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BondingTopics 4 and

14IB Chemistry Year 1 HL

The Brooklyn Latin School

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REVIEW: Ionic Bonding

Wednesday, December 3, 2014

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4.1 Ionic Bonding & Structure

Understandings• Positive ions (cations) form by metals losing valence

electrons.• Negative ions (anions) form by non-metals gaining

electrons.• The number of electrons lost or gained is determined

by the electron configuration of the atom. • The ionic bond is due to electrostatic attraction

between oppositely charged ions.• Under normal conditions, ionic compounds are usually

solids with lattice structures.

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4.1 Ionic Bonding & Structure

Applications• Deduction of the formula and name of an ionic

compound from its component ions, including polyatomic ions.

Guidance• Students should be familiar with the names of these

polyatomic ions: NH4+, OH–, NO3–, HCO3–, CO32–, SO42–, and PO43–

• Explanation of the physical properties of ionic compounds (volatility, electrical conductivity, and solubility) in terms of their structure.

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Valence Electrons• Valence electrons are the electrons in the outermost

principal energy level• All atoms want to be like the Noble Gases, in Group

18, and have a filled s and p sublevel in the valence shell

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Ions• Atoms can gain or lose valence electrons to get a stable

octet.• Elements that have a small number of electrons in their

outer shells (Groups 1, 2, and 13) will lose those electrons and form positive ions called cations. These elements are the metals.

• Elements that have higher numbers of electrons in their outer shells (Groups 15, 16, and 17) will gain electrons and form negative ions called anions. These elements are the non-metals.

• Group 14 elements do not tend to form ions.

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Sneak Peak: Redox• Formation of ionic compounds are always redox

reactions• Oxidation is losing electrons• Reduction is gaining electrons• LEO the Lion goes GER

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Families• The group or family on the Periodic Table helps you

predict what kind of ions will be formed (unlike last year, you will not be given this info!)

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Let’s Practice• What kind of ion is formed by the following elements?1. Lithium2. Sulfur3. Argon4. Oxygen5. Nitrogen

NOTE: When writing an ion, always put the charge number first as the superscript and then the + or -. Example: Na3+ NOT Na+3. This is important for IB!!

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Transition Metals• Remember from Unit 2 Periodic Trends that the

transition metals can form more than one type of ion (lose from the s and d sublevels)

• Refer back to what we learned in Unit 2!

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Unusual Ions• Lead, Pb, despite being in Group 14, forms a stable ion

Pb2+

• Tin, Sn, also in Group 14, can form Sn4+ and Sn2+

• Silver, Ag, forms the ion Ag+• Hydrogen, H, can form H– (hydride) as well as the more

common H+.

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Naming Ions• Cations, or positive ions, keep their name. If an element

can form more than one positive ion, a Roman numeral is written to indicate the chargeo E.g Iron +3 is Iron (III) where Iron +2 is Iron (II)

• Anions, or negative ions, get their ending changed to –ideo E.g. Oxygen -2 is oxide and Nitrogen -3 is nitride

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Polyatomic Ions• Polyatomic ions are groups of atoms that are covalently

bound together (i.e. sharing electrons) with an overall change.

• When naming them, you simply say their name.• You do have to memorize the names and charges of

several polyatomic ions!!!

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Memorize!

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Ionic Compounds• Ionic compounds are formed when there is a transfer

of electrons from metals to nonmetals (or to a group of atoms in the case of polyatomic ions)

• The resulting positive and negative ions are attracted to each other via electrostatic attraction

• The ions now have the electron configuration of Noble gases

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Writing Ionic Compound Formulas

• Since electrons are transferred from one set of atoms to another, the overall ionic compound must remain electrically neutral

• When we are determining the formula for ionic compounds, we simply have to balance to positive and negative charges

• Last year, we called this the kriss kross method!

Writing Ionic Formulas1. Check the periodic table for the ions that each element will form.

Al is in group 3, Oxygen is in group 6

2. Write the number of the charge above the ion. 3+ 2-Al O

3. “Criss cross” the numbers to balance the charges.3 2Al O

4. Write the final formula using subscripts to show the number of each ion.

Al2O3

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Let’s Practice• Write the formula for the compound that forms between

magnesium and nitrogen. • Step 1: Determine the charge each ion forms

Mg +2 N -3• Step 2: Write the cation first and the anionl second• Step 3: Using subscripts, put the number of ions of

each element that allows the charges to balance out; overall charge should equal ZERO!

Mg +2 N -3

Mg3N2

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Reminder• Ionic formulas are ALWAYS empirical because ionic

compounds form crystal structures of repeating ions; this simplest ratio of ions is called the formula unit which simply states the ratio of ions in the crystal

• When using the kriss kross method, only bring down the absolute value; there are no signs in the subscripts

• If polyatomic ions are in the compound, put them in parenthesis and put the subscript outside of that; do not need if the subscript is 1 (one)

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Let’s Practice• Write the formula for ammonium phosphate.

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Lies, oversimplifications and more…

• In Regents Chemistry, we sometimes oversimplified complex concepts to help students understand what was going on; now you need to know the WHY.

• When we are going over bonding now, I am going to try to rid you of the technically incorrect, oversimplified

models you may have from last year.

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The Truth is Out There• Ionic bonding and covalent bonding are not “two

different types of bonds.” • Sodium atoms do not “want to lose an electron.” In fact,

ionization is always endothermic, meaning it actually takes energy to remove the electron from sodium. The fact is, chlorine wants that electron more and takes it from sodium resulting in two stable electron configurations. Energy is release when chlorine takes that electron AND energy is released when the resulting positive ions form the crystal lattice structure. Reactions are all about how much energy you put in and how much you get out!

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More LiesMost textbooks will do the following which ignores the truth:• Showing how sodium atoms give an electron to a

chlorine atom when the reaction is actually between sodium metal and gaseous chlorine molecules, Cl2 (see the image on the right showing sodium metal burning in chlorine gas).

• Showing atoms of sodium and chlorine to both be the same size, i.e. with identical atomic radii.

• Showing a chloride ion the same size as a chlorine atom, and in some cases the same size as a sodium ion.

• Showing all the energy levels in both the sodium and chlorine atoms to be equally spaced.

• Showing the electrons separately in the first energy level and in pairs in the subsequent energy levels.

• Ignoring the changes of state that take place during the reaction.

Source: ThinkIB.net

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Typical Way Ionic Bonding is Shown

What is wrong here?

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Let’s Practice1  Write the formula for each of the compounds in the table on page 3 of your notes with polyatomic ions. 2  Write the formula for each of the following compounds:

(a)  potassium bromide (b)  zinc oxide (c)  sodium sulfate (d) copper(II) bromide (e) chromium(III) sulfate (f) aluminium hydride

3 Name the following compounds: (a) Sn3(PO4)2 (b) Ti(SO4)2 (c) Mn(HCO3)2

(d) BaSO4 (e) Hg2S

4  What are the charges on the positive ions in each of the compounds in Q3 above? 5  What is the formula of the compound that forms from element A in Group 2 and element B in Group 15? 6  Explain what happens to the electron configurations of Mg and Br when they react to form the compound magnesium bromide.

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Lesson 2: Advanced

Ionic BondingThursday, December 4th, 2014

Ionic Bonding Ionic bonds result from the

attractions between positive and negative ions.

Ionic bonding involves 3 aspects:1. Loss of an electron(s) by one element. 2. Gain of electron(s) by a second element.3. Attraction between positive and negative

ions.

