chemistry review atoms, stoichiometry, reactions and bonding

CHEMISTRY REVIEW

Atoms, Stoichiometry, Reactions and Bonding

MODERN VIEW• The atom is mostly

empty space.

• Two regions

• Nucleus- protons and neutrons.

• Electron cloud- region where you might find an electron.

SUB-ATOMIC PARTICLES

• Z - atomic number = number of protons determines type of atom.

• A - mass number = number of protons + neutrons.

• Number of protons = number of electrons if neutral.

SYMBOLS

XA

Z

Na23

11

CHEMICAL BONDS

• The forces that hold atoms together.

• Covalent bonding - sharing electrons.

• Makes molecules.

• Chemical formula- the number and type of atoms in a molecule.

• C2H6 - 2 carbon atoms, 6 hydrogen atoms,

• Structural formula shows the connections, but not necessarily the shape.

Quantum NumbersQuantum Numbers

Each electron in an atom has a unique set of 4 quantum numbers which describe it.

Principal quantum number Angular momentum quantum number Magnetic quantum number Spin quantum number

(n)(l)

(m)(s)

Assigning the NumbersAssigning the Numbers The three quantum numbers (n, l, and m) are integers. The principal quantum number (n) cannot be zero. n must be 1, 2, 3, etc. The angular momentum quantum number (l ) can be any integer between 0 and n - 1. For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1, +2.

Principle, angular momentum, and Principle, angular momentum, and magnetic quantum numbers: magnetic quantum numbers: nn, , ll, and , and mmll

Periodicity

TRENDS IN ATOMIC SIZE

• Influenced by three factors.• Energy Level• Higher energy level is further away.

•Charge on nucleus•More charge pulls electrons in closer.

• Shielding• Layers of electrons shield from nuclear pull.

SHIELDING• The electron on the outside

energy level has to look through all the other energy levels to see the nucleus.

• A second electron has the same shielding.

GROUP TRENDS

• As we go down a group• Each atom has

another energy level,• So the atoms get

bigger.

HLi

Na

K

Rb

PERIODIC TRENDS• As you go across a period the radius gets

smaller.

• Same energy level.

• More nuclear charge.

• Outermost electrons are closer.

Na Mg Al Si P S Cl Ar

IONIC SIZE

• Cations form by losing electrons.• Cations are smaller that the atom they

come from.• Metals form cations.• Cations of representative elements have

noble gas configuration.

IONIC SIZE

• Anions form by gaining electrons.

• Anions are bigger that the atom they come from.

• Nonmetals form anions.

• Anions of representative elements have noble gas configuration.

IONIZATION ENERGY

• The amount of energy required to completely remove an electron from a gaseous atom.

• Removing one electron makes a +1 ion.

• The energy required is called the first ionization energy.

IONIZATION ENERGY

• The second ionization energy is the energy required to remove the second electron.

• Always greater than first IE.

• The third IE is the energy required to remove a third electron.

Symbol First Second ThirdH

HeLiBeBCNOF Ne

1312 2731 520 900 800 1086 1402 1314 1681 2080

5247 7297 1757 2430 2352 2857 3391 3375 3963

11810 14840 3569 4619 4577 5301 6045 6276

WHAT DETERMINES IE

• The greater the nuclear charge the greater IE.

• Distance from nucleus increases IE

• Filled and half filled orbitals have lower energy, so achieving them is easier, lower IE.

• Shielding

GROUP TRENDS

•As you go down a group first IE decreases because•The electron is further away.•More shielding.

PERIODIC TRENDS

• All the atoms in the same period have the same energy level.

• Same shielding.

• Increasing nuclear charge

• So IE generally increases from left to right.

• Exceptions at full and 1/2 fill orbitals.

Firs

t Io

niz

ati

on

en

erg

y

Atomic number

H

He

Li

Be

B

C

N

O

F

Ne Na has a lower

IE than Li Both are s1

Na has more shielding

Greater distance

Na

Affinity tends to increase across a period

Affinity tends to decrease as you go down in a group

Electrons farther from the nucleus experience less nuclear attraction

Some irregularities due to repulsive forces in the relatively small p orbitals

Electron AffinityElectron Affinity - the energy change - the energy change associated with the addition of an electronassociated with the addition of an electron

ELECTRONEGATIVITY

• The tendency for an atom to attract electrons to itself when it is chemically combined with another element.

