Bonding Notes
Bonding - how atoms are held together, determines many aspects of chemical reactivity and physical properties
Bonding ideas are based on models - which are simplifications of the way the molecules behave- we must remember that the model gives us a guide for the behavior of atoms, but exceptions do occur- frequently - since the model is intended to simplify reality.
I. Types of bonds
A. Ionic
1. definition- metal and nonmetal react, both form ions which are attracted to each other - negative and positive ions attract
cation - loses electron (+ charge)
anion - gains electrons (- charge)
2. Why? - bonds form between metal and nonmetals in this way since the energy of the system is lower than in another form - the ion pair has a lower energy than the separated ions. In general systems will find the lowest energies - they favor states in which their energy is lowest
3. The ions:
- group 1 forms +1 ions, group 2 forms +2 ions, group 3 forms +3 ions (loss of electrons - valence shell 8 - stable octet)
- group 6 forms -2 ions, group 7 forms -1 ions (gain electrons - valence 8)
4. Transition metals may have > 1 cation formed - review transition metal chemistry - recall all transition metals can be +2 and then others may also form.
5. Ionic bonds are mainly between metals and nonmetals with a large difference in electronegativity values (larger EN difference , greater IONIC character)
6. Solid phase of ionic compound is typically a lattice of ions - packed three dimensionally. The unique stability of ionic compounds refers to the solid state where many ions are held by ionic forces in all directions
7. Lattice energy - the energy released when an ionic solid forms from its ions (this process releases energy and the resulting solid is less energetic and more stable) Lattice enthalpy increases with increasing ionic charge and with decreasing ionic size. So ions that are small and have a bigger charge have a larger lattice enthalpy (take more energy to break apart) than ions that are larger and have a lesser charge (attraction between ions is greater with smaller ions and more charged ions.)
8. ionic compound = a salt - conducts an electric current when in liquid phase (not in solid phase as ions held tight in lattice structure)
B. Covalent
1. Bonds form as a result of the tendency to seek the lowest possible energy.
2. Covalent bonds do not have ions - rather electrons are shared between atoms
3. Bonds can be equal sharing of electrons (identical atom or atoms with identical electronegativities: Cl-Cl, C-C, C-S. ) These bonds are NONPOLAR or PURE covalent bonds
- each atom shares / needs the electrons equally
4. bonds of unequal sharing result from attraction of atoms with different electronegativities (EN) (unequal desire for electrons) - these bonds are POLAR covalent bonds. The greater the difference in electronegativity between the atoms, the greater the polarity
Polar bonds = unequal sharing of electrons
C-O (2 non metals → covalent bond) polar bond since EN C=2.5 and EN O = 3.5
O is more desirous of electrons than carbon so those electrons are thought to be found closer to oxygen than carbon - the bond has a more negative side close to oxygen and a more positive side next to carbon we use the symbol delta + (δ+) and delta – (δ -) to show the polarity of bonds and molecules. We also use an arrow with a + side (the point represents the - side) One could say that polar covalent bonds are partially ionic in character although one atom is not seen as taking the electrons and one atom as given the electrons
5. Multiple bonds
Covalently bonded molecules can share more than a single pair of electrons: single bond = 1 pair (2 electrons) is shared between adjacent atoms; double bond - 2 pairs of electrons (4 e-) are shared between adjacent atoms; triple bond = three pairs of electrons (6 e-) are shared between adjacent atoms
6. Bond length - as more electrons are shared, the length of the bond (distance between atoms) decreases. Triple bond length< double bond length<single bond length
7. Bond strength is a measure of “Bond energy”→ the average energy required to break the bond between 2 atoms. Bond energies (enthalpies) depend not only upon the 2 atoms in question, but also upon the atoms to which those atoms are bonded - therefore average bond energy is used. Table of bond energies is available; units are kJ/mol (higher the bond enthalpy the stronger the bond - more difficult it is to break)
a. carbon bond strength: single< double<triple since electrons are held within a smaller area and is more difficult to separate.
