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Summary CHE1502 Study Guide and Textbook Mindmaps
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- University Of South Africa (Unisa)
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STRUCTURE AND BONDING 1 1STRUCTURE AND BONDING 2 2
STRUCTURE AND BONDING OF COVALENT BONDS 3
SIGMA (σ) and Pi (π)
1. Structure and Bonding 4
BONDS
RESONANCE 5
STRUCTURE AND BONDING 3 6
Theme 1: Fundamental Concepts in Organic Chemistry
STRUCTURE AND BONDING 4 7
POLARITY 8
INTERMOLECULAR FORCES 9
2. Acids and Bases; Functional Groups ACIDS AND BASES 10
A BORDER WRITTEN IN ALL CAPS, IS COLOURED IN, AND IS SQUARE IS
EQUILIBRIUM POSITIONS OF ACID-BASE REACTIONS 11
THE MAIN TOPIC OF THE PAGE. THIS IS YOUR STARTING POINT. THE
COLOUR OF THIS SQUARE TELLS YOU WHAT CHAPTER YOU ARE ON
REACTIONS OF ACIDS AND BASES 12
The clouds lead to the information ALKANES 1 13
of the topic if you follow its lines
3. Structure and Properties of Alkanes ALKANES 2 14
A cloud-shaped, coloured in
This information does not have a border shows what topic you are
boarder and is not coloured in, but reading ALKANES 3 15
the words are in colour so that your Borders: The borders, shapes, and
.
brain links this information with the colouring in mean different things:
topic name CHEMICAL REACTIONS 1 16
A dotted boarder shows something important that you need to remember FREE RADICAL CHAIN REACTION MECHANISM IN THE HALOGENATION OF ALKANES 17
4. Chemical Reactions
A rounded border that is coloured in shows you an CHEMICAL REACTIONS 2 18
example of something you have just learned
CHEMICAL REACTIONS 3 19
Writing that is underlined gives you questions to practice on your own
STEREOCHEMISTRY 1 20
5. Introduction to Stereochemistry
STEREOCHEMISTRY 2 21
"+" means positive "-" means negative
ORGANIC CHEMISTRY ALKYL HALIDES 1 22
"1x" means "one of"; "2x" means
"two of" etc ALKYL HALIDES 2 23
CHEMISTRY IB
Symbols for elements from the ALKYL HALIDES REACTIONS 24
periodic table are used, for
CHE1502 (UNISA) Theme 2: Saturated Organic Compounds 6. Alkyl Halides
example "H" means "hydrogen",
and "C" means "carbon" SN2 REACTION MECHANISM 25
The molecular formula, H2, should FACTORS INFLUENCING THE SN2 REACTION MECHANISM 26
be written as: (this is the same for
all molecular formulas) SN1 REACTION MECHANISM 27
ALKENES 1 37 SATURATED COMPOUNDS CONTAINING O AND N 28
H3O+ should be written as: I cannot write superscripts or
subscripts, so be aware of the
correct way of writing things PHYSICAL PROPERTIES OF ALKENES 38 ALCOHOLS 1 29
There are some symbols and abbreviations used in these mind
maps. Do not worry! They are all listed here and only ones
Fractions are written as "1/2" that are easy to understand are used. SYNTHESIS OF ALKENES BY ELIMINATION 39 ALCOHOLS 2 30
which can also be written as:
SYNTHESIS OF SYNTHESIS AND REACTIONS OF ALCOHOLS 31
ALKENES BY
ELIMINATION - 40
EXAMPLES AND 7. Saturated Compounds Containing Oxygen and Nitrogen REACTIONS OF ALCOHOLS 32
is written as "p-x-" 8. Alkenes
MECHANISMS
REACTIONS OF ALCOHOLS WITH HYDROHALIC ACIDS 33
is written as "sp2" REACTIONS OF ALKENES 1 41
ALCOHOLS EXAMPLES AND MECHANISMS 34
is written as "1s2 2s 2p2 REACTIONS OF ALKENES 2 42
ETHERS 35
is written as "109.4o" or "109.4 degrees" REACTIONS OF ALKENES 3 43
AMINES 36
REACTIONS OF ALKENES 4 44
is written as "pKa"
ALKYNES 1 45
is written as "Sn1"
Theme 3: Unsaturated Organic Compounds
9. Alkynes ALKYNES 2 46
ADDITION REACTIONS OF ALKYNES 47
CARBONYL
Please feel free to contact me should you feel there is COMPOUNDS AND 48
Credits: Credit for all information and images
any unfair or illegal use of the content found in this THEIR DERIVATIVES
in this study guide goes to the UNISA
guide. I do not intend any plagiarism, rather, I wish to
CHE1502 Study Guide and to the
use the work of others as a guide to help others
recommended textbook, Organic Chemistry 9e ALDEHYDES AND KETONES 1 49
understand the content in a "bigger picture" context. I
by L.G. Wade
will gladly remove any work that is seen to be unethical.