• How fair it shares.

• Big electronegativity means it pulls the electron toward it.

• Atoms with large negative electron affinity have larger electronegativity.

GROUP TREND

• The further down a group the farther the electron is away and the more electrons an atom has.

• More willing to share.

• Low electronegativity.

PERIODIC TREND

• Metals are at the left end.

• They let their electrons go easily

• Low electronegativity

• At the right end are the nonmetals.

• They want more electrons.

• Try to take them away.

• High electronegativity.

Ionization energy, electronegativityElectron affinity INCREASE

Atomic size increases, shielding constant

Ionic size increases

Prepare yourself toPrepare yourself to

C

c = C = speed of light, a constant (3.00 x 108 m/s)

= frequency, in units of hertz (hz, sec-1)

= wavelength, in meters

Electromagnetic radiation propagates Electromagnetic radiation propagates through space as a wave moving at through space as a wave moving at the speed of light.the speed of light.

Types of electromagnetic radiation:Types of electromagnetic radiation:

PROBLEM 37

E = h EE = Energy, in units of Joules (kg·m= Energy, in units of Joules (kg·m22/s/s22))

hh = Planck’s constant (6.626 x 10-34 J·s)= Planck’s constant (6.626 x 10-34 J·s)

= frequency, in units of hertz (hz, sec= frequency, in units of hertz (hz, sec-1-1))

The energy (The energy (E E ) of electromagnetic ) of electromagnetic radiation is directly proportional to radiation is directly proportional to the frequency (the frequency () of the radiation.) of the radiation.

Long Wavelength

= Low Frequency

= Low ENERGY

Short Wavelength

= High Frequency

= High ENERGY

Wavelength TableWavelength Table

Relating Frequency, Relating Frequency, Wavelength and EnergyWavelength and Energy

c hE

hc

E

Common re-arrangements:

E

hc

PROBLEM 43

PES• method that provides information on all the

occupied energy levels of an atom (that is, the ionization energies of all electrons in the atom) is known as photoelectron spectroscopy; this method uses a photon (a packet of light energy) to knock an electron out of an atom.

PHOTOELECTRON SPECTRUMThe photoelectron spectrum is a plot of the number of electrons emitted versus their kinetic energy. In the diagram below, the “X” axis is labeled high to low energies so that you think about the XY intersect as being the nucleus.

http://www.chem.arizona.edu/chemt/Flash/photoelectron.html



Interpretations from the data:1. There are no values on the y axis in the tables above. Using the Periodic Table and Table 1, put numbers on the y axis.2. Label each peak on the graphs above with s, p, d, or f to indicate the suborbital they represent..3. What is the total number of electrons in a neutral potassium atom?

Orbital names s, p, d, and f stand for names given to groups of lines in the spectra of the alkali metals. Early chemists called the line groups sharp, principal, diffuse, and fundamental.

1-2-

6-

1s 2s 3s4s

2p 3p

IONS

• Atoms or groups of atoms with a charge.

• Cations- positive ions - get by losing electrons(s).

• Anions- negative ions - get by gaining electron(s).

• Ionic bonding- held together by the opposite charges.

• Ionic solids are called salts.

POLYATOMIC IONS

• Groups of atoms that have a charge.

• Yes, you have to know common ones.

Periodic TablePeriodic Table

METALS• Conductors

• Lose electrons

• Malleable and ductile

NONMETALS• Brittle

• Gain electrons

• Covalent bonds

SEMI-METALS OR METALLOIDS

Alkali Metals

Alkaline Earth Metals

Halogens

Transition metals

Noble Gases

Inner Transition Metals

+1+2 -1-2-3

NAMING COMPOUNDS

• Two types

• Ionic - metal and non metal or polyatomics.

