Presentation on theme: "Chapter 8 Bonding. What is a Bond? l A force that holds atoms together. l Why? l We will look at it in terms of energy. l Bond energy- the energy required."— Presentation transcript:
Chapter 8 Bonding
What is a Bond? l A force that holds atoms together. l Why? l We will look at it in terms of energy. l Bond energy- the energy required to break a bond. l Why are compounds formed? l Because it gives the system the lowest energy.
Ionic Bonding l An atom with a low ionization energy reacts with an atom with high electron affinity. l A metal and a non metal l The electron moves. l Opposite charges hold the atoms together.
Coulomb's Law l E= 2.31 x J · nm(Q 1 Q 2 )/r l Q is the charge. l r is the distance between the centers. l If charges are opposite, E is negative l exothermic l Same charge, positive E, requires energy to bring them together.
What about covalent compounds? l The electrons in each atom are attracted to the nucleus of the other. l The electrons repel each other, l The nuclei repel each other. l The reach a distance with the lowest possible energy. l The distance between is the bond length.
0 Energy Internuclear Distance
0 Energy Internuclear Distance
0 Energy Internuclear Distance
0 Energy Internuclear Distance
0 Energy Internuclear Distance Bond Length
0 Energy Internuclear Distance Bond Energy
Covalent Bonding l Electrons are shared by atoms. l These are two extremes. l In between are polar covalent bonds. l The electrons are not shared evenly. l One end is slightly positive, the other negative. Indicated using small delta
H - F + -
Electronegativity l The ability of an electron to attract shared electrons to itself. l Pauling method l Imaginary molecule HX l Expected H-X energy = H-H energy + X-X energy 2 = (H-X) actual - (H-X) expected
Electronegativity is known for almost every element l Gives us relative electronegativities of all elements. l Tends to increase left to right. l decreases as you go down a group. l Most noble gases arent discussed. l Difference in electronegativity between atoms tells us how polar the bond is.
Electronegativity difference Bond Type Zero Intermediate Large Covalent Polar Covalent Ionic Covalent Character decreases Ionic Character increases
Dipole Moments l A molecule with a center of negative charge and a center of positive charge is dipolar (two poles), l or has a dipole moment. l Center of charge doesnt have to be on an atom. l Will line up in the presence of an electric field.
H - F
How It is drawn H - F + -
Which Molecules Have Dipoles? l Any two atom molecule with a polar bond. l With three or more atoms there are two considerations. 1)There must be a polar bond. 2)Geometry cant cancel it out.
Geometry and polarity l Three shapes will cancel them out. l Linear
Geometry and polarity l Three shapes will cancel them out. l Planar triangles 120º
Geometry and polarity l Three shapes will cancel them out. l Tetrahedral
Geometry and polarity l Others dont cancel l Bent
Geometry and polarity l Others dont cancel l Trigonal Pyramidal
Ions l Atoms tend to react to form noble gas configuration. l Metals lose electrons to form cations l Nonmetals can share electrons in covalent bonds. –When two non-metals react.(more later) l Or they can gain electrons to form anions.
Ionic Compounds l We mean the solid crystal. l Ions align themselves to maximize attractions between opposite charges, l and to minimize repulsion between like ions. l Can stabilize ions that would be unstable as a gas. l React to achieve noble gas configuration
Size of ions l Ion size increases down a group. l Cations are smaller than the atoms they came from. l Anions are larger. l across a row they get smaller, and then suddenly larger. l First half are cations. l Second half are anions.
Periodic Trends l Across the period nuclear charge increases so they get smaller. l Energy level changes between anions and cations. Li +1 Be +2 B +3 C +4 N -3 O -2 F -1
Size of Isoelectronic ions l Positive ions have more protons so they are smaller. Al +3 Mg +2 Na +1 Ne F -1 O -2 N -3
Forming Ionic Compounds l Lattice energy - the energy associated with making a solid ionic compound from its gaseous ions. M + (g) + X - (g) MX(s) l This is the energy that pays for making ionic compounds. l Energy is a state function so we can get from reactants to products in a round about way.
