Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 1: Ionic Bonding Review
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.
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.
Notes:
• ___________________________________________________________ are the electrons in the
outermost principal energy level
• All atoms want to be like the ______________________________________________________,
in Group 18, and have a filled s and p sublevel in the valence shell
Ions:
Groups 1, 2 and
3
Groups 15, 16,
and 17
Group 14
1
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Let’s Practice
What kind of ion is formed by the following elements?
1. Lithium
2. Sulfur
3. Argon
4. Oxygen
5. Nitrogen
Transition Metals:
Unusual Ions:
Naming Ions:
Cations
2
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Anions
Polyatomic
Ions
Memorize!
Ionic Compounds:
Writing Ionic Compounds:
• 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
3
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
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+
2Al
O
3. “Criss cross” the numbers to balance the charges.
3
2
Al
O
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.
•
Write the formula for ammonium phosphate.
1 Write the formula for each of the compounds on the last page (with the
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?
4
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
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
Ionic bonding involves:
1.
2.
3.
Stable octet rule:
Crystal lattice:
5
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
•
•
•
Dies:________________
Guided Notes
In Chemistry, ___________________________________________________, meaning the ions
want to surround themselves with other ions of opposite charge
The ions take on a predictable _____________________________________________
crystalline structure known as the ___________________________________________
The ___________________________________________________________ of the lattice tells
you how many ions each ion in the crystal is surrounded by
Lattice Energy:
• ______________________________________________ is a measure of the strength of
attraction between ions in the lattice of an ionic compound
• ______________________________________________ is higher for ions that are small and
highly charged and weaker for ions that are larger and have a lower charge
Characteristics of Ionic Bonds:
1. Solid at room
temperature
2. Higher melting
points and boiling
points than
covalent
compounds
3. Conduct
electricity in
molten or
solution state
4. Soluble in polar
compounds
6
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
5. Brittle
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
___________________________ 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 _________________________
and an appropriate state symbol to denote
the solvent is used.
• In the case of solvents like oil or hexane,
C6H14, which are _______________________ 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 __________________________________.
Ionic Character:
Bond Continuum:
7
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
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
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.
Lesson 3: Review of 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.
8
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Guidance
• Bond polarity can be shown either with partial charges, dipoles, or vectors.
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
Energetics of Bond Formation:
Lewis Dot Structures:
9
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Multiple Bonds:
• A ________________________________________________ is a sharing of two electrons
• A ________________________________________________is a sharing of two pairs of, or
four, electrons
• A ________________________________________________is a sharing of three pairs of, or
six, electrons
• You can never have a quadruple bond!
Bond Properties
• ________________________________________________: a measure of the distance
between the two bonded nuclei.
• ________________________________________________: usually described in terms of
________________________________________________, 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 Bonds:
• Double bonds are shorter and stronger than single bonds and triple bonds
are shorter and stronger than double bonds
10
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Polar Bonds:
• A bond is______________________________________ when the two elements sharing
electrons share them unevenly
• Elements with different electronegativity values will share electrons
unevenly
• The term ______________________________________________ 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 �–, and the less electronegative atom has
become partially positive or �+.
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
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?
11
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
•
•
•
•
Dies:________________
Guided Notes
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
___________________________________________ 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 __________________________________________________ than
nonpolar bonds!
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
Lesson 4: Covalent Bonding Structures
12
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
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.
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 nonbonding 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:
Steps For Drawing Lewis Dot Diagrams:
13
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Step 1
Step 2
Step 3
Step 4
Step 5
Step 6
Skeleton Help:
1.
2.
3.
4.
5.
Draw the Lewis structure for the molecule CCl4.
14
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Draw the following Lewis Dot diagrams
1. CH4
2. NH3
3. H2O
4. CO2
5. HCN
6. OH- (HINT: What does the -1 charge mean?)
7. SO42Coordinate Covalent Bonds
Octet Rule Exceptions:
Element
Exception
Example
Less than an octet:
15
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
•
Dies:________________
Guided Notes
Atoms that have less than a stable octet are said to be
___________________________________________________________________ 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
16
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 5: VSEPR Theory
VSEPR Theory:
Central Atoms: How many electron domains exist in the central atom of the
following molecules whose Lewis structures are shown?
VSEPR Theory
• The repulsion applies to electron domains, which can be ______________________
_______________________________________________________________________________________
_______________________________________________________________________________________.
• The ______________________________________________________________________________
around the central atom determines the geometrical arrangement of the
electron domains.
• The shape of the molecule is determined by the_______________________________
between the bonded atoms.
