chemical bonding. what is a bond? force that holds atoms together results from the simultaneous...
TRANSCRIPT
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– Endothermic
• When a bond is formed energy is released– 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 following1.Mg
2.C
3.Ne
Dot Diagrams - Ions
• For ions, use brackets and place the charge outside the brackets
• Examples: 1. Na+
2. O2-
3. H+
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”
Identifying Bond Type
• Ionic – metal and a nonmetal
• Covalent – 2 nonmetals– Nonpolar Covalent (equal)
• Same atoms (diatomics, triatomics)
– Polar Covalent (unequal)• Unequal electronegativity
• Metallic – metals
Identifying Bond Types
• Indicate the type of bond present in each: 1. HCl
2. CCl43. MgCl24. O2
5. Hg
6. H2O
Ionic Bonds• Transfer of 1 or more electrons
from a metal to a nonmetal• The greater the electronegativity
difference between atoms, the greater the ionic character
• Example: Sodium Chloride (NaCl)
Na ClX
Na electron transferred to Cl
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+ + Cl1-
2. Mg2+ + Cl1-
3. Ca2+ + CO32-
4. Al3+ + O2-
Criss Cross Method
1. Write the symbol for the cation and anion
2. Write each ion’s 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. Mg2+ + Cl-
2. Ca2+ + CO32-
3. Al3+ + O2-
4. Calcium ion + hydroxide ion
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. Iron (II) Chloride
2. Iron (III) Oxide
3. Copper (II) Oxide
4. 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– All solids at STP
• Hard
• Brittle
Melting and Boiling Points of Compounds
Compound Name
Formula Type of Compound mp
(oC)
bp
(oC)Magnesium Flouride MgF2 Ionic 1261 2512
Sodium Chloride NaCl Ionic 801 1686
Calcium Iodide CaI2 Ionic 784 1373
Iodine MonoChloride ICl Covalent 27 370
Carbon tetrachloride CCl4 Covalent -23 350
Hydrogen Flouride HF Covalent -83 293
Hydrogen Sulfide H2S Covalent -86 212
Methane CH4 Covalent -182 109
Properties of Salts (cont’d)
• 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
Covalent Bonds• Sharing of electrons between 2 nonmetals
Non-Polar Covalent
• Electrons are shared equally• Equal distribution of electrons
• All diatomic molecules have non-polar covalent bonds
Nonpolar Covalent Examples
1. Flourine (F2)
2. Hydrogen (H2)
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
• The greater the electronegativity difference, the more polar the bond is
Polar Covalent Examples
1. HCl
2. H2O
Dipoles
• Form when the charge in a bond is asymmetrical
– Present in polar bonds– Partial positive and partial negative charges
Polar Bonds / Dipoles
• Isn’t a whole charge just a partial charge • means a partially positive • means a partially negative
Example:
H - Cl
+---→• 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. CCl4b. CH4
2. Which has a stronger dipole? a. HCl
b. HBr
Properties of Molecular Substances (Covalent Compounds)
• Soft
• Low melting points and boiling points– Many exist as gases or liquids at STP
• Poor 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- is only used for the second element– Example: CO = carbon monoxide
Mono- 1 Hexa- 6
Di- 2 Hepta- 7
Tri- 3 Octa- 8
Tetra- 4 Nona- 9
Penta- 5 Deca- 10
Examples
1. CCl42. H2O
3. NO
4. N2O5
5. 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)– All solids at STP (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: H2, CO2, HCl
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 CH H
H
H109.5º
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
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
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
• Polar molecules will have stronger IMFs than nonpolar molecules– The greater the polarity the stronger the IMF
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. HCl
b. 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
HHO
+ -
+
H HO+-
+
Hydrogen Bonds
• VERY STRONG
• Molecules with H bonds will have high boiling points, melting points, and surface tension– See Water Heating
• Example: NH3
H-bonds Examples
1. Which sample has Hydrogen Bonds?
a. H2 b. HF c. F2 d. HCl
2. Rank in order from strongest (1) to weakest (3). a. Hydrogen Bonds
b. Covalent Bonds
c. 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) c. KCl(l)b. HCl(aq) d. KCl(aq)
2. When sodium chloride dissolves in water the chloride ion is attracted toa. The positive part of the water, the O atomb. The negative part of the water, the O atomc. The positive part of the water, the H atomd. 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