• Covalent- we will just learn the rules for 2 non-metals.

IONIC COMPOUNDS

• If the cation is monoatomic- Name the metal (cation) just write the name.

• If the cation is polyatomic- name it.

• If the anion is monoatomic- name it but change the ending to –ide.

• If the anion is poly atomic- just name it

• Practice.

COVALENT COMPOUNDS

• Two words, with prefixes.

• Prefixes tell you how many.

• mono, di, tri, tetra, penta, hexa, septa, octa, nona, deca

• First element whole name with the appropriate prefix, except mono.

• Second element, -ide ending with appropriate prefix.

• Practice

ACIDS

• Substances that produce H+ ions when dissolved in water.

• All acids begin with H.

• Two types of acids:

• Oxyacids

• Non-oxyacids

NAMING ACIDS

• If the formula has oxygen in it

• write the name of the anion, but change • ate to -ic acid• ite to -ous acid

• Watch out for sulfuric and sulfurous

• H2CrO4• HMnO4

• HNO2

NAMING ACIDS

• If the acid doesn’t have oxygen

• add the prefix hydro-

• change the suffix -ide to -ic acid

• HCl

• H2S

• HCN

HYDRATES

• Some salts trap water when they form crystals.

• These are hydrates.

• Both the name and the formula needs to indicate how many water molecules are trapped.

• In the name we add the word hydrate with a prefix that tells us how many water molecules.

HYDRATES

• In the formula you put a dot and then write the number of molecules.

• Calcium chloride dihydrate = CaCl22

• Chromium (III) nitrate hexahydrate = Cr(NO3)3 6H2O

THE MOLE

• The mole is a number.

• A very large number, but still, just a number.

• 6.022 x 1023 of anything is a mole

• A large dozen.

• The number of atoms in exactly 12 grams of carbon-12.

MOLAR MASS

• Mass of 1 mole of a substance.

• Often called molecular weight.

• To determine the molar mass of an element, look on the table.

• To determine the molar mass of a compound, add up the molar masses of the elements that make it up.

FIND THE MOLAR MASS OF

• CH4

• Mg3P2

• Ca(NO3)3

• Al2(Cr2O7)3

• CaSO4 · 2H2O

PERCENT COMPOSITION• Percent of each element a compound is composed

of.

• Find the mass of each element, divide by the total mass, multiply by a 100.

• Easiest if you use a mole of the compound.

• Find the percent composition of CH4

• Al2(Cr2O7)3

• CaSO4 · 2H2O

WORKING BACKWARDS

• From percent composition, you can determine the empirical formula.

• Empirical Formula the lowest ratio of atoms in a molecule.

• Based on mole ratios.

• A sample is 59.53% C, 5.38%H, 10.68%N, and 24.40%O what is its empirical formula.

• A 0.2000 gram sample of a compound (vitamin C) composed of only C, H, and O is burned completely with excess O2 . 0.2998 g of CO2 and 0.0819 g of

H2O are produced. What is the empirical formula?

EMPIRICAL TO MOLECULAR FORMULAS

• Empirical is lowest ratio.

• Molecular is actual molecule.

• Need Molar mass.

• Ratio of empirical to molar mass will tell you the molecular formula.

• Must be a whole number because...

EXAMPLE

• A compound is made of only sulfur and oxygen. It is 69.6% S by mass. Its molar mass is 184 g/mol. What is its formula?

CHEMICAL EQUATIONS

• Are sentences.

• Describe what happens in a chemical reaction.

• Reactants Products

• Equations should be balanced.

• Have the same number of each kind of atoms on both sides because ...

ABBREVIATIONS

• (s) • (g) • (aq)• heat

•

• catalyst

MEANING

• A balanced equation can be used to describe a reaction in molecules and atoms.

• Not grams.

• Chemical reactions happen molecules at a time

• or dozens of molecules at a time

• or moles of molecules.

STOICHIOMETRY

• Given an amount of either starting material or product, determining the other quantities.

• use conversion factors from• molar mass (g - mole)• balanced equation (mole - mole)

• keep track.