• _______________________________________________________________________________________
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 Pair vs. Bonded Pair
• 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!
17
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Two Electron
Domains
Three
Electron
Domains
Four Electron
Domains
18
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
19
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Steps To Determine Shape:
1.
2.
3.
4.
5.
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?
1. tetrahedral (b) bent (c) linear (d) trigonal pyramidal (e) triangular planar
20
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Polarity
• ____________________________________________ is whether the pair of electrons are
shared evenly among the two nonmetals covalently bonded together
o A bond is considered polar if the electronegativity difference between
the two elements is less than 0.4
• ____________________________________________ refers to whether there is an overall
uneven charge distribution in the molecule called a
_____________________________________________________
o A molecule with all nonpolar bonds will always be nonpolar overall
o 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:
1. The polarity of the bonds
2. The shape of the molecule
3. When molecules are symmetrical, even though the bonds are polar,
the charge distributions effectively cancel each other out; see below
•
•
The shapes we have learned so far that can even be nonpolar when the bonds
are polar are _________________________________________________________________________
________________________________________________________________________________________
____________________________________________________________________ molecules can
not be nonpolar with polar bonds.
Net Dipole
Let’s Practice
21 Predict whether the following will be polar or non-polar molecules:
(a) PH3
(b) CF4
(c) HCN
(d) BeCl2
21
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
(e) C2H4
(f) ClF
(g) F2
(h) BF3
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.
Lesson 6: Resonance
Warm-Up: Try to draw the following Lewis Dot structures:
1. CO322. O3
What do you notice?
22
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Delocalized Electrons
Resonance:
Benzene:
•
Notes:
______________________________________
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.
23
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
•
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
1. _______________________________ – in diamond, each carbon atom is covalently
bonded to four other carbon atoms in a continuous tetrahedral shape
2. _______________________________ – 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
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
24
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
25
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
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?).
Lesson 7: 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.
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
26
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
London
Dispersion
Forces
DipoleDipole
27
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Hydrogen
Bonds
Ion-Dipole
28
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
London Dispersion Forces Strength:
Since London Dispersion Forces are so weak, the nonpolar molecules that are held
together by them have:
1. Low melting points
2. Low boiling points
3. High vapor pressure
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 de Waals’ Forces:
29
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Why are hydrogen bonds so strong?
Boiling Point Trends:
Water and Biochemistry
30
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Let’s Practice
1. 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
3. What is the strongest intermolecular force holding together the following
substances?
1. NH3
2. CCl4
3. C4H9OH
4. N2
5. PH3
6. CO2
7. CH3F
8. HF
9. Na2O
10. SiO2
31
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 8: Physical Properties of Covalent Compounds
Solubility
Electrical
Conductivity
32
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Let’s Practice
33
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 9: 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.
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 Bonds:
Metallic bond strength is determined by:
1.
2.
3.
34
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Metallic Properties:
Alloys:
35
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Common alloys:
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.
36
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 10: Expanded Octets and Shapes
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 nonbonding 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.
37
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Expanded Octets:
 While the octet is the most common arrangement for atoms when entering
________________________________________________ utilizing the
________________________________________________
• Since the atom needs to use the d sublevel to expand its octet, only elements
in ________________________________________________and below can have expanded
octets
• Elements such a sulfur and phosphorus can create compounds with 5 or
6 electrons.
5 Electron Domains:
6 Electron Domains:
38
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Molecular Polarity:
39
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Let’s Practice
40
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 12: Formal Charge
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
Way 1
Way 2
Formal Charge
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)
41
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Looking Back at 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:
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?
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.
42
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 13: Sigma and Pi Bonds
Molecular Orbital:
Sigma
Bonds
Pi Bonds
Pictures
Sigma Bonds
•
Pi Bonds
Pi bonds are weaker than sigma bonds, as their electron density is further
from the nucleus.
43
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
Lesson 14: 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
Carbon Atom:
• Step 1: A process known as ______________________________ 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
•
•
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 _______________________________________________________ which are
the same as each other, but different from the original orbitals.
This mixing of orbitals is known as _____________________________________________
44
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
•
Dies:________________
Guided Notes
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.
Sp3
Sp2
Sp
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.
Molecular Shapes
• tetrahedral arrangement ↔ sp3 hybridized
• triangular planar arrangement ↔ sp2 hybridized
• linear arrangement ↔ sp hybridized
45
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
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?
46
Nomen:____________________________________________________
IB Chemistry Year 1 HL – Unit 4 Bonding (Topic 4 and 14)
Dies:________________
Guided Notes
47