. Mg C Ne
Dot Diagrams - Ions For ions, use brackets and place the charge outside the brackets Examples: 1. 2. 3. Na+ O2H+ Octet Rule
Atoms will gain or lose electrons in order to have a full valence shell like the nobles gases Take the shortest route Metals lose electrons to form positive ions (Cations) Nonmetals gain electrons to form negative ions (Anions)
Exceptions 1st principle energy level only holds 2 electrons Transition elements can lose valence (s) and inner (d) electrons this is why they have multiple oxidation states Some atoms may be stable with less than an octet many compounds with B Some atoms may be stable with more than an octet
elements beyond period 2, especially P and S, the additional electrons are added to the d sublevel Molecules with an odd number of electrons they will be unstable Types of Bonds Ionic - Electrons are transferred from a metal to a nonmetal Covalent - Electrons are shared between 2 nonmetals
Polar Covalent electrons are shared unequally Nonpolar Covalent electrons are shared equally Metallic - Electrons are mobile within a metal, Sea of Electrons Dog Analogy Ionic Bonds
Polar Covalent Bonds Dogs of equal strength Metallic Bonds Unevenly matched but willing to share Nonpolar Covalent Bonds
big greedy dog stealing the other dogs bone Mellow dogs with plenty of bones to go around See the Dogs Identifying Bond Type Ionic metal and a nonmetal Covalent 2 nonmetals Metallic metals OR Use electronegativity differences
Ionic: 1.7 or more Polar Covalent: 0.5-1.6 Nonpolar Covalent: 0.0-0.4 Identifying Bond Types Indicate the type of bond present in each: 2. HCl
CCl4 3. MgCl2 4. O2 5. Hg H 2O 1.
6. Ionic Bonds Transfer of 1 or more electrons from a metal to a nonmetal Electronegativity difference is 1.7 Example: Sodium Chloride (NaCl) Na electron transferred to Cl Na Cl X
Monatomic Ions One atom in an ion Look at the valence electrons to determine the charges Examples: K+, O2- Polyatomic Ions More than one atom in the ion Reference Table E Charge belongs to the entire ion, not an individual atom Within the polyatomic ion the atoms are held together by covalent bonds When writing it, place ( ) around the entire ion, with the charge outside Examples: (NH4)+, (H3O)+, (CO3)2-
Writing Ionic Formulas You need an equal amount of positive and negative charges, so that the compound is neutral Ionic Formulas are always written as empirical formulas (reduced) Examples 1. Na1+ + Cl12. Mg2+ + Cl13. Ca2+ + CO324. Al3+ + O2- Criss Cross Method 1. Write the symbol for the cation and anion 2. Write each ions charge as a superscript 3. Criss-cross the charges to become
subscripts of the other ion Do not put (+) or (-) charges in the final formula 4. Reduce to least common multiple (empirical formula) Ionic Formulas Write the formula for the compound formed from the following ions: 1. 2. 3. 4.
Mg2+ + ClCa2+ + CO32Al3+ + O2Ca + OH Naming Ionic Compounds Name the cation first, the anion second Cation keeps its name, anion changes its ending to ide (Chlorine Chloride) Do not change the ending of polyatomic ions Examples: 1. NaCl 2. CaCO3 3. MgF2 Stock System only used for positive ions Some cations have more than one
positive oxidation states A roman numeral is used to indicate the charge of the positive ion Stock System Examples 1. 2. 3. 4. Iron (II) Chloride Iron (III) Oxide Copper (II) Oxide a. What charge does copper have in copper II sulfate? b. What is the formula for copper II sulfate?