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Stable Octet Rule• Atoms tend to either gain or lose

electrons in their highest energy level to form ions

• Atoms prefer having 8 electrons in their highest energy level

Na atom 1s2 2s2 2p6 3s1 One electron extraCl atom 1s2 2s2 2p6 3s2 3p5 One electron short of a stable octet

Na+ Ion 1s2 2s2 2p6 Stable octetCl- Ion 1s2 2s2 2p6 3s2 3p6 Stable octet

Examples

Positive ions attract negative ions forming ionic bonds.

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Ionic Bonding Ionic substances are made of repeating arrays of

positive and negative ions.

An ionic crystal lattice

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Ionic Bonding The array is repeated over and over to form the

crystal lattice.

Each Na+ ion is surrounded by 6 other Cl- ions. Each Cl- ion is surroundedby 6 other Na+ ions

Model of aSodium chloridecrystal

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Ionic Compound Structure

• In Chemistry, opposites attract, meaning the ions want to surround themselves with other ions of opposite charge

• The ions take on a predictable three-dimensional crystalline structure known as the ionic lattice (see below)

• The coordination number of the lattice tells you how many ions each ion in the crystal is surrounded byFor example, in the sodium

chloride lattice, the coordination

number is six because each Na+

ion is surrounded by six Cl– ions

and each Cl– ion is surrounded by

six Na+ ions.

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Lattice Energy• Lattice energy is a measure of the strength of

attraction between ions in the lattice of an ionic compound

• Lattice energy is higher for ions that are small and highly charged and weaker for ions that are larger and have a lower charge

NaCl

Ionic Bonding• The shape and form of the crystal lattice

depend on several factors:

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1. The size of the ions2. The charges of the

ions3. The relative numbers

of positive and negative ions

Characteristics of Ionic Bonds

1. Crystalline at room temperatures

2. Higher melting points and boiling points than covalent compounds

3. Conduct electrical current in molten or solution state but not in the solid state

4. More soluble in polar solvents such as water

5. Brittle

Water solutions of ionic compounds areusually electrolytes.

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Physical Properties – A Closer Look

Melting and Boiling Point• The electrostatic force of attraction that holds together

all the positive and negative ions in the lattice are very strong; it takes a lot of energy to separate the ions

• Example: The melting point of NaCl is over 800oC.• MP and BP is generally higher when the charge on

ions are greater• Ionic compounds have low volatility (tendency to

vaporize)

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Coulomb’s Law• In Physics, we learned that the force between two

charged particles is equal to:

• Where q1 and q2 are the charges, r is the distance between the two charged particles

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Solubility• Solubility is determined by the degree to which the two

substances mixing are attracted to each other (“like dissolves like”)

• Ionic compounds can dissolve in polar substances because the positively and negatively charged ions are attracted to the partially positive and negative regions in the polar covalent compound

• Example: When NaCl dissolves in water, the partially negative oxygen atoms in the molecule can dislodge the positive ions from the crystal structure

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Solvation of Ionic Compounds

• When an ionic compound dissolves in a polar liquid, the ions get dislodged from the crystal lattice structure.

• In Water: As these ions separate from the lattice, they become surrounded by water molecules and are said to be hydrated and the state symbol (aq) is used.

• In Other Polar Solvent: If a liquid other than water is able to dissolve the solid, the ions are said to be solvated and an appropriate state symbol to denote the solvent is used.

• In the case of solvents like oil or hexane, C6H14, which are non-polar and so have no charge separation, there is no attraction between the liquid and the ions. So here the ions remain tightly bound to each other in the lattice, and the solid is insoluble.

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Electrical Conductivity• Substances can conduct electricity when they are able

to move a charge• Ionic compounds in the solid phase are locked into

place and therefore cannot conduct electricity• Ionic compounds can only conduct charge in the liquid

(molten) phase or when dissolved in a polar solvent

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Brittle• Ionic compounds are brittle, meaning they easily

shatter when a force is applied.• Take a look at the picture below. Why do you think this

is?

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Ionic Character• Ionic compounds typically form between metals, which

have a tendency to lose electrons, and nonmetals, which have a tendency to gain electrons

• In order for two elements to form a binary ionic compound, they must have two very difference electronegativity values

• In general, an electronegativity difference of greater than 1.8 is considered to be ionic

• The larger the electronegativity, the more ionic the bond.

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Bond Continuum• We are going to jump into covalent bonding next.• We will see that the distinction between ionic and

covalent bonds is not black and white, but is best described as a bonding continuum with all intermediate types possible.

• In general, the larger the electronegativity difference, the more ionic the bond. The smaller the en difference, the more covalent the bond.

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Let’s Practice• Explain which of the following pairs will be most likely

to form an ionic bond. A Be and F B Si and OC N and ClD K and S

Consider the difference in electronegativity of each pair:A 1.6 and 4.0 B 1.9 and 3.4 C 3.0 and 3.2 D 0.8 and 2.6 D has the greatest difference, so the compound K2S will be the most ionic.

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More Practice7  Which fluoride is the most ionic?A NaF B CsF C MgF2 D BaF2

8  Which pair of elements reacts most readily? A Li + Br2 B Li + Cl2 C K + Br2 D K + Cl2

9 You are given two white solids and told that only one of them is an ionic compound. Describe three tests you could carry out to determine which it is.

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Answers• 9. Test the melting point: ionic solids have high melting

points. Test the solubility: ionic compounds usually dissolve in water but not in hexane. Test the conductivity: ionic compounds in aqueous solution are good conductors, as are ionic compounds when they are molten

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Lesson 3: Review

Covalent Compounds

Friday, December 5, 2014

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Covalent BondingUnderstandings: • A covalent bond is formed by the electrostatic

attraction between a shared pair of electrons and the positively charged nuclei.

• Single, double, and triple covalent bonds involve one, two, and three shared pairs of electrons respectively.

• Bond length decreases and bond strength increases as the number of shared electrons increases.

• Bond polarity results from the difference in electronegativities of the bonded atoms.

Guidance • Bond polarity can be shown either with partial charges,

dipoles, or vectors.

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Covalent BondingApplications and skills: • Deduction of the polar nature of a covalent bond from

electronegativity values. Guidance • Electronegativity values are given in section 8 of the

data booklet.

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REVIEW: Covalent Bonding

• In a covalent bond, both elements are trying to gain electrons in order to achieve the stable noble gas configuration

• Instead of a transfer of electrons, an electron pair is shared between the two nuclei.

• A group of atoms that are held together by covalent bonds are called a molecule

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Energetics of Bond Formation

• The formation of a covalent bond stabilizes the two atoms entering into the bond so energy is released.

• Energy is always required to break a covalent bond.• The two nuclei are both attracted to the electron pair

while also repelling each other; the bond forms at an equilibrium point that balances the attraction and repulsion

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Lewis Dot Structures• Lewis dot structures are helpful to visualize molecules

and covalent bonds• Atoms enter into covalent bond to try to get a total of

eight valence electrons (octet rule) so we only show valence electrons in the Lewis Dot structure

• Bonds are shown as a line between the two bonded atoms and lone pairs of electrons that do not enter into the bond are shown as dots

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Multiple Bonds• A single bond is a sharing of two electrons• A double bond is a sharing of two pairs of, or four,

electrons• A triple bond is a sharing of three pairs of, or six,

electrons• You can never have a quadruple bond!

