IB Chemistry Year 1 HL The Brooklyn Latin School

IB Chemistry Year 1 HL The Brooklyn Latin School
Bonding Topics 4 and 14 IB Chemistry Year 1 HL The Brooklyn Latin School

REVIEW: Ionic Bonding Wednesday, December 3, 2014

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.

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.

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

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.

Sneak Peak: Redox Formation of ionic compounds are always redox reactions Oxidation is losing electrons Reduction is gaining electrons LEO the Lion goes GER

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!)

Let’s Practice What kind of ion is formed by the following elements?
Lithium Sulfur Argon Oxygen 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!!

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!

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+.

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 charge E.g Iron +3 is Iron (III) where Iron +2 is Iron (II) Anions, or negative ions, get their ending changed to –ide E.g. Oxygen -2 is oxide and Nitrogen -3 is nitride

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!!!

Memorize!

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

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 Formulas
1. 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 2 4. Write the final formula using subscripts to show the number of each ion. Al2O3

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! Mg3N2

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)

Let’s Practice Write the formula for ammonium phosphate.

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.

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!

More Lies Most 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

Typical Way Ionic Bonding is Shown
What is wrong here?

Let’s Practice 1 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: 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.

Lesson 2: Advanced Ionic Bonding
Thursday, December 4th, 2014

Ionic Bonding Ionic bonding involves 3 aspects:
Ionic bonds result from the attractions between positive and negative ions. Ionic bonding involves 3 aspects: Loss of an electron(s) by one element. Gain of electron(s) by a second element. Attraction between positive and negative ions. 30

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 Examples Na atom 1s2 2s2 2p6 3s One electron extra Cl atom s2 2s2 2p6 3s2 3p One electron short of a stable octet Na+ Ion s2 2s2 2p Stable octet Cl- Ion s2 2s2 2p6 3s2 3p Stable octet Positive ions attract negative ions forming ionic bonds. 31

Ionic Bonding Ionic substances are made of repeating arrays of positive and negative ions. An ionic crystal lattice 32

Ionic Bonding The array is repeated over and over to form the crystal lattice. Model of a Sodium chloride crystal Each Na+ ion is surrounded by 6 other Cl- ions. Each Cl- ion is surroundedby 6 other Na+ ions 33

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 by For 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.

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 size of the ions The charges of the ions
The shape and form of the crystal lattice depend on several factors: The size of the ions The charges of the ions The relative numbers of positive and negative ions 36

Characteristics of Ionic Bonds
Crystalline at room temperatures Higher melting points and boiling points than covalent compounds Conduct electrical current in molten or solution state but not in the solid state More soluble in polar solvents such as water Brittle Water solutions of ionic compounds are usually electrolytes. 37

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)

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

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

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.

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

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?

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.

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.

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 O C N and Cl D 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.

More Practice 7 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.

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

Lesson 3: Review Covalent Compounds
Friday, December 5, 2014

Covalent Bonding Understandings:
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.

Covalent Bonding Applications 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.

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

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

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

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!

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

Multiple Bond Double bonds are shorter and stronger than single bonds and triple bonds are shorter and stronger than double bonds

Even in the same molecules, double bonds are shorter and stronger!

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+.

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!

Let’s Practice Use the electronegativity values to put the following bonds in order of decreasing polarity. N–O in NO2 N–F in NF3 H–O in H2O N–H in NH3

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?

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!

Bond Continuum

Let’s Practice 10 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

Solutions 10. D 11. C

Lesson 4: Covalent Bonding Structures
Tuesday, December 9, 2014

4.3 Covalent Structures Understandings:
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.

4.3 Covalent Structures Guidance:
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.

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

Steps For Drawing Lewis Dot Structures
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! Draw the skeletal structure to show how the atoms are linked together. Use a pair of crossed, a pair of dots or a straight line between all of the atoms bonded together Add more electron pairs around the atoms until everyone has 8 valence electrons (except for H which only gets 2 electrons!) Check to make sure everyone has an octet! If they do not, make double and/or triple bonds! Make sure you are showing the # of electrons from Step 1

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 Halogens are rarely in the center Hydrogen is never in the center If you have carbon, put that in the center Try to give Group 17 elements one bond, Group 16 elements two bonds, 15 three bonds and 14 four bonds. Try to make the molecule somewhat symmetrical if possible

Easy Example Draw the Lewis structure for the molecule CCl4.
Step 1: Determine # of valence e’s. 12+(7*4) = 32 Step 2: Set up skeleton structure Step 3: Use a straight line to represent each bond. You may also use dots or x’s Step 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

Let’s Practice Draw the following Lewis Dot diagrams CH4 NH3 H2O CO2
HCN OH- (HINT: What does the -1 charge mean?) SO42-

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

Octet Rule Exceptions Element 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

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)

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

Lesson 5: VSEPR Wednesday, December 10, 2014

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!

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.

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?

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

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!

Two Electron Domains The furthers away two electron domains can be from each other is 180 degrees, giving those molecules a linear shape

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.

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°

Summary 2 Linear 180 3 Trigonal Planar 120 1 Bent or T-Shaped 117 4
# of Electron Domains # of Bonded Domains # of Long Pairs Shape Bond Angles 2 Linear 180 3 Trigonal Planar 120 1 Bent or T-Shaped 117 4 Tetrahedral 109.5 Trigonal Pyramidal 107 Bent/T-Shaped 105 5 Trigonal Bipyramidal 90 & 120 6 Octahedral 90

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

Steps To Determine Shape
1 Draw the Lewis structure 2 Count the total number of electron domains on the central atom. 3 Determine the electron domain geometry as follows: 2 electron domains → linear 3 electron domains → triangular planar 4 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.

