materials held together by the simultaneous attraction of electrons to two nuclei
What electrons?
How attracted? 2 ways!
Bonding
Chemical Energy -a form of potential energy (PE) PE- energy stored in
molecules Substances possess
energy (chemical) due to their composition and structure
Energy Changes in Bonding
When a chemical bond is formed, energy is released.
When a chemical bond is broken, energy is absorbed.
When 2 atoms are held together by a chemical bond, generally at a lower energy condition than when separated.
Bonding and Stability Generally, systems at
low energy levels are more stable than at high energy levels.
So bonding will more often occur if change leads to lower energy condition (MORE STABLE).
Bond is formed -more energy released means more stable, stronger bond
-less energy given off means less stable weaker bond
Is Oxygen complete? Forms bond, releases energy
exothermic reaction Later, breaks bond and
decomposes endothermic rxn (reverse of exothermic)
Exact same amount of energy required for reaction either released or absorbed (3351 kiloJoules)
2Al2O3(s) 4Al(s) + 3O2(g) +3351kJ
4Al(s) + 3O2(g) 2Al2O3(s) -3351kJ
*from Ref Table I
Electronegativity Measure of the ability
of an atom to attract the electrons that form a bond between it and another atom.
Highest electronegativity is 4.0
Reference Table S
Can use electronegativities to predict type of bonds formed.
Bonds Between Atoms Electrons (valence)
involved in bond formation can be transferred, shared equally or shared unequally between 2 atoms.
Ionic Bonds aka Electrovalent/Electrostatic Transfer of one or more electrons from metals to nonmetals form ions ions attract (+ and -) {electrostatic force} ionic bond formed
Ionic Bonds continued Form between
elements with electronegativity difference 1.7 with few exceptions.
*Remember metals with nonmetals.
Ionic bonds may form between monatomic or polyatomic ions.
Monatomic ions one atom with charge Na+, F-, Al+3
Polyatomic ions compound of 2 or
more covalently bonded atoms with a charge.
OH-, NH4+, S2O3
-2
Reference Table E.
AgNO3(aq) + NaCl(aq) --> AgCl(aq) + NaNO3(aq)
What is polyatomic ion in above reaction?
What happens to compounds above in water?
Lattice (Binding) Energyof an Ionic Solid
Measure of the energy required to completely separate a mole of a solid ionic compound into its separate ions.
The higher the lattice energy, the stronger the ionic bond.
Taken from mikeblaber.org on 7/27/11.
Ionic Solids Structural unit: made up of ions. Ionic bonding High melting points-->strong forces Geometric structure, ions held in crystal lattice by
electrostatic attraction Do not conduct electricity. When melted or dissolved in water, crystal lattice
is destroyed and ions move freely allowing for electrical conductivity.
Brittle Ex: NaCl, KClO3,MgO, KBr, Li2SO4
Covalent Bonds Simultaneous attraction of 2 nuclei for the same electrons resulting in the sharing of those electrons.
Difference in electronegativities is less than 1.7(some exceptions).
Bonding Continuum
Building Molecules with Lewis Dot Structures
Nonpolar Covalent Bonds Electrons shared
equally between atoms of the same element.
Ex: Diatomics “identical twins”
H2,N2,O2,F2,Cl2,Br2,I2
Difference in electronegativities is ZERO.
Diatomics can have single, double or triple covalent bonds.
Structures are symmetrical/nonpolar.Building Molecules with Lewis Dot Structures
Sigma (s) and Pi (p) Bonds•Single bonds are sigma bonds, electron density is concentrated along the line that represents the bond joining the two atoms.(overlapping s orbitals)•Double bonds contain one sigma and one pi bond. A pi bond occurs when the electron density is concentrated above and below the line that represents the bond joining the two atoms. (overlapping p orbitals)•Triple bonds are one sigma and two pi bonds.
Polar Covalent bonds can help create both polar(asymmetrical) and non-polar(symmetrical) structures.
Remember that the words polar and nonpolar can be used to describe both bonds and overall symmetry of molecules!
Polar Molecules due to either bonding or symmetry
Types of Molecular Shape that influence Overall Symmetry
Molecule Bonding Diagram Shape Symmetry of Molecule
Water polar bentasymmetrical
(polar)
Carbon dioxide polar linear
symmetrical (non-polar)
Methane polar tetrahedralsymmetrical (non-polar)
3D image
3D image
Resonance Structures A hybrid of the possible drawings because
no one Lewis structure can represent the situation.