Ionic Salts Salts are ionic compounds made up of cations and anions The ratio of cations to anions is always such that an ionic compound has no overall charge Many of the ions are bonded together to form a crystal Properties of Ionic Salts Ionic Bonds are very strong Very high melting and boiling points Hard Brittle Properties of Salts (contd)
Do not conduct electricity as solids Do conduct electricity when the salt melts or is dissolved in water (liquid phase or aqueous) In order to conduct electricity a substance must have free moving charged particles In the solid phase the ions are not free to move Melting and Boiling Points of Compounds Compound Name Formula Type of Compound
mp (oC) bp (oC) Magnesium Flouride MgF2 Ionic 1261 2512
Covalent -182 109 Covalent Bonds Sharing of electrons between 2 nonmetals Electronegativity difference is 1.6 Non-Polar Covalent Electrons are shared equally
Uniform distribution of electrons Bond is symmetrical Electronegativity difference of 0-0.4 All diatomic molecules have non-polar covalent bonds Nonpolar Covalent Examples 1.
Flourine (F2) a. b. 2. e-neg difference = Dot diagram: Hydrogen (H2) a. b. e-neg difference = Dot diagram:
Polar Covalent Unequal Sharing of electrons Unequal distribution of electrons Partial positive and partial negative charges The side with the higher electronegativity will have a greater share of the electron(s)
resulting in a partial negative charge Electronegativity difference of 0.5-1.6 Polar Covalent Examples 1. HCl a. b. 2. e-neg difference: Dot diagram:
H2O a. b. e-neg difference: Dot diagram: Dipoles Form when the charge in a bond is asymmetrical Present in polar bonds Partial positive and partial negative charges
Polar Bonds / Dipoles Isnt a whole charge just a partial charge means a partially positive means a partially negative Example: - Cl H +--- The Cl pulls harder on the electrons (more eneg) The electrons spend more time near the Cl Dipole Examples
1. Which molecule contains more polar bonds? a. CCl4 b. CH4 2. Which has a stronger dipole? a. b. HCl HBr Properties of Molecular Substances (Covalent Compounds)
Soft Low melting points and boiling points Many Poor exist as gases conductors of heat and electricity (in all phases) Examples: H2O, CCl4, NH3, C6H12O6, O2 Molecular Formulas
(Covalent Compounds) Contain covalent bonds Tells you how many atoms are present in a single molecule Named similarly to ionic compounds, except use prefixes to indicate the number of atoms per molecule Prefixes Mono- 1 Hexa-
6 Di- 2 Hepta- 7 Tri- 3 Octa-
8 Tetra- 4 Nona- 9 Penta- 5 Deca-
10 Mono- is only used for the second element Example: CO = carbon monoxide Examples 1. 2. 3. 4.
5. CCl4 H2O NO N2O5 BBr3 Structural Formulas Specifies how atoms are bonded together Dashes represent bonds 2 atoms can share up
to 3 pairs of electrons Single Bonds 2 atoms share 1 pair of electrons (2 electrons) Examples: 1. Ammonia (NH3) 2. Chlorine (Cl2) 3. Hydrochloric Acid (HCl) Double Covalent Bonds 2 atoms share 2 pairs of electrons (4
electrons) 2 bonds between 2 atoms Examples: 1. Carbon Dioxide (CO2) 2. Oxygen (O2) Triple Covalent Bond 2 atoms share 3 pairs of electrons (6 electrons) 3 bonds between 2 atoms Examples: 1. Nitrogen (N2) 2. Ethyne (C2H2)
Bond Length/Strength Length: Single > Double > Triple The more electrons in a bond, the greater the attraction, therefore shorter As you move down a group bond length increases Due
to increasing molecular size Strength: Triple is the strongest, most stable, requires the most energy to break Network Solids Covalently bonded atoms are linked into a giant network (macromolecules)
Examples: Diamond (C), Graphite (C), Silicon Carbide (SiC), and Silicon Dioxide (SiO2) Network Solids Properties: Hard High melting and boiling points Do not conduct heat and electricity
Metallic Bonding Sea of Electrons Electrons are free to move through the solid. + + + + + + + + + + + +