EXAMPLES• One way of producing O2(g) involves the

decomposition of potassium chlorate into potassium chloride and oxygen gas. A 25.5 g sample of Potassium chlorate is decomposed. How many moles of O2(g) are produced?

• How many grams of potassium chloride?

• How many grams of oxygen?

EXAMPLES• A piece of aluminum foil 5.11 in x 3.23 in x

0.0381 in is dissolved in excess HCl(aq). How many grams of H2(g) are produced?

• How many grams of each reactant are needed to produce 15 grams of iron form the following reaction? Fe2O3(s) + Al(s) Fe(s) + Al2O3(s)

YIELDHow much you get from an chemical reaction

LIMITING REAGENT

• Reactant that determines the amount of product formed.

• The one you run out of first.

• Makes the least product.

• Book shows you a ratio method.

• It works.

• So does mine

LIMITING REAGENT

• To determine the limiting reagent requires that you do two stoichiometry problems.

• Figure out how much product each reactant makes.

• The one that makes the least is the limiting reagent.

EXAMPLE

• Ammonia is produced by the following reaction

N2 + H2 NH3 What mass of ammonia can be produced from a mixture of 100. g N2 and 500. g

H2 ?

• How much unreacted material remains?

EXCESS REAGENT

• The reactant you don’t run out of.• The amount of stuff you make is the yield.• The theoretical yield is the amount you

would make if reaction went to completion.• The actual yield is what you make in the

lab.

PERCENT YIELD• % yield = Actual x 100%

Theoretical

• % yield = what you got x 100%what you could have got

TYPES OF REACTIONS

Precipitation reactions

• When aqueous solutions of ionic compounds are poured together a solid forms.

• A solid that forms from mixed solutions is a precipitate

• If you’re not a part of the solution, your part of the precipitate

PRECIPITATION REACTIONS

• NaOH(aq) + FeCl3(aq) NaCl(aq) + Fe(OH)3(s)

• is really

• Na+(aq)+OH-(aq) + Fe+3 + Cl-(aq) Na+

(aq) + Cl- (aq) + Fe(OH)3(s)

• So all that really happens is

• OH-(aq) + Fe+3 Fe(OH)3(s)

• Double replacement reaction

SOLUBILITY RULES

All nitrates are soluble

Alkali metals ions and NH4+ ions are

soluble

Halides are soluble except Ag+, Pb+2, and

Hg2+2

Most sulfates are soluble, except

Ag+,Pb+2, Hg+2 ,Sr+2, Ca+2 and Ba+2

SOLUBILITY RULES

Most hydroxides are slightly soluble (insoluble) except NaOH and KOH

Sulfides, carbonates, chromates, and phosphates are insoluble

Lower number rules supersede so Na2S is soluble

THREE TYPES OF EQUATIONS

• Molecular Equation- written as whole formulas, not the ions.

• K2CrO4(aq) + Ba(NO3)2(aq)

• Complete Ionic equation show dissolved electrolytes as the ions.

• 2K+ + CrO4-2 + Ba+2 + 2 NO3

-

BaCrO4(s) + 2K+ + 2 NO3-

• Spectator ions are those that don’t react.

THREE TYPE OF EQUATIONS

• Net Ionic equations show only those ions that react, not the spectator ions

• Ba+2 + CrO4-2 BaCrO4(s)

• Write the three types of equations for the reactions when these solutions are mixed.

• iron (III) sulfate and potassium sulfide Lead (II) nitrate and sulfuric acid.

TYPES OF REACTIONS

Acid-Base

• For our purposes an acid is a proton donor.

• a base is a proton acceptor usually OH-

• What is the net ionic equation for the reaction of HCl(aq) and KOH(aq)?

• Acid + Base salt + water

• H+ + OH- H2O

ACID - BASE REACTIONS

• Often called a neutralization reaction Because the acid neutralizes the base.

• Often titrate to determine concentrations.

• Solution of known concentration (titrant),

• is added to the unknown (analyte),

• until the equivalence point is reached where enough titrant has been added to neutralize it.

TITRATION

• Where the indicator changes color is the endpoint.