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Bond Properties• Bond length: a measure of the distance between the

two bonded nuclei.• Bond strength: usually described in terms of bond

enthalpy, is a measure of the energy required to break the bond.

• As we go down a group, molecules form longer bond lengths

• As bond length increases, bond enthalpy decreases

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Multiple Bond• Double bonds are shorter and stronger than single

bonds and triple bonds are shorter and stronger than double bonds

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Even in the same molecules, double bonds are shorter and

stronger!

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Polar Bonds• A bond is polar when the two elements sharing

electrons share them unevenly• Elements with different electronegativity values will

share electrons unevenly• The term dipole is often used to indicate the fact that

this type of bond has two separated opposite electric charges.

• The more electronegative atom with the greater share of the electrons, has become partially negative or 𝛅–, and the less electronegative atom has become partially positive or 𝛅+.

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Electronegativity Difference

• As the difference in electronegativity between the two element increase, the bond becomes more polar

• NOTE: This does not mean the overall molecule is polar! More on this later!

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Let’s Practice• Use the electronegativity values to put the following

bonds in order of decreasing polarity. a) N–O in NO2

b) N–F in NF3

c) H–O in H2O

d) N–H in NH3

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Challenge Question• Oxygen is a very electronegative element with a value

of 3.4 on the Pauling scale. Can you determine the formula of a compound in which oxygen would have a partial positive charge?

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Polar or Nonpolar?• In Regents Chemistry, we stated that any bonds with an

electronegativity difference of < 0.4 were considered nonpolar.

• In IB Chem, the only bonds that are considered truly nonpolar are between the same element (i.e. N2, O2 etc.). These are called pure covalent molecules.

• However, there are some bonds, notables C-H that have an EN difference of < 0.4 that behave basically nonpolar (even though carbon is slightly more electronegative). Remember – it is all a continuum!

• Polar bonds have more ionic character than nonpolar bonds!

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Bond Continuum

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Let’s Practice10  Which substance contains only ionic bonds?

A NaNO3 B H3PO4 C NH4Cl D CaCl2 11  Which of the following molecules contains the shortest bond between carbon and oxygen?

A CO2 B H3COCH3 C CO D CH3COOH 12  For each of these molecules, identify any polar bonds and label them using δ+ and δ– appropriately:

(a) HBr (b) CO2 (c) ClF (d) O2 (e) NH3 13  Use the electronegativity values in Section 8 of the IB data booklet to predict which bond in each of the following pairs is more polar.

(a) C–H or C–Cl (b) Si–Li or Si–Cl (c) N–Cl or N–Mg

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Solutions10. D11. C

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Lesson 4: Covalent Bonding

StructuresTuesday, December 9, 2014

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4.3 Covalent StructuresUnderstandings: • Lewis (electron dot) structures show all the valence electrons in a

covalently bonded species.

Guidance: • The term ‘electron domain’ should be used in place of ‘negative charge

centre’.• Electron pairs in a Lewis (electron dot) structure can be shown as dots,

crosses, a dash, or any combination.• Coordinate covalent bonds should be covered. • The ‘octet rule’ refers to the tendency of atoms to gain a valence shell

with a total of 8 electrons. • Some atoms, like Be and B, might form stable compounds with

incomplete octets of electrons. • Resonance structures occur when there is more than one possible

position for a double bond in a molecule. • Shapes of species are determined by the repulsion of electron pairs

according to VSEPR theory. • Carbon and silicon form giant covalent/network

covalent/macromolecular structures.

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4.3 Covalent StructuresGuidance: • Allotropes of carbon (diamond, graphite, graphene, C60

buckminsterfullerene) and SiO2 should be covered.

Applications and skills: • Deduction of Lewis (electron dot) structure of molecules and ions

showing all valence electrons for up to four electron pairs on each atom.

• The use of VSEPR theory to predict the electron domain geometry and the molecular geometry for species with two, three, and four electron domains.

• Prediction of bond angles from molecular geometry and presence of non-bonding pairs of electrons.

• Prediction of molecular polarity from bond polarity and molecular geometry.

• Deduction of resonance structures, including C6H6, CO32–, and O3. • Explanation of the properties of giant covalent compounds in terms

of their structures.

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Lewis Dot Diagrams• Lewis Dot diagrams are a way to show how atoms are

bonded in a molecule• Also called a “dot and cross” diagram

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Steps For Drawing Lewis Dot Structures

1. Calculate total number of valence electrons in the molecule. Consider these the money you have at the bank. You cannot spend more than you have and you have to spend it all!

2. Draw the skeletal structure to show how the atoms are linked together.

3. Use a pair of crossed, a pair of dots or a straight line between all of the atoms bonded together

4. Add more electron pairs around the atoms until everyone has 8 valence electrons (except for H which only gets 2 electrons!)

5. Check to make sure everyone has an octet! If they do not, make double and/or triple bonds!

6. Make sure you are showing the # of electrons from Step 1

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Skeleton Help• Often, the most difficult part is deciding how to arrange

the atoms. When we jump into the HL level material, we will learn about formal charge and how that can help us out with more complicated diagrams

• For now, try to put the least electronegative element in the middle

1. Halogens are rarely in the center2. Hydrogen is never in the center3. If you have carbon, put that in the center4. Try to give Group 17 elements one bond, Group 16

elements two bonds, 15 three bonds and 14 four bonds. 5. Try to make the molecule somewhat symmetrical if

possible

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Easy ExampleDraw the Lewis structure for the molecule CCl4.Step 1: Determine # of valence e’s. 12+(7*4) = 32Step 2: Set up skeleton structure Step 3: Use a straight line to represent each bond. You may also use dots or x’sStep 4: Add the remaining amount of electrons. Step 5: Check to make sure everyone has an octet! If not, start to make double and triple bonds!Step 6: Check that you are showing exactly 32 e’s

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Let’s PracticeDraw the following Lewis Dot diagrams1. CH4

2. NH3

3. H2O

4. CO2

5. HCN6. OH- (HINT: What does the -1 charge mean?)7. SO4

2-

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Coordinate Covalent Bonds

• In a coordinate covalent bond, one atoms donates BOTH of the electrons in the covalent bond

• These are also called dative bonds• Sometimes, you might show this in a Lewis Dot

structure with an arrow pointing from the element donating both electrons

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Octet Rule ExceptionsElement

Exception Example

Be Can be happy with 2 bonds

BeCl2

B Can be happy with 3 bonds

BF3

S Can expand octet; form 6 bonds

SF6

P Expands octet, forms 5 bonds

PCl5

Xe Can bond, forms 6 bonds XeF4

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Less Than An Octet• Atoms that have less than a stable octet are said to be

electron deficient and serve as excellent Lewis Acids (electron pair accepter)

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Let’s Practice!14  Draw the Lewis structures of: (a) HF (b) CF3Cl (c) C2H6 (d) PCl3 (e) C2H4 (f) C2H2

15  How many valence electrons are in the following molecules? (a) BeCl2 (b) BCl3 (c) CCl4 (d) PH3 (e) SCl2 (f) NCl3

16  Use Lewis structures to show the formation of a coordinate bond between H2O and H+.

17  Draw the Lewis structures of: (a) NO3

– (b) NO+ (c) NO2– (d) O3 (e) N2H4

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Lesson 5: VSEPR

Wednesday, December 10, 2014

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Molecule Shapes• Lewis Dot structures are helpful because they let us

know how the atoms are arranged and where there are lone pairs of electrons but tell us very little about the shape of the molecules

• Why is shape important? Structure determines functionality!