Properties of Metallic Solids Very Strong Good conductors of heat and electricity because electrons are free to move about Luster High melting point (except Hg) Malleable, Ductile VSEPR Theory In a small molecule, the electron pairs are as far away from each other as possible VSEPR
= Valence Shell Electron Pair Repulsion Linear Drawn on a straight line All molecules of only 2 atoms are linear Many 3 atom molecules are linear, if there are no unshared electron pairs on the central atom If both ends are the same, the molecule is nonpolar (Symmetrical = Nonpolar) If the ends are different, the molecule will be polar (Asymmetrical = Polar) Bond Angle = 180o
See Molecules Examples: H , CO , HCl 2 2 Trigonal Planar Trigonal Planar A central atom is bonded to 3 other atoms, with no extra electrons on the central atom Forms a flat Y shape (triangle shape) If the ends are all the same, NONPOLAR If the ends are different, POLAR Bond Angle = 120o
See Molecules Examples: BCl3, BH2F Pyramidial A central atom is bonded to 3 other atoms and the central atom has an unshared electron pair 3-D, like a pyramid Always POLAR Bond Angle = 107o See Molecules Example: NH3
Tetrahedral Tetrahedral A central atom bonded to 4 other atoms 3-D shape allows the electron pairs to get as far away from each other as possible H H
109.5 C H H Tetrahedral If all the ends are the same, NONPOLAR If the ends are different, POLAR Bond Angle = 109.5o See Molecules Examples: 1. CH4
2. CH3Cl Bent Bent A central atom is bonded to 2 other atoms and the central atom has 2 unshared electron pairs Always POLAR Bond angle = 105o See Molecules Example: H2O Intermolecular Attractions/
Forces Forces between molecules Determines boiling point, melting point, vapor pressure, surface tension The stronger the intermolecular attractions, the higher the boiling point All intermolecular attractions are
weaker than actual bonds Dipole-Dipole Forces Occurs between 2 polar molecules The positive end of one molecule is attracted to the negative end of another molecule The greater the electronegativity difference is, the more polar the bond
will be and the stronger the dipole will be Example: HCl Dipole Examples 1. Which would have the strongest intermolecular forces? Explain Why. a. b. HCl HBr 2. Which would have the weakest intermolecular forces? Explain Why. a. H2S
b. H2O Hydrogen Bonds Special, Strong type of Dipole Attractions Attraction of a covalently bonded H atom to a F, O, or N atom on another covalent compound + -
H O H + - + H O +
H Hydrogen Bonds VERY STRONG Molecules with H bonds will have high boiling points, melting points, and surface tension Example: NH3 H-bonds Examples 1. Which sample has Hydrogen
Bonds? a. H2 b. HF c. F2 d. HCl 2. Which is the strongest? a. b. c. Hydrogen Bonds Covalent Bonds
Dipole-Dipole Attractions Molecule Ion Attractions Attraction between a polar compound and an ion (ionic salt) Polar substances (such as water) attract ions from ionic compounds in solution This allows ionic substances to dissolve in polar solvents (water)
The anion is attracted to the positive end of the polar solvent The cation is attracted to the negative end of the polar solvent The ion dissociates (falls apart) Example: NaCl(aq) Molecule-Ion Examples 1. Molecule-Ion attractions are present in which sample? a. HCl(l) b. HCl(aq) c. KCl(l) d. KCl(aq)
2. When sodium chloride dissolves in water the chloride ion is attracted to a. The positive part of the water, the O atom b. c. d. The negative part of the water, the O atom The positive part of the water, the H atom The negative part of the water, the H atom Van Deer Waals Forces
Very weak Exist between non-polar molecules Caused by momentary dipoles Increases as molecular mass increases VDW Examples 1. Rank in order from weakest to strongest:
Hydrogen Bonds Covalent Bonds Van deer Waals Forces Dipole-Dipole Attractions 2. Which would have the strongest intermolecular forces? a. H2 b. Cl2 c. F2 d. Br2
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