• Not always at the equivalence point.

• A 50.00 mL sample of aqueous Ca(OH)2 requires 34.66 mL of 0.0980 M Nitric acid for neutralization.

What is [Ca(OH)2 ]? • # of H+ x MA x VA = # of OH- x MB x VB

ACID-BASE REACTION

• 75 mL of 0.25M HCl is mixed with 225 mL of 0.055 M Ba(OH)2 . What is the concentration of the excess

H+ or OH- ?

TYPES OF REACTION

Oxidation-Reduction called Redox

• Ionic compounds are formed through the transfer of electrons.

• An Oxidation-reduction reaction involves the transfer of electrons.

• We need a way of keeping track.

OXIDATION STATES

• A way of keeping track of the electrons.

• Not necessarily true of what is in nature, but it works.

• need the rules for assigning (memorize). The oxidation state of elements in their standard

states is zero. Oxidation state for monoatomic ions are the same as

their charge.

OXIDATION STATES Oxygen is assigned an oxidation state of -2 in its

covalent compounds except as a peroxide. In compounds with nonmetals hydrogen is assigned

the oxidation state +1. In its compounds fluorine is always –1. The sum of the oxidation states must be zero in

compounds or equal the charge of the ion.

OXIDATION STATES

• Assign the oxidation states to each element in the following.

• CO2

• NO3-

• H2SO4

• Fe2O3

• Fe3O4

OXIDATION-REDUCTION• Transfer electrons, so the oxidation states change.

• Na + 2Cl2 2NaCl

• CH4 + 2O2 CO2 + 2H2O

• Oxidation is the loss of electrons.

• Reduction is the gain of electrons.

• OIL RIG

• LEO GER

HALF-REACTIONS• All redox reactions can be thought of as happening in

two halves.

• One produces electrons - Oxidation half.

• The other requires electrons - Reduction half.

• Write the half reactions for the following.

• Na + Cl2 Na+ + Cl-

• SO3- + H+ + MnO4

- SO4- + H2O + Mn+2

BALANCING REDOX EQUATIONS

• In aqueous solutions the key is the number of electrons produced must be the same as those required.

• For reactions in acidic solution an 8 step procedure. Write separate half reactions For each half reaction balance all reactants except H

and O

Balance O using H2O

ACIDIC SOLUTION

Balance H using H+

Balance charge using e-

Multiply equations to make electrons equal Add equations and cancel identical species Check that charges and elements are balanced.

PRACTICE• The following reactions occur in aqueous solution.

Balance them

• Cr(OH)3 + OCl- + OH

- CrO4

-2 + Cl

- +

H2O

• MnO4- + Fe+2 Mn+2 + Fe+3

• Cu + NO3- Cu+2 + NO(g)

• Pb + PbO2 + SO4-2 PbSO4

• Mn+2 + NaBiO3 Bi+3 + MnO4-

REDOX TITRATIONS• Same as any other titration.

• the permanganate ion is used often because it is its

own indicator. MnO4- is purple, Mn+2 is colorless.

When reaction solution remains clear, MnO4- is

gone.

• Chromate ion is also useful, but color change, orangish yellow to green, is harder to detect.

EXAMPLE• The iron content of iron ore can be determined by

titration with standard KMnO4 solution. The iron ore is

dissolved in excess HCl, and the iron reduced to Fe+2 ions. This solution is then titrated with KMnO4 solution,

producing Fe+3 and Mn+2 ions in acidic solution. If it requires 41.95 mL of 0.205 M KMnO4 to titrate a solution made with 0.6128 g of iron ore, what percent of the ore was iron?

CHAPTER 8

Bonding

WHAT IS A BOND?

• A force that holds atoms together.

• We will look at it in terms of energy.• Bond energy - the energy required to break a bond.

• Why are compounds formed?• Because it gives the system the lowest energy.

IONIC BONDING

• An atom with a low ionization energy reacts with an atom with high electron affinity.

• The electron moves.

• Opposite charges hold the atoms together.