b. C-O and C=O in carboxylic acid
C-O length is longer than C=O (0.143 vs. 0.122)
C=O strength is more than C-O (799 vs. 358)
***use bond energies to determine changes in energy for chemical rxn and therefore predict whether the reaction will occur spontaneously or not
8. Electron distribution in molecules see examples:
O2, N2, CO2, C2H4, C2H2 - see also Lewis dot
C. Metallic bonds
1. Definition: forces that keep metal atoms together
2. Properties of metals
a. malleability (sheets),
b. ductility (wires),
c. good conductors
d. high mp
3. Properties lead to the observation that metals are bonded in a strong and non- directional manner (difficult to separate atoms but easy to move them around)
4. “electron sea model” - metal cations in a sea of valence electrons
a. cations are arranged in a lattice- a regular frame in which rows of cations can be moved (hammered, pulled)
b. electrons can move to conduct current metal. Cations can be hammered, pulled into wires - non rigid structure
c. the electrons do not “belong” to a single cation, but are delocalized and may move about
II. Shapes of molecules VSPER theory and Lewis dot
A. Lewis dot structures
1. What is it? - shows the arrangement of the valence electrons in an atom - helps identify stable molecules
2. Octet rule - atoms will be most stable energetically if they have a complete octet in their valence (outermost) shell.
3. Electrons can be distinguished as bonding pairs - involved in a bond between atoms, and lone pairs -electrons not involved in bonding.
4. Guidelines for writing lewis dot structures
a. total # valence electrons are more important than identifying where the valence electrons come from.
b. start with single bonds and progress to multiple bonds
c. hydrogen needs only 2 electrons to be complete, all others need eight
d. show all electrons (even lone pairs) in structure
5. Exceptions to the octet rule
a. boron forms bonds with only 3 pairs of electrons -(BF3) due to small size and large + charge
b. > octet
SF6 - stable molecule - extra electrons can be placed in the empty d orbitals (so third row and heavier elements MAY exceed octet rule by using empty d orbitals) - examples - ClF3, XeO3, RnCl2, BeCl2 ICl4
B. Resonance - more than one Lewis structure may be written for a compound
1. the “actual” structure is an average of the resonance structures
2. electrons are not always “stuck” between two atoms, but do have the ability to move about a molecule - and act in a delocalized manner.
3. delocalized electrons move around a molecular structure and do not fit our model of Lewis structures (but in nature - true structures are explained this way) - examples CO3-2 and benzene (C6H6)
C. VSEPR theory - used to predict the shapes of molecules (molecular geometry)
1. the shape of a molecule is important in its reactivity and chemical properties
2. VSEPR (valence shell electron pair repulsion model) - structure is determined by minimizing electron pair repulsions
3. electrons will be arranged as far away from each other as possible (to minimize their interaction (will repulse negative charge)
4. the shapes and bond angles *use models to show shapes
a. linear 180 between electron pairs BeCl2 (flat)
b. trigonal planar - 120 degrees between 3 pairs (flat) BF3
c. tetrahedral - 109.5 degrees between 4 pairs (3 dimensional) CH4
d. trigonal bipyramid - 90 degrees and 120 degrees - 5 pairs
e. octahedral - 90 degrees 6 pairs
5. when the electron pairs are non bonding but lone pairs, the electrons take more space and change the shape of the molecule, for example NH3 3 bonding pairs and 1 lone pair - forms TRIGONAL PYRAMIDAL shape (107 ° rather than 109.5 ° between bonding pairs)
H2O 2 bonding and 2 lone pairs 104.5 ° between bonding pairs since lone pairs take up more space than bonding pairs – forms a BENT or ANGULAR shape
XeF4 - 4 bonding and 2 lone pairs - square planar molecule ( 90° bond angle)
***lone pairs will occupy positions that have the most space when available in trigonal bipyramid the 120 ° position will be occupied by lone pairs
6. multiple bonds - for determination of structure, multiple bonds count as a single pair
7. molecule polarity determination: polarity of molecule depends on the type and arrangement of bonds.
a. when all bonds are identical and evenly spaced (ie. Symmetrical), the molecule is nonpolar
b. when bonds are not identical or not evenly spaced, the molecule is polar has a dipole moment
III. Intermolecular forces
Whereas bonds are intramolecular (within molecules), intermolecular forces are forces between molecules that hold molecules together (not atoms within a molecule).