Na(s) + ½F 2 (g) NaF(s) First sublime NaNa(s) Na(g) H = 109 kJ/mol Ionize Na(g) Na(g) Na + (g) + e - H = 495 kJ/mol Break F-F Bond½F 2 (g) F(g) H = 77 kJ/mol Add electron to F F(g) + e - F - (g) H = -328 kJ/mol
Na(s) + ½F 2 (g) NaF(s) Lattice energy Na + (g) + F - (g) NaF(s) H = -928 kJ/mol
Calculating Lattice Energy l Lattice Energy = k(Q 1 Q 2 / r) l k is a constant that depends on the structure of the crystal. l Qs are charges. l r is internuclear distance. l Lattice energy is with smaller ions l Lattice energy is greater with more highly charged ions.
Calculating Lattice Energy l This bigger lattice energy pays for the extra ionization energy. l Also pays for unfavorable electron affinity. l O 2- (g) is unstable, but will form as part of a crystal
Partial Ionic Character l There are probably no totally ionic bonds between individual atoms. l Calculate % ionic character. l Compare measured dipole of X-Y bonds to the calculated dipole of X + Y - the completely ionic case. l % dipole = Measured X-Y x 100 Calculated X + Y - l In the gas phase.
% Ionic Character Electronegativity difference 25% 50% 75%
How do we deal with it? l If bonds cant be ionic, what are ionic compounds? l And what about polyatomic ions? l An ionic compound will be defined as any substance that conducts electricity when melted. l Also use the generic term salt.
The Covalent Bond l The forces that causes a group of atoms to behave as a unit. l Why? l Due to the tendency of atoms to achieve the lowest energy state. l It takes 1652 kJ to dissociate a mole of CH 4 into its ions l Since each hydrogen is hooked to the carbon, we get the average energy = 413 kJ/mol
l CH 3 Cl has 3 C-H, and 1 C - Cl l the C-Cl bond is 339 kJ/mol l The bond is a human invention. l It is a method of explaining the energy change associated with forming molecules. l Bonds dont exist in nature, but are useful. l We have a model of a bond.
What is a Model? l Explains how nature operates. l Derived from observations. l It simplifies them and categorizes the information. l A model must be sensible, but it has limitations.
Properties of a Model l A human inventions, not a blown up picture of nature. l Models can be wrong, because they are based on speculations and oversimplification. l Become more complicated with age. l You must understand the assumptions in the model, and look for weaknesses. l We learn more when the model is wrong than when it is right.
Covalent Bond Energies l We made some simplifications in describing the bond energy of CH 4 l Each C-H bond has a different energy. CH 4 CH 3 + H H = 435 kJ/mol CH 3 CH 2 + H H = 453 kJ/mol CH 2 CH + H H = 425 kJ/mol CH C + H H = 339 kJ/mol l Each bond is sensitive to its environment.
Averages l There is a table of the averages of different types of bonds pg. 365 l single bond- one pair of electrons is shared. l double bond- two pair of electrons are shared. l triple bond- three pair of electrons are shared. l More bonds, more bond energy, but shorter bond length.
Using Bond Energies We can estimate H for a reaction. l It takes energy to break bonds, and end up with atoms (+). l We get energy when we use atoms to form bonds (-). If we add up the energy it took to break the bonds, and subtract the energy we get from forming the bonds we get the H. l Energy and Enthalpy are state functions.