COULOMB'S LAW

• E= 2.31 x 10-19 J · nm(Q1Q2)/r• Q is the charge.• r is the distance between the centers.

• If charges are opposite, E is negative• exothermic

• Same charge, positive E, requires energy to bring them together.• endothermic

WHAT ABOUT COVALENT COMPOUNDS?

• The electrons in each atom are attracted to the nucleus of the other.

• The electrons repel each other,

• The nuclei repel each other.

• They reach a distance with the lowest possible energy.

• The distance between is the bond length.

0

En

ergy

Internuclear Distance

0

En

ergy


Bond Length

0

En

ergy


Bond Energy

Bond Length DiagramBond Length Diagram

COVALENT BONDING

• Electrons are shared by atoms.

• There are two extremes.

• In between are polar covalent bonds.• The electrons are not shared evenly.

• One end is slightly positive, the other negative.• Indicated using small delta

Electronegativity:Electronegativity:

Ionic BondsIonic Bonds Electrons are transferred

Electronegativity differences are generally greater than 1.7 The formation of ionic bonds is always exothermic!

DIPOLE MOMENTS

• A molecule with a center of negative charge and a center of positive charge is dipolar (two poles),• or has a dipole moment.

• Center of charge doesn’t have to be on an atom.

• Will line up in the presence of an electric field.

SH2.1 2.5

H HS

+

-

H HS

+

+

- HH

S +

+-H

HS+

+ -H

H S++

-

H HS

+

+

-

HH

S

+

+

-

HH

S

++

-

HH

S

+

+

-

HH

S

++

-

HH

S

+

+

-

HH

S

+

HOW IT IS DRAWN

+

-

HH

S

+

WHICH MOLECULES HAVE DIPOLE MOMENTS?

• Any two atom molecule with a polar bond.

• With three or more atoms there are two considerations.• There must be a polar bond.• Geometry can’t cancel it out.

GEOMETRY AND POLARITY• Three shapes will cancel them out.

• Linear

GEOMETRY AND POLARITY

• Three shapes will cancel them out.

• Planar triangles

120º

GEOMETRY AND POLARITY• Three shapes will cancel them out.

• Tetrahedral

GEOMETRY AND POLARITY

• Others don’t cancel

• Bent

GEOMETRY AND POLARITY• Others don’t cancel

• Trigonal Pyramidal

IONS• Atoms tend to react to form noble gas configuration.

• Metals lose electrons to form cations

• Nonmetals can share electrons in covalent bonds.

• When two non metals react.(more later)

• Or they can gain electrons to form anions.

IONIC COMPOUNDS• We mean the solid crystal.

• Ions align themselves to maximize attractions between opposite charges,

• and to minimize repulsion between like ions.

• Can stabilize ions that would be unstable as a gas.

• React to achieve noble gas configuration

PERIODIC TRENDS• Across the period nuclear charge increases so they get

smaller.

• Energy level changes between anions and cations.

Li+1

Be+2

B+3

C+4

N-3O-2 F-1

SIZE OF ISOELECTRONIC IONS

• Iso - same

• Iso electronic ions have the same # of electrons

• Al+3 Mg+2 Na+1 Ne F-1 O-2 and N-3

• All have 10 electrons.

• All have the configuration 1s22s22p6

SIZE OF ISOELECTRONIC IONS• Positive ions have more protons so they are smaller.

Al+3

Mg+2

Na+1 Ne F-1 O-2 N-3

USING BOND ENERGIES• We can find H for a reaction.

• It takes energy to break bonds, and end up with atoms (+).

• 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.

• Energy and Enthalpy are state functions.

FIND THE ENERGY FOR THIS

2 CH2 = CHCH3

+

2NH3 O2+

2 CH2 = CHC N

+

6 H2O

C-H 413 kJ/molC=C 614kJ/molN-H 391 kJ/mol

O-H 467 kJ/molO=O 495 kJ/molCN 891 kJ/mol

C-C 347 kJ/mol

LEWIS STRUCTURE• Shows how the valence electrons are arranged.

• One dot for each valence electron.

• A stable compound has all its atoms with a noble gas configuration.