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) CO32– (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

Polarity Bond polarity is whether the pair of electrons are shared evenly among the two nonmetals covalently bonded together 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 moment A molecule with all nonpolar bonds will always be nonpolar overall 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

Molecular Polarity Molecular polarity depends on:
The polarity of the bonds 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

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.

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.

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)

Let’s Practice Polar Nonpolar
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 Polar Nonpolar

Trickier Question Nonpolar Polar
22 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

Lesson 6: Lewis Dot Structures and Resonance
Wednesday, December 17

Warm-Up Try to draw the following Lewis Dot structures: CO32- O3
What do you notice?

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 different in 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.

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.

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 CO32–.

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

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.

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).

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.)

Two Network Solids To Know!!!!
You need to know the following two network solids for the IB exam: Diamond – in diamond, each carbon atom is covalently bonded to four other carbon atoms in a continuous tetrahedral shape 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

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.

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

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, as in graphite, forming hexagons with bond angles of 120°. Single layer. Electrical Conductivity Good None Semi-conductor Very good Thermal Conductivity Poor Very low Very good; best known Special Properties Soft Hardest known Light and strong Thinnest material ever Uses Pencils! Cutting class Jewels Nanotech Lots of new stuff!

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

Let’s Practice 25 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?).

Answers 25 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.

Lesson 7: Intermolecular Forces
Thursday, December 18

Warm-Up What were the five intermolecular forces we learned about in Regents Chemistry last year?

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.

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.

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.

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

London Dispersion Forces

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

Properties Since London Dispersion Forces are so weak, the nonpolar molecules that are held together by them have: Low melting points Low boiling points High vapor pressure Polar molecules also have London dispersion forces but these are far weaker than some of the other forces at work!

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 molecules attracting each other, generating a force known as a dipole–dipole attraction.

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

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)

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!

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.

Hydrogen bonds explain boiling point trends

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!

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!

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!!

All of these forces are weaker than bonds!!

Let’s Practice Methoxymexane (CH3–O–CH3) boils at a much lower temperature than ethanol (CH3CH2–O–H). Use your knowledge of intermolecular forces to explain why. Put the following molecules in order of increasing boiling point and explain your choice: CH3CHO, CH3CH2OH and CH3CH2CH3

Let’s Practice What is the strongest intermolecular force holding together the following substances? NH3 CCl4 C4H9OH N2 PH3 CO2 CH3F HF Na2O SiO2

Lesson 8: Physical Properties of Covalent Compounds
Friday, December 19

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!

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.

Summary

Let’s Practice

Answers

QUIZ!! 20 minutes!

Lesson 9: Metallic Bonding
Tuesday, December 23

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.

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.

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.

Referred to as “sea of electrons”

Metallic Bond Strength
Metallic bond strength is determined by: The number of delocalized electrons The charge on the cation 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.

Metallic Properties These delocalized electrons give metallic substances very unique properties including: Good electrical conductivity Good thermal conductivity Malleable Ductile High melting points Shiny, lustruous

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

Some Common Alloys Alloy Composition Uses 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

Alloys Small 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 Silicon Magnesium
The electrons in semimetals are much less mobile than in metals, hence they are semiconductors 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

Let’s Practice 31 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.

Lesson 10: HL Expanded Octets and Shapes
Tuesday, January 6

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.

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.

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.

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.

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

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

Shapes See-saw T-Shaped Linear

6 Electron Domains Molecules with six electron domains will position them in an octahedral shape with angles of 90°.

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

Shapes Square Pyramidal Square Planar

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. For example, PCl5 and SF6 are both nonpolar. However, here the molecule may be non-polar even with different atoms bonded to the central atom depending on HOW they are bonded! For example, SBrF5 has a net dipole and is polar; PCl3F2 has no net dipole due to cancellation, so it is non-polar.

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!

Let’s Practice

Answers

Lesson 11: HL Formal Charge
Wednesday, January 7

Warm-Up Pair Share When we can draw more than one Lewis Dot structure for a compound, how do we know which one is correct?

Sulfur Dioxide Which one is correct?
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?

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!

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: (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); (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)

Looking Back at SO2

SO2 We can conclude that structure (ii) where all atoms have a formal charge of zero is the most stable structure for SO2.

Let’s Practice Use the concept of formal charge to determine which of the following Lewis (electron dot) structures for XeO3 is preferred?

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?

Let’s Practice 39 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 SO42–, 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.

Answers

Lesson 12: Sigma and Pi Bonds
Thursday, January 8

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

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!!!

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.

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.

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.

Lesson 13: Hybridization
Friday, January 9

14.2 Hybridization Understandings:
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.

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?

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

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

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!

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.

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.

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.

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.

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).

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.

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.

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

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?

More Practice

Even More Fun!

Lesson 15: Hybridization and Stuctures
Tuesday, January 13, 2015

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

Nitrate Ion - Example NO3+

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!

Lesson 16: Ozone Wednesday, January 14, 2015

Lesson 18: Review Thursday, January 15, 2014

Lesson 18: Review Friday, January 16, 2015

IB Chemistry Year 1 HL The Brooklyn Latin School

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