Taken from sv.wikipedia.org on 7/27/11.
Taken from en.wikipedia.org on 7/27/11.
The VSPER Model-molecule will assume the shape that most minimizes electron pair repulsions
Total number of single bonds, double bonds, and lone pairs on the central atom
Structural pair geometry Shape
2 Linear
3 Trigonal Planar
4 Tetrahedral
5 Trigonal Bipyramidal
6 Octahedral
planar
VSPER is Valence Shell Electron Pair Repulsion.
Coordinate Covalent Bonds When 2 shared electrons forming covalent bond are both
donated by one of the atoms “Free loader” Once formed, same as ordinary covalent bond Often involved with forming polyatomic ions
Bonding by Jarod Gagnon
ElectronegativityIt is the ability
of an atom to takesome electrons to make
it’s outer shell stableBut how is it able?
What is electronegatvity for?Sometimes the difference is 1.7 or more
That means one atom is so much strongerit takes the other’s electrons who has them no longer.
If the number is lower than thatthe electrons just share and sit where they sat
Then a covalent bond has begun,but that isn’t the end of the fun.
Sometimes a freeloader comes to the tableand shares some electrons so it can be stable.
The previous compound had no electrons to lack,which is good for the freeloader gives nothing back.
Covalent Bonds form molecules. Molecule-discrete particle formed by covalently
bonded atoms where atoms share electrons so that final electron configuration of each atom is similar to an inert gas
Examples: O2, HCl, H2O, CH4, NH3, C6H12O6, CO2
Molecules form molecular substances. Molecular substances may be gases, liquids or
solids depending on attraction that exists between the molecules.
Just as we did for Ionic Bonding, let’s compare properties & characteristics of Molecular substances as solids!
Covalent or Molecular Solids Structural unit: made
of molecules (covalent bonds)
Low melting points Relatively weak forces Soft Poor heat conductors Do not conduct
electricity (good electrical insulators)
Ex: I2, H2O, CO2, cellulose C5H10O5
Predicting Bonds and Structure?
Building Molecules with Lewis Dot Structures
Network Bonding Certain solids consist of covalently bonded atoms
linked in a network that extends throughout sample with an absence of discrete particles
“One Big Giant” Molecule
Just as we did for Ionic & Covalent bonding, let’s compare properties & characteristics of Network solids!
Network Solids Structural unit: made
up of atoms, “One Giant Molecule”
Very high melting pts. Very strong covalent
bonds (Network bonds) Do not conduct in any
phase Very hard Ex: Diamond (C),
Quartz/Sand (SiO2), Silicon Carbide (SiC)
Silicon Carbide grinding wheels
Metallic Bonding Occurs between atoms that have a small number of valence electrons (metals) leaving them with many vacant valence orbitals and low ionization energies
“Electron Sea Model”
Electron Sea Model Held together by positive kernel and negative
valence electrons. Do you remember kernel? Valence e- free to move from atom to atom. How
can this explain a metal’s conductivity, ductility and malleability?
Just as we did for Ionic, Covalent & Network bonding, let’s compare properties & characteristics of Metallic solids!
Metallic Solids Structural units: made of
positive kernels and valence electrons.
Intermediate melting pts. Relatively intermediate
forces (metallic bonds) Conduct electricity and
heat in all phases. Malleability, Ductility
and Luster. Ex: Cu, Na, Fe, K, Au
Can you complete a summary chart that compares characteristics of all four types of solids?
4 Types of Solids: Ionic, Molecular, Network & Metallic solids
At least 6 properties/characteristics:
type of bonding structural unit melting point conductivity examples other
So far we have talked about bonds between atoms inside molecules or particles! Intramolecular Forces:
found within molecule or particle called chemical bonds.
Think intramurals or intravenous!
6 examples: Ionic Polar Covalent Non-polar covalent Coordinate covalent Network Metallic
Molecular Attractions aka Intermolecular Forces
Bonds (Intramolecular forces) build particles. Now how do these particles come together to build
something we can actually see?
forces beween moleculesThink international or interscholastic
Dipoles Dipoles are polar
molecules “2 poles.” Have asymmetrical
distribution of electrical charge within molecule.
Dipole-dipole attraction is force of attraction between polar molecules.