• Enzymes work because they “fit” another molecule in.

• We touched on molecule shapes last year but will now dive deeper!

• What determines molecule shapes? Electron repulsion! Like charges repel!

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Valence Shell Electron Pair Repulsion

(VSEPR) Theory • This theory is based on the simple notion that because

electron pairs in the same valence shell carry the same charge, they repel each other and so spread themselves as far apart as possible.

• When we said “pair” of electrons that is an oversimplification; in fact when you have multiple bonds they often act together. It is more appropriate to say electron domain so we will be using this term in class

• What matters in determining shape is the total number of electron domains, and this can be determined from the Lewis structure.

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Central Atoms• We often look at the central atom for simple

compounds to determine the shape of the molecule• How many electron domains exist in the central atom of

the following molecules whose Lewis structures are shown?

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VSEPR• The repulsion applies to electron domains, which can be

single, double, or triple bonding electron pairs, or non-bonding pairs of electrons.

• The total number of electron domains around the central atom determines the geometrical arrangement of the electron domains.

• The shape of the molecule is determined by the angles between the bonded atoms.

• Non-bonding pairs (lone pairs) have a higher concentration of charge than a bonding pair because they are not shared between two atoms, and so cause slightly more repulsion than bonding pairs. The repulsion decreases in the following order: lone pair–lone pair > lone pair–bonding pair >

bonding pair–bonding pair

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Lone Pairs vs. Bonded Pairs

• Since lone pairs have a higher concentration of charge, molecules with lone pairs will have slight distortions in the expected bonding angles because they will push away the other electron domains more!

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Two Electron Domains• The furthers away two electron domains can be from

each other is 180 degrees, giving those molecules a linear shape

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Three Electron Domains• Molecules with three electron domains will position

them at 120° to each other, giving a triangular planar shape to the electron domain geometry.

• If only two of the three domains are bonding, the shape of the molecule will be bent or V-shaped and the bond angle is 117 degrees.

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Four Electron Domains• Molecules with four electron domains will position them

at 109.5° to each other, giving a tetrahedral shape to the electron domains

• If all four sides are bonding electrons, the molecule will also have a tetrahedral shape (e.g. CH4) with bond angles of 109.5°

• If three sides are bonding, the molecule will have a trigonal pyramidal shape (e.g. NH3) with bond angles of 107°

• If two sides are bonding, the molecule will have a bent or V-Shape (e.g. H2O) with bond angles of 105°

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Summary# of Electron Domains

# of Bonded Domains

# of Long Pairs

Shape Bond Angles

2 2 0 Linear 180

3 3 0 Trigonal Planar 120

3 2 1 Bent or T-Shaped 117

4 4 0 Tetrahedral 109.5

4 3 1 Trigonal Pyramidal

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4 2 2 Bent/T-Shaped 105

5 5 0 Trigonal Bipyramidal

90 & 120

6 6 0 Octahedral 90

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Expanded Octets• More on the shapes where there are 5 and 6 electron

domains later in the unit when we go into the HL level material

• In order to get those shapes, the atoms need expanded octets

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Steps To Determine Shape

1  Draw the Lewis structure2  Count the total number of electron domains on the central atom. 3  Determine the electron domain geometry as follows:

2 electron domains → linear3 electron domains → triangular planar4 electron domains → tetrahedral

4  Determine the molecular geometry from the number of bonding electron domains. 5  Consider the extra repulsion caused by the lone pairs and adjust the bond angles accordingly.

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Let’s Practice!18  Predict the shape and bond angles of the following molecules: (a) H2S (b) CF4 (c) HCN (d) NF3 (e) BCl3 (f) NH2Cl (g) OF2

19  Predict the shape and bond angles of the following ions: (a) CO3

2– (b) NO3– (c) NO2

+ (e) ClF2+ (d) NO2

– (f) SnCl3–

20 How many electron domains are there around the central atom in molecules that have the following shapes? (a) tetrahedral (b) bent (c) linear (d) trigonal pyramidal (e) triangular planar

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Polarity• Bond polarity is whether the pair of electrons are

shared evenly among the two nonmetals covalently bonded togethero A bond is considered polar if the electronegativity difference between the two

elements is less than 0.4

• Molecular polarity refers to whether there is an overall uneven charge distribution in the molecule called a dipole momento A molecule with all nonpolar bonds will always be nonpolar overallo A molecule with polar bonds can be nonpolar overall IF the molecule is

symmetrical, meaning the forces pulling on the electron pairs all balance out leading to no net force pulling on the bonding pairs

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Molecular PolarityMolecular polarity depends on:

1. The polarity of the bonds2. The shape of the molecule

• When molecules are symmetrical, even though the bonds are polar, the charge distributions effectively cancel each other out; see below

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What shapes can be nonpolar?

• In order for a molecule to have polar bonds but be nonpolar overall, it has to be symmetrical, meaning that the same atom is attached to the central atom

• In addition, there cannot be lone pairs attached to the central atom as those electron domains are different from the bonded domain

• The shapes we have learned so far that can even be nonpolar when the bonds are polar are linear, trigonal pyramidal and tetrahedral

• Trigonal pyramidal and bent molecules can not be nonpolar with polar bonds.

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Net Dipole• If either the molecule contains bonds of different

polarity, or its bonds are not symmetrically arranged, then the dipoles will not cancel out, and the molecule will be polar.

• Another way of describing this is to say that it has a net dipole moment, which refers to its turning force in an electric field.

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It helps to think about dipole in terms of vectors (FLASHBACK to Freshman year Physics!)

If the Fnet of all the resultant force vectors is equal to zero, there is no overall dipole moment and the molecule is nonpolar.

Another way to conceptualize this is by thinking about a “tug of war” (see left)

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Let’s Practice• 21 Predict whether the following will be polar or non-

polar molecules: (a) PH3

(b) CF4

(c) HCN (d) BeCl2

(e) C2H4

(f) ClF (g) F2

(h) BF3

a) Polarb) Nonpolarc) Polard) Nonpolare) Nonpolarf) Polarg) Nonpolarh) Nonpolar

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Trickier Question22 The molecule C2H2Cl2 can exist as two forms known as cis–trans isomers, which are shown below. (Sneak Preview for Orgo: The double bond locks this molecule in place preventing the atoms from rotating around each other. Can’t wait for Orgo!)

Determine whether either of these has a net dipole moment.

Nonpolar Polar

108

Lesson 6: Lewis Dot Structures

and Resonance

Wednesday, December 17

109

Warm-UpTry to draw the following Lewis Dot structures:1. CO3

2-

2. O3

What do you notice?

110

Delocalized Electrons• In these Lewis Dot structures, we find that we can draw

more than one type of structure.

• Which one is correct?• Turns out…neither of them is exactly correct. These

structures suggest that the molecule should contain one oxygen–oxygen double bond and one oxygen–oxygen single bond, which we would expect to be differentin bond length and bond strength. However, experimental data reveal that ozone actually contains two equal oxygen–oxygen bonds, intermediate in length and strength between single and double bonds.

111

112

Delocalized Electrons• What is going on here?• In some molecules, bonding electrons are less

restricted and are shared between more than one bonding position.

• These electrons are said to be delocalized. Free from the constraints of a single bonding position, delocalized electrons spread themselves out, giving greater stability to the molecule or ion.

