Accelerated ChemistryChapter 6 NotesMr. Seidel

(Student edition)

Chapter 6 problem set: 2, 4, 6, 15, 19, 20, 21, 25, 27, 28, 34, 37, 39, 41, 42, 45, 47, 48, 49

Useful diagrams: You should check out all of the figures in this chapter. Look at the tables as well.

6.1 Introduction to chemical bonding

Most elements are not found alone in nature. They are “stuck” to other atoms.

Chemical Bond -

Types of chemical bonds:

Ionic -

Covalent -

Metallic -

Covalent bonds may be polar or nonpolar

Polar - unequal sharing of electrons (HCl)

Nonpolar - equal sharing of electrons (H2)

Book uses electronegativity differences to predict:

4.0 – 1.7 = 100 to 50% ionic character1.7-.3 = 50 to 5% i.c.>0.3 = less than 5% i.c.

so > 1.7 = ionic 1.7 -0 .4 = polar covalent0.3 or less = nonpolar

examples: NaClHClCl2

easier way to predict:

ionic = pc = npc =

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Why does chemical bonding occur? It usually results in going to a ______energy, more ______state.

Graphs of:

Forming bonds ( )Breaking bonds ( )

6.2 Covalent Bonding and Molecular Compounds

molecule -

monatomic molecules -

diatomic molecules -

polyatomic molecules -

chemical formulas show the relative #’s of atoms in a chemical compound

ex.C6H12O6

Pb(NO3)2

(NH4)2Cr2O7

The formation of a covalent bond:

Bond Length Vs. Bond Energy Graph

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Diatomic Molecules and Orbital Notation (orbital overlap diagrams):

H2F2

O2N2

Octet Rule- .

Electron cloud representations

F2HF

HF - orbital overlap diagram

Lewis Structures for covalent compounds

Basic rules (according to Seids) 1. each atom wants 8 electrons (except H wants 2)

2. each atoms goes for close to the right # of bonds

3. the least electronegative atoms goes in the middle or

4. the atom that makes the most bonds goes in the middle

Draw H2O, PCl3, SiH2F2, CS2, C2H6, C2H4, C2H2, CH2O, HCN, FON

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Drawing polyatomic ions

count electrons - if the charge is - 3, add 3 electrons

example PO4-3p-5,0 - 6x4=24, total = 29 but need to add 3 = 32

less bonds than atoms want = negative charge more bonds than atoms want = positive charge

CKbe - check bonds, check electrons

P wants 3 bonds, has 4 - + 1 chargeEach O wants 2, has 1 - so each O = -1 Total = - 3

DrawNH4+

coordinate covalent bond - 2 shared electrons in a bond are donated by 1 atom

Draw OH-1, sulfate, nitrate, nitrite, carbonate, bicarbonate, H2SO4, H3PO4

6.3 Ionic Bonding

Ionic compound -

- formula unit - lowest whole # ratio of ions

metals - lose electrons - why? ______

nonmetals - gain electrons - why?______

metals lose electrons until they become like a (usually 8 valence electrons)

nonmetals gain electrons until they do the same

both go to s2p6 - 8 valence electrons - called an

the tendency to arrange electrons so each atom has 8 is called the

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formation of an ionic bond: (Bohr model)

an easier way..... (Dot Pictures)

The ionic bonding picture looks like this....

Other examples:

energy is involved in all chemical reactions

Na + Cl yields NaCl + 769 kJ

lattice energy -

NaCl = - 769 kJ/moleNaF = - 922 kJ/moleKCl = -718 kJ/mole

smaller ions have LE

property summary of ionic compounds - hard, shatter (not ), conduct when ______, ______melting point, no smell

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6.4Metallic Bonding - “Sea of electrons theory” - Momma Seidel and her big, wooden spoon

Remember – in ionic bonds some atoms want e- and some don’t – in covalent bonds, all atoms share – in metals, no one atom wants the e-

Properties of metals - conduct electricity, heat, hardness, luster, malleability, ductility

6.5The Properties of Molecular Compounds

VSEPR Model - Valence shell electron pair repulsion theory (vesper) - electron pairs get as far away from each other as possible

consider shared pairs lone pairs shape bond angle(s)drawing

AB

AB2

AB2

AB2

AB3

AB3

AB4

AB5

AB6

to predict - structure must be drawn first - examples: CCl4, H2S, HBr, SO2, ClO4-1,BH3, PF5, etc.

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Intermolecular Forces- forces that hold molecules together

- happens in covalent compounds

intermolecular forces - can be weak or strong

intramolecular forces - always strong (bonds)

Types of IMF

1. dipole-dipole

dipole- when electrons are unevenly distributed

2. Hydrogen Bonding H-bonding is a “ ” dipole-dipole

Are we having FON? - H-bonding happens any time H is bonded to ______.

why? A large difference in electronegativity between F, O, or N and H results in one end of the molecule being very negative, while the other end is very positive.

Why N and not Cl? They have identical electronegativities!

Well, N is so much smaller than Cl so the negative charge is spread over a smaller area which exerts more force.

Effect of H-bonds on physical properties? Water expands when it freezes H-bonding is also responsible for the shapes of proteins

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we learned polar molecules have dipole-dipole IMF holding them together.

we learned about H-bonding or “super” dipole-dipole IMF

These two types of IMF usually result in substances being solids or liquids at room temp.

most nonpolar covalent substances are gases at room temp. as the forces holding them together are not strong enough to keep the molecules attracted - hence they are gases

O2, H2, N2 - straight nonpolar substances

CO2 - have dipoles, but nonpolar due to its molecular geometry

now, a third type of IMF

3. Van der Waals Forces (London Forces) - Temporarily induced dipoles caused by the motion of electrons.

more electrons = more attraction so, bigger atoms have stronger Van der Waals forces

NIB Seid’s Summary - see, chemistry makes sense after all!

IMFMolecule/Bond TypeMoleculeBoiling Pt. (Co)

Noble GasHe-269

Ar-186

nonpolarH2-253

O2-183

Cl2-34

polarHF19.5

ICl97

ionicNaCl1413

MgF22237

metallicCu2567

Fe2750

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NIB

Bond Energy - basic idea - what is the strength of chemical bonds?

bond energy - energy needed to break a bond - measured in kJ/mole

bond strength and stability:

stronger bond - more stable -needs more energy to break the bond

weaker bond - takes little energy to break the bond so the chemical is unstable

chemical changes favor lower energy states - exothermic reactions are favored

Bond Strength

which is stronger? - single, double, or triple bond? uh.....

which is shortest bond length? s, d, or t? - the answer is

which is stronger, short or long bonds? uh...

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