- these forces are weaker than bonds
- changes in phase are accompanied by breaking of intermolecular forces
A. dipole-dipole attraction
1. Attractive force between polar molecules wherein the positive side of a polar molecule is attracted to the negative side of an adjacent polar molecule
2. The farther apart the molecules, the weaker the force (forces among solid molecules and among liquid molecules would be much stronger that forces in gases - proximity of molecules)
B. hydrogen bond (misnomer since not a “true” bond)
1. special type of dipole attraction
2. dipole of high electronegative atoms N,O, F and H atom.
H2O experiences H bonds - a stronger dipole attraction (given the high EN of O)
NH3 has H bonds
HF has H bonds
3. both the high polarity (EN difference) and the small size of H atom(which allows the dipole to be relatively close to the H atom) account for the strength of the force
4. The H bond intermolecular force forms between the lone pair electrons (of the O, F, or N atom in one molecule) and the H in another molecule. So since water has 2 lone pairs in the molecule, it can form 2 H bonds between molecules whereas ammonia, with just 1 lone pair, can form only 1 H bond between molecules.
5. Substances with lone pair electrons (but not HF, HO, or HN bonds) may form H bonds between different molecules, such as carboxylic acids and water molecules (the lone pairs on the carbonyl group may form an IM force with water molecules and hence water solubility is enhanced by a molecules availability to form H bonds
C. van der Waal`s forces - forces among nonpolar molecules (aka London dispersion forces) *** note- all molecules have vdW forces but are considered insignificant in polar molecules
1. Instantaneous dipoles can form (since electrons move within their orbitals)
2. This dipole can INDUCE another dipole to form and create a chain rxn of induced dipoles
3. Surface area of molecule determines strength of vdW. Long molecules have greater vdW forces (greater area for separation of induced dipole) than ball shaped molecules.
4. Strength of van der Waal`s forces < dipole < hydrogen bonding
D. Intermolecular forces and Boiling Point
1. Boiling Point = temp at which solid → liquid
2. Greater IM forces → higher BP
3. Since forces in H bonds are biggest, those molecules with H bond would be expected to have higher BP than similar molecules without H bonds.
4. van der Waals forces increase with increasing molecular mass ( mass increase → increase in # of electrons → greater chance of instantaneous dipole → greater chance of induced dipoles → stronger intermolecular force)
see BP of H2O vs H2S
NH3 vs. PH3
C3H8 vs. CH3CHO vs. C2H5OH
IV. Physical Properties based on bonding.
A. bonding and intermolecular force strength determine the physical properties
B. bonding
1.ionic - high mp, non-volatile, conduct as molten and in solution, solubility- recall rules - many ionic compounds are water soluble
2. covalent- low mp - usually liquid and low melting solids at room temp, non-conductors of electricity, more soluble in nonpolar solvent (mineral oil lab)
3. metals - metallic bonds - mp varies with # of valence electrons; good conductors - all phases (sea of electrons), not soluble in water (may make mixtures of metals - alloys like brass, steel)
4. network solid allotropes of carbon (macromolecules - more later)
C. intermolecular forces
1. mp- within group (halogen) increases down table as Molecular Mass increases, vdW forces increase
2. H bond>dipole> van der Waals
3. normally mp determination is a very small range at which solid → liquid
4. with impure substances mp are over a wide range of temperature
5. ionic > polar covalent > nonpolar covalent (pure covalent)
6. bp – liquids → gas - intermolecular forces are key again