113

Resonance• Resonance occurs when more than one valid Lewis

structure can be drawn for a particular molecule. The true structure is an average of these, known as a resonance hybrid.

• Let’s try to draw the resonance structures for the carbonate ion CO3

2–.

114

Benzene• Benzene is an important organic molecule that has

delocalized electrons• The resonance structure gives this molecule additional

stability that would otherwise not be predicted

115

Benzene• When we learn about sigma and pi bonds later in the

unit when we move onto the HL level, we will see that the electrons in the double bonds of all the carbons are all delocalized and spread out among the carbon atoms

• Bond order is a measure of the number of electrons involved in bonds between two atoms. Values for bond order are: single bonds = 1, double bonds = 2, triple bonds = 3. Resonance hybrids have fractional values of bond order. The carbon-carbon bonds in benzene have a bond order of 1.5.

116

Let’s Practice• Compare the structures of CH3COOH and CH3COO– with

reference to their possible resonance structures. • Put the following species in order of increasing carbon–

oxygen bond length: CO CO2 CO32– CH3OH

• By reference to their resonance structures, compare the nitrogen–oxygen bond lengths in nitrate(V) (NO3

–) and nitric(V) acid (HNO3).

117

Network Solids• Network solids are a special type of covalent

compound where there are no molecules or discrete particles but rather, all the atoms are covalently bonded together into one giant crystal

• It is referred to as a giant molecular or network covalent structure or macromolecular structure.

• These network solids have very different characteristics than regular covalent compounds (i.e. higher boiling points, higher melting points, very hard, etc.)

118

Two Network Solids To Know!!!!

• You need to know the following two network solids for the IB exam:

1. Diamond – in diamond, each carbon atom is covalently bonded to four other carbon atoms in a continuous tetrahedral shape

2. Silicon Dioxide – SiO2 (commonly referred to as silica or quartz) where each Si atom is covalently bonded to four O atoms, and each O to two Si atoms

119

Quartz/Silica• SiO2 refers to the ratio of atoms within the giant

molecule – it is an empirical formula and the actual number of atoms present will be a very large multiple of this. As the atoms are strongly held in tetrahedral positions that involve all four silicon valence electrons, the structure has the following properties:

• strong;• insoluble in water;• high melting point;• non-conductor of electricity.

• These are all properties we associate with glass and sand – different forms of silica.

120

Allotropes• Allotropes are different forms of the same elements

that have different structures in the same phase• The different structures mean the different allotropes

have different physical and chemical properties• For example, O2 gas and O3 gas are bonded differently

and have very different characteristics• You need to know the four allotropes of carbon

for the IB exam – diamond, graphite, fullerene, and graphene

121

Characteristic

Graphite Diamond Fullerene Graphene

Structure Each C is sp2 hybridized, bonded to 3 other Cs and found in layers that can slide of each other; bond angle 120

Each C is sp3 hybridized and bonding to 4 other C atoms; bond angle 109.5

Each C atom is sp2 hybridized and bonded in a sphere of 60 carbon atoms, consisting of 12 pentagons and 20 hexagons.

Each C atom is covalently bonded to 3 others, asin graphite, forming hexagons with bond angles of 120°. Single layer.

Electrical Conductivity

Good None Semi-conductor

Very good

Thermal Conductivity

Poor Very good Very low Very good; best known

Special Properties

Soft Hardest known

Light and strong

Thinnest material ever

Uses Pencils! Cutting classJewels

Nanotech Lots of new stuff!

122

123

124

125

Graphene• Since graphene is relatively new (discovered in 2004),

the IB might want you to know more about it. It is at the current edge of technology

126

Let’s Practice25  Describe the similarities and differences you would expect in the properties of silicon and diamond. 26  Explain why graphite and graphene are good conductors of electricity whereas diamond is not (HINT: which ones have resonance?).

127

Answers25  Similarities: strong, high melting points, insoluble in water, non-conductors of electricity, good thermal conductors. Differences: diamond is stronger and more lustrous; silicon can be doped to be an electrical conductor. 26  Graphite and graphene have delocalized electrons that are mobile and so they conduct electrical charge. In diamond all electrons are held in covalent bonds and so are not mobile.

128

Lesson 7: Intermolecula

r ForcesThursday, December 18

129

Warm-Up• What were the five intermolecular forces we learned

about in Regents Chemistry last year?

130

Topic 4.4 – Intermolecular Forces

Understandings: • Intermolecular forces include London (dispersion)

forces, dipole–dipole forces, and hydrogen bonding. Guidance • The term ‘London (dispersion) forces’ refers to

instantaneous dipole–induced dipole forces that exist between any atoms or groups of atoms and should be used for non-polar entities. The term ‘van der Waals’ is an inclusive term, which includes dipole–dipole, dipole–induced dipole, and London (dispersion) forces.

• The relative strengths of these interactions are London (dispersion) forces < dipole–dipole forces < hydrogen bonds.

131

Topic 4.4 – Intermolecular Forces

Applications and skills: • Deduction of the types of intermolecular force present

in substances, based on their structure and chemical formula.

• Explanation of the physical properties of covalent compounds (volatility, electrical conductivity, and solubility) in terms of their structure and intermolecular forces.

132

Intermolecular Forces• Covalent bonds hold together atoms within a molecule

(intramolecular)• However, in between molecules, intermolecular forces

hold them together and allow them to enter into the liquid and solid phase

• The type of intermolecular force depends on the polarity of the molecule; the strength of the intermolecular force depends on the type and the size of the molecules

• The strength of intermolecular forces determines the physical properties of a substance. Volatility, solubility, and conductivity can all be predicted and explained from knowledge of the nature of the forces between molecules.

133

London Dispersion Forces

• London dispersion forces form between nonpolar molecules and are the result of temporary or induced dipole moments

• Nonpolar molecules do not have a dipole moment, meaning there is no partial positive or negative side; HOWEVER, since electrons are always moving, at any given moment there may be an uneven distribution of charge giving one side a partial negative and the other side a partial positive.

• This in turn can affect the electron clouds of the nearby atoms causing an induced dipole

134

London Dispersion Forces

135

London Dispersion • The intermolecular force is the weak attraction between

these temporary dipoles• These forces increase as molecules get larger; larger

electron clouds increase the probability of a charge imbalance

136

137

PropertiesSince London Dispersion Forces are so weak, the nonpolar molecules that are held together by them have:1. Low melting points2. Low boiling points3. High vapor pressure

• Polar molecules also have London dispersion forces but these are far weaker than some of the other forces at work!

138

Dipole-Dipole Attractions• Form between polar molecules• In polar molecules, one end of the molecule is electron

deficient with a partial positive charge (δ+), while the other end is electron rich with a partial negative charge (δ–). This is known as a permanent dipole.

• It results in opposite charges on neighboring moleculesattracting each other, generating a force knownas a dipole–dipole attraction.

139

Strength of Dipole-Dipole• Dipole-dipole attractions are stronger than London

dispersion forces meaning these substances have a higher bp and mp

• The strength of the dipole-dipole attractions is due to the degree of polarity, size of the molecule, orientation and more

140

Van der Waals’ Forces• You may also hear the term van der Waals’ forces used

when talking about intermolecular attraction• This is an umbrella term used to describe any

intermolecular force not due to full ion-ion attraction or covalent bonding

• Sometimes, in the case of very large molecules, these forces will occur intramolecularly (i.e. large proteins)

141

Hydrogen Bonding• Hydrogen bonds are a special, exceptionally strong type

of intermolecular force that only occurs when you have hydrogen covalently bonding to an extremely electronegative element, nitrogen, fluorine or oxygen.