Find the energy for this 2 CH 2 = CHCH 3 + 2NH 3 O2O2 + 2 CH 2 = CHC N + 6 H 2 O C-H 413 kJ/mol C=C 614kJ/mol N-H 391 kJ/mol O-H 467 kJ/mol O=O 495 kJ/mol C N 891 kJ/mol C-C 347 kJ/mol
Localized Electron Model l Simple model, easily applied. l A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. l Three Parts 1) Valence electrons using Lewis structures 2) Prediction of geometry using VSEPR 3) Description of the types of orbitals (Chapt 9)
Lewis Structure l Shows how the valence electrons are arranged. l One dot for each valence electron. l A stable compound has all its atoms with a noble gas configuration. l Hydrogen follows the duet rule. l The rest follow the octet rule. l Bonding pair is the one between the symbols.
Rules l Sum the valence electrons. l Use a pair to form a bond between each pair of atoms. l Arrange the rest to fulfill the octet rule (except for H and the duet). lH2OlH2O l A line can be used instead of a pair.
A useful equation l (happy-have) / 2 = bonds l CO 2 C is central atom l POCl 3 P is central atom l SO 4 2- S is central atom l SO 3 2- S is central atom l PO 4 3- P is central atom l SCl 2 S is central atom
Exceptions to the octet l BH 3 l Be and B often do not achieve octet l Have less than an octet, for electron deficient molecules. l SF 6 l Third row and larger elements can exceed the octet l Use 3d orbitals? l I 3 -
Exceptions to the octet l When we must exceed the octet, extra electrons go on central atom. l (Happy – have)/2 wont work l ClF 3 l XeO 3 l ICl 4 - l BeCl 2
Resonance l Sometimes there is more than one valid structure for an molecule or ion. l NO 3 - l Use double arrows to indicate it is the average of the structures. l It doesnt switch between them. l NO 2 - l Localized electron model is based on pairs of electrons, doesnt deal with odd numbers.
Formal Charge l For molecules and polyatomic ions that exceed the octet there are several different structures. l Use charges on atoms to help decide which. l Trying to use the oxidation numbers to put charges on atoms in molecules doesnt work.
Formal Charge l The difference between the number of valence electrons on the free atom and that assigned in the molecule or ion. l We count half the electrons in each bond as belonging to the atom. l SO 4 -2 l Molecules try to achieve as low a formal charge as possible. l Negative formal charges should be on electronegative elements.
Examples l XeO 3 l NO 4 3- l SO 2 Cl 2
VSEPR l Lewis structures tell us how the atoms are connected to each other. l They dont tell us anything about shape. l The shape of a molecule can greatly affect its properties. l Valence Shell Electron Pair Repulsion Theory allows us to predict geometry
VSEPR l Molecules take a shape that puts electron pairs as far away from each other as possible. l Have to draw the Lewis structure to determine electron pairs. l bonding l nonbonding lone pair l Lone pair take more space. l Multiple bonds count as one pair.
VSEPR l The number of pairs determines –bond angles –underlying structure l The number of atoms determines –actual shape
VSEPR Electron pairs Bond Angles Underlying Shape 2180° Linear 3120° Trigonal Planar ° Tetrahedral 5 90° & 120° Trigonal Bipyramidal 6 90°Octagonal
Actual shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 220linear 330trigonal planar 321bent 440tetrahedral 431trigonal pyramidal 422bent
Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 550trigonal bipyrimidal 541See-saw 532T-shaped 523linear
Actual Shape Electron Pairs Bonding Pairs Non- Bonding Pairs Shape 660Octahedral 651Square Pyramidal 642Square Planar 633T-shaped 621linear
Examples l SiF 4 l SeF 4 l KrF 4 l BF 3 l PF 3 l BrF 3
No central atom l Can predict the geometry of each angle. l build it piece by piece.
How well does it work? l Does an outstanding job for such a simple model. l Predictions are almost always accurate. l Like all simple models, it has exceptions. l Doesnt deal with odd electrons
Polar molecules l Must have polar bonds l Must not be symmetrical l Symmetrical shapes include –Linear –Trigonal planar –Tetrahedral –Trigonal bipyrimidal –Octahedral –Square planar