• Hydrogen follows the duet rule.

• The rest follow the octet rule.

• Bonding pair is the one between the symbols.

RULES• Sum the valence electrons.

• Use a pair to form a bond between each pair of atoms.

• Arrange the rest to fulfill the octet rule (except for H and the duet).

• H2O

• A line can be used instead of a pair.

• POCl3 P is central atom

• SO4-2

S is central atom

• SO3-2

S is central atom

• PO4-2

P is central atom

• SCl2 S is central atom

EXCEPTIONS TO THE OCTET• BH3 • Be and B often do not achieve octet• Have less than and octet, for electron deficient

molecules.

• SF6 • Third row and larger elements can exceed the octet• Use 3d orbitals?

• I3-

EXCEPTIONS TO THE OCTET

• When we must exceed the octet, extra electrons go on central atom.

• ClF3

• XeO3

• ICl4-

• BeCl2

RESONANCE• Sometimes there is more than one valid

structure for an molecule or ion.• NO3

-

• Use double arrows to indicate it is the “average” of the structures.

• It doesn’t switch between them.• NO2

-

• Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.

FORMAL CHARGE• The difference between the number of valence

electrons on the free atom and that assigned in the molecule.

• We count half the electrons in each bond as “belonging” to the atom.

• SO4-2

• Molecules try to achieve as low a formal charge as possible.

• Negative formal charges should be on electronegative elements.

EXAMPLES

• XeO3

• NO4-3

• SO2Cl2

VSEPR

• Lewis structures tell us how the atoms are connected to each other.

• They don’t tell us anything about shape.

• The shape of a molecule can greatly affect its properties.

• Valence Shell Electron Pair Repulsion Theory allows us to predict geometry

VSEPR• Molecules take a shape that puts electron pairs as far

away from each other as possible.

• Have to draw the Lewis structure to determine electron pairs.

• bonding

• nonbonding lone pair

• Lone pair take more space.

• Multiple bonds count as one pair.

VSEPR

• The number of pairs determines• bond angles• underlying structure

• The number of atoms determines • actual shape

VSEPRElectron

pairsBond

AnglesUnderlyingShape

2 180° Linear

3 120° Trigonal Planar

4 109.5° Tetrahedral

590° &120°

Trigonal Bipyramidal

6 90° Octagonal

A

C

C

B

SP HYBRIDIZATION• End up with two lobes 180º apart.

• p orbitals are at right angles

• Makes room for two bonds and two sigma bonds.

• A triple bond or two double bonds.

SP2 HYBRIDIZATION

• When three things come off atom.

• trigonal planar

• 120º

• on one bond

SP2 HYBRIDIZATION

• C2H4

• Double bond acts as one region of electrons.

• trigonal planar

• Have to end up with three blended orbitals.

• Use one s and two p orbitals to make sp2 orbitals.

• Leaves one p orbital perpendicular.

SP3 HYBRIDIZATION

• We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry.

• We combine one s orbital and 3 p orbitals.

• sp3 hybridization has tetrahedral geometry.

HOW WE GET TO HYBRIDIZATION

• We know the geometry from experimentation• We know the orbitals of the atom• hybridizing atomic orbitals can explain

the geometry.• So if the geometry requires a

tetrahedral shape, it is sp3 hybridized• This includes bent and trigonal

pyramidal molecules because one of the sp3 lobes holds the lone pair.

SP3D HYBRIDIZATION

• The model predicts that we must use the d orbitals.

• sp3d hybridization, ex: PCl5

•Trigonal bipyrimidal•can only bond.•can’t bond.•basic shape for five bonds.

SP3D2 • gets us to six bonds around

• octahedral

TWO TYPES OF BONDS

• Sigma bonds from overlap of orbitals.

• Between the atoms.

• Pi bond ( bond) above and below atoms

• Between adjacent p orbitals.

• The two bonds of a double bond.

CO2• C can make two and two • O can make one and one

CCOO OO

Let’s Try Some ProblemsLet’s Try Some Problems

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chemistry review atoms, stoichiometry, reactions and bonding

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