• Hydrogen bonding is stronger than any van der Waals’ force and accounts for the exceptionally high boiling points of water, ammonia (NH3), alcohols and HF.

• This type of intermolecular force is extremely important when looking at Biochemistry as it holds together our DNA and is important in the folding of proteins!

142

Why so strong?• The large electronegativity of the O, N or F make the

hydrogen atom bonded to them extremely electron deficient.

• The hydrogen atom has no other electrons other than the one in that bonding pair. It now exerts a very strong attractive force on another bonding domain from a neighboring molecule.

143

Hydrogen bonds explain boiling point trends

144

H2O• Without hydrogen bonding, water would be a gas at

room temperature meaning life as we know it on this planet could not exist

• Also, hydrogen bonds account for the fact that solid ice is one of the only substances in the universe LESS DENSE than liquid water. Without them, ice would not float and our oceans would freeze from the bottom up!

145

Biochemistry• The hydrogen bond is important in the Chemistry of life. • TIP FOR NEXT YEAR: This will be covered in Biology but

is also an option for next year – Option B: Biochemistry. Keep in mind that since you are covering this in Bio, it means that if we cover a different option in Chem, you have two different choices for Paper 3.

• Hydrogen bonding cause our proteins to fold and holds together our very DNA!

• The strands of our DNA are held together by hydrogen bonds. These bonds are strong enough to hold the strands together but not so strong that they cannot be separated when we need to replicate a gene!

146

Advance Molecular Orbital Theories on H-

Bonds• The extreme electronegativity different between F, O

and N does not fully explain the strength and behavior of hydrogen bonds.

• In reality, hydrogen bonds go beyond just being a stronger version of dipole-dipole forces; they have some properties of covalent bonds.

• This is an area Chemistry still being investigated!!

147

All of these forces are weaker than bonds!!

148

Let’s Practice1. Methoxymexane (CH3–O–CH3) boils at a much lower

temperature than ethanol (CH3CH2–O–H). Use your knowledge of intermolecular forces to explain why.

2. Put the following molecules in order of increasing boiling point and explain your choice: CH3CHO, CH3CH2OH and CH3CH2CH3

149

Let’s PracticeWhat is the strongest intermolecular force holding together the following substances?1. NH3

2. CCl43. C4H9OH

4. N2

5. PH3

6. CO2

7. CH3F

8. HF9. Na2O

10.SiO2

150

Lesson 8: Physical

Properties of Covalent

Compounds Friday, December 19

151

Properties of Covalent Compounds - Solubility

• Solubility – polar covalent compounds are generally soluble in other polar covalent compounds and can mix with ionic compounds; nonpolar covalent compounds are soluble in other nonpolar covalent compounds. “Like dissolves like”

• The larger the polar covalent compounds, the less soluble they become in other polar covalent compounds because the dipole becomes a less important component of the molecule; the nonpolar parts of the molecule overwhelm the dipole-dipole attractions

• Network solids are not soluble!

152

Electrical Conductivity• Covalent compounds do not contain ions so they are

not able to conduct electricity • HOWEVER…some covalent compounds can ionized

under certain conditions (HINT: remember our acids!) and can conduct electricity in solution. Examples include HCl, H2SO4, etc.

• Some network solids such as graphite and graphene are electrical conductors due to their mobile electrons.

153

Summary

154

Let’s Practice

155

Answers

156

QUIZ!!20 minutes!

157

Lesson 9: Metallic Bonding

Tuesday, December 23

158

Topic 4.5 Metallic Bonding

Understandings: • A metallic bond is the electrostatic attraction between a

lattice of positive ions and delocalized electrons. • The strength of a metallic bond depends on the charge

of the ions and the radius of the metal ion. • Alloys usually contain more than one metal and have

enhanced properties. Guidance • Examples of various alloys should be covered.

159

Topic 4.5 Metallic Bonding

Applications and skills: • Explanation of electrical conductivity and malleability in

metals. • Explanation of trends in melting points of metals. Guidance • Trends should be limited to s- and p-block elements. • Explanation of the properties of alloys in terms of non-

directional bonding.

160

Metallic Bonding - Intro• Metals tend to have low ionization energies and tend to

lose electrons in order to gain Noble gas configurations.• In the elemental state, when there is no other element

present to accept the electrons and form an ionic compound, the outer electrons are held only loosely by the metal atom’s nucleus and so tend to ‘wander off’ or, more correctly, become delocalized.

• The metal atoms become positively charged ions and form a regular lattice structure through which these electrons can move freely.

• There is a force of electrostatic attraction between the lattice of cations and the delocalized electrons, and this is known as metallic bonding.

161

Referred to as “sea of electrons”

162

Metallic Bond Strength• Metallic bond strength is determined by:1. The number of delocalized electrons2. The charge on the cation3. The radius of the cation

Example 1: If we compare the melting point of Na which gets 1 delocalized electron per atom and Mg which gets 2 delocalized electrons per atom, Na melts at 98 degree C and Mg melts at 650 degrees C.Example 2: Melting points decrease down a group as the attraction between the nucleus and the delocalized electrons decrease.

163

Metallic Properties• These delocalized electrons give metallic

substances very unique properties including:oGood electrical conductivityoGood thermal conductivityoMalleableoDuctileoHigh melting pointso Shiny, lustruous

164

165

Alloys• When metals form homogeneous solutions, they are

called alloys• Alloys are made by mixing the metals together while

they are in the molten or liquid state• Even if the positive cations have different sizes, since

they are not held in a strict lattice • Alloys often have different characteristics than

the original metals including being more stable and stronger; it is more difficult for the atoms to move over each other when they are different sizes

166

Some Common AlloysAlloy Compositio

nUses

Brass Copper and zinc

Door handles, screws

Steel Iron, Carbon and other metals

Bridges, buildings

Dental Amalgam

Mercury, silver and tin

Used to be used in teeth fillings

167

AlloysSmall amounts of a another element added to a metal can change its overall properties.

For example, adding a small amount of carbon to iron, will significantly increase its hardness and strength forming steel.

168

Semimetals

The electrons in semimetals are much less mobile than in metals, hence they are semiconductors

Silicon

Magnesium

169

Comparison of Types of Bonding  Ionic Covalent Metallic

Formation Anion & cation Transferred electrons

Shared electrons

Cations in a sea of mobile valence electrons

Source Metal + nonmetal

Two nonmetals Metals only

Melting point Relatively high

Relatively low Generally high

Solubility Dissolve best in water and polar solutions

Dissolve best in non-polar solvents

Generally do not dissolve

Conductivity Water solutions conduct electricity

Solutions conduct electricity poorly or not at all

Conduct electricity well

Other properties

Strong crystal lattice

Weak crystal structure

Metallic properties; luster, malleability etc. 170

Bonding Types Are Continuous

• There are no clear boundaries between the three types of bonding.

• Chemical bonding may be thought of as a triangle.

• Each vertex represents one of the three types of chemical bonds.

• There are all degrees of bonding types between these extremes.

171

172

Let’s Practice31  Which is the best definition of metallic bonding?

A  the attraction between cations and anions B  the attraction between cations and delocalized electrons C  the attraction between nuclei and electron pairs D  the attraction between nuclei and anions

32  Aluminium is a widely used metal. What properties make it suitable for the following applications? (a) baking foil (b) aircraft bodywork (c) cooking pans(d) tent frames 33 Suggest two ways in which some of the properties of aluminium can be enhanced.