Hydrogen Bonding Special case of dipole-
dipole that occurs when hydrogen is bonded to a small, highly electronegative atom (N, O, F)
Slightly stronger than other dipole-dipole attraction.
Hydrogen bonding is FON!!! Or I’ve had NOF of Hydrogen bonding!!!
Why doesn’t H2O fit with pattern of others? All are polar, same family, all have same bonding. Smaller mass, should be lower boiling pt.? H2O has hydrogen bonding.
London Dispersion Forcesaka Van Der Waals
Weak attractive forces exist between non-polar molecules (no dipoles, no H bonding)
Caused by momentary dipole, a chance distribution of electrons.
Make it possible for small, non-polar molecules to exist as solids or liquids under low temp and high pressure ex. H2, He, O2, N2
London Dispersion Forces act over short distances. Can these forces be
increased? London Dispersion
forces increase with increase in size of molecules and # of electrons or a decrease in distances between molecules.
Pentane
Molecule-Ion Attraction Ionic compounds
(salts) are generally soluble in polar solvents.
Remember (aq)! Why soluble?
Molecule-Ion Attraction Explained Polar solvents are asymmetrical (+ & - ends). + & - end of liquid are attracted to + & - ends of
ionic salt. This pulls apart ions, breaking crystal lattice structure (salt is dissolved and broken!)
If water is solvent, creates hydrated ions: water molecule surrounded by ions.
Now that ions are broken free of lattice, what can the solution now do?
Chemical Formulas Both a qualitative and
a quantitative expression of the composition of an element or compound.
How much sulfuric acid is represented?
One H2SO4 molecule containing 2 H atoms, 1 S atom & 4 O atoms OR One mole of H2SO4 molecules containing 2 moles of H atoms, 1 mole of S atoms and 4 moles of O atoms.
2 Types of Chemical Formulas Molecular Formula:
indicates total number of atoms of each element needed to form a molecule.
Empirical Formula: simplest ratio in which atoms combine to form a compound
Empirical formulas do not always exist in nature.
For the next section in this unit, you will need to do the following: Write chemical
formulas. Name chemical
compounds. Balance chemical
equations.
Writing Formulas for Ionic Compounds If you use the criss-cross method, remember to
simplify to lowest values. (Use parentheses if multiple polyatomic ions are needed).
1. Write the symbol for the metal ion.2. Write the symbol for the nonmetal or
polyatomic ion.3. Check the oxidation numbers of each ion. If they
add up to zero, this is the formula. A Roman numeral after the name of the metal ion
denotes its oxidation number.
Writing Formulas for Ionic Compounds continued
4. Use the proper subscripts after the symbol for each ion so that when multiplied times the oxidation #, the total algebraic sum is zero. (Use parentheses if multiple polyatomic ions are needed).
Examples: calcium carbonate, ammonium sulfite, Nickel (III) sulfide, Copper (II) chloride
Naming Ionic Compounds Text p.176
1. Write the name of the cation (metal or ammonium ion).
2. Write the name of the anion (nonmetal or polyatomic ion). Nonmetals ending to ide.
3. If the metal can have more than one charge (oxidation number) place a Roman Numeral after its name to denote the charge.
4. Examples: FeBr3, K2Cr2O7, Mn(C2H3O2)3
Naming Covalent Compounds Text pp.206-207
For covalent compounds composed of two elements, name 1st element and then 2nd element’s name ending is changed to –ide.
Use prefixes to indicate number of atoms of each element.
Examples: CO, CO2, CCl4, P2S5,
1 = mono- 2 = di- 3 = tri- 4 = tetra- 5 = penta-
6 = hexa- 7 = hepta- 8 = octa- 9 = nona- 10 = deca-
Balancing Chemical EquationsText pp.267-274
Equations must be balanced to support the Law of Conservation of Mass (matter or mass can not be created nor destroyed, only rearranged).
To balance, you need to make # of atoms of each element the same on both reactants’ and products’ side.
Balancing Chemical Equations Continued
Can’t change the formula, only the coefficients. Coefficient- a small whole number that appears as
a factor in front of a formula in a chemical rxn. Hint: When balancing, start with uncommon
elements first. If polyatomic ions appear on both reactant and product side, balance as a group. Often it is helpful to save H and O until the end.
When finished, your answer should be simplified into smallest whole #’s.
Let’s try some examples.