173

174

Lesson 10: HL Expanded Octets and

ShapesTuesday, January 6

175

14.1 Further Aspects of Covalent Bonding - HL

Understandings: • Covalent bonds result from the overlap of atomic orbitals. A sigma bond

(σ) is formed by the direct head-on/end-to-end overlap of atomic orbitals, resulting in electron density concentrated between the nuclei of the bonding atoms. A pi bond (π) is formed by the sideways overlap of atomic orbitals, resulting in electron density above and below the plane of the nuclei of the bonding atoms.

Guidance: • The linear combination of atomic orbitals to form molecular orbitals should

be covered in the context of the formation of sigma (σ) and pi (π) bonds. • Formal charge (FC) can be used to decide which Lewis (electron dot)

structure is preferred from several. The FC is the charge an atom would have if all atoms in the molecule had the same electronegativity. FC = (number of valence electrons) – 1⁄2(number of bonding electrons) – (number of non-bonding electrons). The Lewis (electron dot) structure with the atoms having FC values closest to zero is preferred.

• Exceptions to the octet rule include some species having incomplete octets and expanded octets.

176

14.1 Further Aspects of Covalent Bonding - HL

Guidance • Molecular polarities of geometries corresponding to five

and six electron domains should also be covered. • Delocalization involves electrons that are shared

by/between all atoms in a molecule or ion as opposed to being localized between a pair of atoms.

• Resonance involves using two or more Lewis (electron dot) structures to represent a particular molecule or ion. A resonance structure is one of two or more alternative Lewis (electron dot) structures for a molecule or ion that cannot be described fully with one Lewis (electron dot) structure alone.

177

14.1 Further Aspects of Covalent Bonding - HL

Applications and skills: • Prediction whether sigma (σ) or pi (π) bonds are formed from the

linear combination of atomic orbitals. • Deduction of the Lewis (electron dot) structures of molecules and

ions showing all valence electrons for up to six electron pairs on each atom.

• Application of FC to ascertain which Lewis (electron dot) structure is preferred from different Lewis (electron dot) structures.

• Deduction using VSEPR theory of the electron domain geometry and molecular geometry with five and six electron domains and associated bond angles.

• Explanation of the wavelength of light required to dissociate oxygen and ozone.

• Description of the mechanism of the catalysis of ozone depletion when catalysed by CFCs and NOx.

178

Expanded Octets• While the octet is the most common arrangement for

atoms when entering into covalent bonds, there are some exceptions to the octet rule that include expanded octets utilizing the d-sublevel

• Since the atom needs to use the d sublevel to expand its octet, only elements in Row 3 and below can have expanded octets

• Elements such a sulfur and phosphorus can create compounds with 5 or 6 electrons.

179

5 Electron Domains• The electron domain in a 5 electron domain compound,

such as PCl5, will position themselves in a trigonal bipyramidal shape with bond angles of 90°, 120°, and 180°

• As with the other shapes, the molecular shape will be slightly different depending on how many domains on bonded vs. lone pairs

180

Additional Shapes – 5 Electron Domains

• 5 bonded domains – electron configuration triangular bipyramidal and molecular configuration triangular bipyramidal

• 4 bonded domains - electron configuration triangular bipyramidal and molecular configuration unsymmetrical tetrahedron or see-saw

• 3 bonded domains - electron configuration triangular bipyramidal and molecular configuration T-shaped structure

• 2 bonded domains - electron configuration triangular bipyramidal and molecular configuration linear

181

Shapes

See-saw T-Shaped

Linear

182

6 Electron Domains• Molecules with six electron domains will position them

in an octahedral shape with angles of 90°.

183

6 Electron Domains – Molecular Shapes

• 6 bonded domains – electron configuration octahedral and molecular configuration octahedral

• 5 bonded domains – electron configuration octahedral and molecular configuration square pyramidal

• 4 bonded domains – electron configuration octahedral and molecular configuration square planar

184

Shapes

Square Pyramidal Square Planar

185

186

187

188

Molecular Polarity• If there are no lone pairs and all the atoms attached to

the central atom are the same, the molecules are non-polar as there is no net dipole for 5 and 6 electron domain shapes. o For example, PCl5 and SF6 are both non- polar.

• However, here the molecule may be non-polar even with different atoms bonded to the central atom depending on HOW they are bonded!o For example, SBrF5 has a net dipole and is polar; PCl3F2 has no net dipole due to

cancellation, so it is non-polar.

189

Molecular Polarity (cont.)• Whereas before, the presence of lone pairs always

meant the molecule was polar, if the lone pairs are on opposite sides of the molecule and the same element is bound to the central atom, the molecule can be nonpolar!

190

Let’s Practice

191

Answers

192

Lesson 11: HL Formal ChargeWednesday, January 7

193

Warm-UpPair ShareWhen we can draw more than one Lewis Dot structure for a compound, how do we know which one is correct?

194

Sulfur Dioxide• Let’s take a look at SO2, sulfur dioxide. Knowing that

sulfur can sometimes expand its octet, we can draw the Lewis Dot structure two ways:

Way 1

Way 2

Which one is correct?

195

Formal Charge• Formal charge is used to help us determine which

structure is most stable• Formal charge treats each covalent bond as an equal

electron distribution with each atom “owning” 1 of the electrons in the bond.

• Each atom also “owns” every electron that is in a lone pair around it

• When calculating formal charge we count how many electrons each atom “owns” and compare that to the original number of valence electrons it started with; it is generally more favorable if those numbers are the same!

196

Calculating Formal Charge

• The number of valence electrons (V) is determined from the element’s group in the Periodic Table.

• The number of electrons assigned to an atom in the Lewis (electron dot) structure is calculated by assuming that: o (a)  each atom has an equal share of a bonding electron pair (one electron per

atom), even if it is a coordinate bond (1⁄2B); o (b)  an atom owns its lone pairs completely (L).

• This means that the number of electrons assigned = 1⁄2 number of electrons in bonded pairs (1⁄2B) + number of electrons in lone pairs (L)

• So overall: FC = V – (1⁄2B + L)

197

198

Looking Back at SO2

199

SO2• We can conclude that structure (ii) where all atoms

have a formal charge of zero is the most stable structure for SO2.

200

Let’s Practice• Use the concept of formal charge to determine which of

the following Lewis (electron dot) structures for XeO3 is preferred?

201

Other Considerations• In addition to formal charge, when looking at

compounds where multiple Lewis Dot structures can be drawn, we should also look at electronegativity

• So, a useful guideline to follow is that the most stable of several Lewis (electron dot) structures is the structure that has:

• the lowest formal charges and• negative values for formal charge on the

more electronegative atoms. • Which is the correct structure here?

202

Let’s Practice39  Use the concept of formal charge to explain why BF3 is an exception to the octet rule. 40  Draw two different Lewis (electron dot) structures for SO4

2–, one of which obeys the octet rule for all its atoms, the other which has an octet for S expanded to 12 electrons. Use formal charges to determine which is the preferred structure.

203

Answers

204

Lesson 12: Sigma and Pi

Bonds Thursday, January 8

205

Molecular Orbitals• When atoms covalently bond together, the s and p

electron orbitals actually change and hybridize in order to form molecular or bonding orbitals

• The molecular orbital is formed from one electron from each atom overlapping and combining in an orbital with lower energy than the two atomic orbitals where they started out

206

Sigma Bonds• When two atomic orbitals overlap along the bond axis –

an imaginary line between the two nuclei – the bond is described as a sigma bond, denoted using the Greek letter σ.

• This type of bond forms by the overlap of s orbitals, p orbitals, and hybrid orbitals (to be described in the next section) in different combinations.

• It is always the bond that forms in a single covalent bond!!!

207

Sigma Bond Picture

Sigma Bonds• s orbital

overlap

• p orbital overlap

• s and p overlap

Sigma Bonds (σ) bond• All single bonds are sigma• In order to be a sigma bond, the bond must overlap

along the bond axis.

210

Pi Bonds• When two p orbitals overlap sideways, the electron

density of the molecular orbital is concentrated in two regions, above and below the plane of the bond axis.

• This is known as a pi bond, denoted using the Greek letter π.

• This type of bond only forms by the overlap of p orbitals alongside the formation of a sigma bond.

• In other words, pi bonds only form within a double bond or a triple bond.

211

Pi Bond Picture

pi (π) bond• Two p orbitals overlap sideways• Electron density becomes concentrated in two regions.

Where?

pi (π) bond• Pi bonds can only form alongside formation of a sigma

bond.• Pi bonds therefore only form within a double or triple

bond.

Pi Bonds vs. Sigma Bonds

• Pi bonds are weaker than sigma bonds, as their electron density is further from the nucleus.

215

Lesson 13: Hybridization

Friday, January 9

216

14.2 HybridizationUnderstandings: • A hybrid orbital results from the mixing of different

types of atomic orbitals on the same atom. Applications and skills: • Explanation of the formation of sp3, sp2, and sp hybrid

orbitals in methane, ethene, and ethyne. • Identification and explanation of the relationships

between Lewis (electron dot) structures, electron domains, molecular geometries, and types of hybridization.

Guidance • Students need only consider species with sp3, sp2, and

sp hybridization.

217

Hybridization• We know that both s and p atomic orbitals take part in

covalent bonding• However, they are (a) of different energy levels with s

being lower in energy and (b) have very different shapes. This would lead to uneven bonds for an atom that can make four bonds! So what happens?

218

Closer Look: Carbon• Carbon starts with 2 electrons in the 2s sublevel and 2

electrons in the 2p sublevel. But we know that they need to form 4 covalent bonds to get a stable octet so something has to happen to these orbitals to get each electron into an orbital alone!

• Step 1: A process known as excitation occurs in which an electron is promoted within the atom from the 2s orbital to the vacant 2p orbital. Now each atomic orbital has 1 electron and can bond with another atom.

• Step 2: Atomic orbitals hybridize to form four bonding orbitals each with the same amount of energy called sp3 orbitals

219

220

Hybridization• The details of hybridization are complex and depend on

an understanding of quantum mechanics, but in essence unequal atomic orbitals within an atom mix to form new hybrid atomic orbitals which are the same as each other, but different from the original orbitals.

• This mixing of orbitals is known as hybridization. • Hybrid orbitals have different energies, shapes, and

orientation in space from their parent orbitals and are able to form stronger bonds by allowing for greater overlap.

• You only need to know sp, sp2 and sp3 hybrid orbitals for IB but know that there are also sp3d and sp3d2

221

Advanced Work• We are only going to learn the valence bond theory and

not the more robust molecular orbital theory• In reality, the hybridization is much more complex and

when orbitals overlap they interfere constructively forming bonding orbitals and destructively forming anti-bonding orbitals

• All this is way beyond even the scope of IB but may be required in advance Chemistry courses in college!

222

S and P Hybrid Orbitals• Hybridization of one s orbital with three p orbitals

produces four so-called sp3 hybrid orbitals that are equal to each other. Their shape and energy have properties of both s and p, but are more like p than s.

• Hybridization of one s orbital with two p orbitals will produce three equal sp2 hybrid orbitals. In this case, one p orbital is left unhybridized and can participate in pi bonding. You will see this when there are double bonds.

• Finally, hybridization of one s orbital with one p orbital will produce two equal sp hybrid orbitals. In this case, two p orbitals are left unhybridized and can participate in pi bonding. You will see this when there are triple bonds.

223

sp3 hybridization • When carbon forms four single bonds, it undergoes sp3

hybridization, producing four equal orbitals.

• These orbitals orientate themselves at 109.5°, forming a tetrahedron. Each hybrid orbital overlaps with the atomic orbital of another atom forming four sigma bonds.

• For example, methane, CH4.

224

225

sp2 hybridization • When carbon forms a double bond, it undergoes sp2

hybridization, producing three equal orbitals.

• These orbitals orientate themselves at 120°, forming a triangular planar shape. Each hybrid orbital overlaps with a neighboring atomic orbital, forming three sigma bonds.

• For example, ethene, C2H4.

226

227

sp hybridization • When carbon forms a triple bond, it undergoes sp

hybridization, producing two equal orbitals.

• These orbitals orientate themselves at 180°, giving a linear shape. Overlap of the two hybrid orbitals with other atomic orbitals forms two sigma bonds.

• For example, ethyne, C2H2.

228

229

Triple Bonds• In a triple bond, each carbon has two unhybridized p

orbitals that are orientated at 90° to each other. As these overlap each other sideways, two pi bonds form representing four lobes of electron density around the atoms.

• These coalesce into a cylinder of negative charge around the atom, making the molecule susceptible to attack by electrophilic reagents (those that are attracted to electron-dense regions).

230

Expanded Octets• If an atom has 5 electron domains, one d orbital gets

involved and the hybridization is sp3d which produces five equivalent orbitals orientated to the corners of a triangular bipyramid.

• If an atom has 6 electron domains, two d orbitals get involved and the hybridization is sp3d2 which produces six equivalent orbitals orientated to the corners of a octahedral.

231

Lone Pairs• Lone pairs DO get counted in the electron hybridization• For example, in ammonia, NH3, the non-bonding pair

on the N atom resides in a sp3 hybrid orbital.

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Molecule Shapes• Although we have focused mostly here on examples

from organic chemistry (those involving carbon), this concept of hybridization can be used to explain the shape of any molecule.

• And conversely the shape of a molecule can be used to determine the type of hybridization that has occurred. The relationship is as follows:

• tetrahedral arrangement ↔ sp3 hybridized• triangular planar arrangement ↔ sp2 hybridized • linear arrangement ↔ sp hybridized

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Let’s Practice• Urea (see below) is present in solution in animal urine.

What is the hybridization of C and N in the molecule, and what are the approximate bond angles?

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More Practice

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Even More Fun!

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Lesson 15: Hybridization and Stuctures

Tuesday, January 13, 2015

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Looking Back – Delocalized Electrons

Now that we have learned about hybridization, we are going to take a look back at some of the more advance chemical topics.

Resonance• Now we can see that resonance generally occurs when

there is sp2 or sp3 hybridization• The sigma bonds form in the same plane whereas the

unhybridized p-orbitals

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Nitrate Ion - Example• NO3

+

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Benzene• In Benzene, all of the Carbon atoms have sp2

hybridization and the unhybridized p orbitals form pi bonds above and below the bonding plane

• These electrons become delocalized and are shared among all the carbon atoms!

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Lesson 16: Ozone

Wednesday, January 14, 2015

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Lesson 18: Review

Thursday, January 15, 2014

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Lesson 18: ReviewFriday, January 16, 2015