ALDEHYDES AND KETONES 2 50
10. Carbonyl Compounds and their Derivatives CARBOXYLIC ACIDS 1 51
CARBOXYLIC ACIDS 2 52
CARBOXYLIC ACID
53
DERVIATIVES 1
CARBOXYLIC ACID
54
DERVIATIVES 2
, A filled outer electron shell has
Atoms join with other atoms to form bonds to
the configuration of a noble gas
have a completely filled outer shell of electrons
(noble gas configuration)
High melting points
Bonding: The combination of atoms - each atom gets a filled outer shell of electrons
Inorganic Compounds Solutions in water conduct electricity
Ionic Bonding: Electrostatic attraction
between ions with opposite charges Ionic bonds
Atoms can bond via 2 modes
Introduction
Covalent Bonding: Sharing of electrons Low-melting points
The movement of valence electrons from a region of high Chemical Reactions: Bond Organic Compounds Solutions in water do not conduct electricity
electron density to a region of lower electron density formation or breakage
Covalent bonds
Plays a big role in in the nature of a Electronegativity: The tendency of an atom to attract a pair of
chemical reaction and how it reacts bonding electrons to itself in a molecule (see the blue topic below)
Lose 1 electron to become a Atoms have protons (+ charge, found in the nucleus), neutrons (no charge, found in
cation with a filled outer shell Group 1A elements have nucleus), and electrons (- charge, found in orbitals around the nucleus)
1s outer electron shell
Chemical Bond Formation Types of orbitals: s, p, d, f - only
the s and p-orbitals are important
Gain 1 electron to become an (for this course) Orbitals are found in
anion with a filled outer shell Group 7A elements have different electron shells
7 valence electrons Ionic Bonding
Example: The electrostatic attraction between the sodium cation Ionic Bond: Electrostatic attraction between
(positive ion) and chlorine anion (negative ion) to form NaCl (table salt) a sodium cation and chlorine anion The number of orbitals increases in
Each higher shell has higher
shells that are larger and further
energy and is larger in size
away from the nucleus
Non-metals (more electronegative) Metals (electropositive) -
- gain electrons lose electrons
Second electron shell contains 1x
s-orbital and 3x p-orbitals
Covalent Bonding: Sharing of electrons so each atom has a noble gas configuration (octet) First electron shell contains 1x
s-orbital
The 3 p-orbitals point in the x-, y-,
Compounds formed by non-metals (the molecules in a covalent bond are all non-metals) and z- axes
Covalent Bonding
Example of a covalent bond: 2 H atoms bond Each orbital has maximum 2 This is Pauli's exclusion
Iodine can also bond with another I to form I2:
to form H gas (H2) electrons principle
The 3x p-orbitals have the
2p-orbitals are higher in
same energy - they are
energy than 2s-orbitals
Principles of Atomic Structure called degenerate orbitals
STRUCTURE AND
BONDING 1
Carbon can form many compounds because of
the different combinations of s- and p-orbitals
Electron Configuration: The Valence Electrons: Electrons in the
distribution of electrons of an atom outermost shell of an atom
Hund's Rule: Electrons will occupy all orbitals of
the same energy before pairing begins
The electron configurations of the first 10 elements of the periodic table:
The attraction that an atom has on
the bonding electrons and
electronegativity value of an atom
Non-polar Bond: Bonding electrons are
attracted with the same intensity by each atom
Polar Covalent Bond: Unequal sharing of negatively charged bonding electrons by two atoms
Trend of electronegativity: Increases as you go
higher and to the right of the periodic table
Electronegativity
δ + and δ - show which atom is more
The "δ" means "slightly"
electronegative ( δ -) or electropositive ( δ +)
Practice: Use electronegativities to predict the direction of the
Symbol for electronegativity: dipole moments of the following bonds:
1
Electronegativity = >1.7 - ionic Electronegativity = 0.4-1.7 - polar Electronegativity = <0.4 - non-polar
Example: The carbon-fluorine bond is polar
The CH3F bond is written as: because the difference in electronegativity is 1.5
, Help predict how bonds will break or how a reaction will take place
Cl: [Ne} 3s2 3p5 O: [He] 2s2 2p4 Electronic configurations in each atom in ClO4- :
Drawing: How structures are drawn showing the bonds
Number of electrons available for bonding = Number of valence electrons of Cl and electrons not involved in bonding
+ Number of valence electrons of O + number of charge (i.e. -1 for ClO4-) Example: Draw the Lewis structure
of a molecule with the molecular
formula of CH3CONH2 Each atom has an octet in the drawing of a
= 7 (1xCl) + (4 x 6) (4xO) + 1 (charge) = 32 electrons molecule
Therefore the Lewis structure is: Each valence electron of an atom is
indicated with a "•"
Example: Draw the Lewis
structure of the ClO4 ion
A pair of bonding atoms is
indicated with a "--"
Single Bond: Two electrons shared between atoms (see
Lewis Structures
the purple example), denoted with "--"
The central Cl atom forms 4 sigma (single) bonds with each O
Drawing Lewis structures Double Bond: To get an octet structure, two pairs of electrons
may need to be shared between 2 atoms, denoted with "="
Cl contributes 7 electrons - the central Cl atom that has the 7 +1 electron
(from the -1 charge) available for bonding with O
Triple Bond: Three pairs of electrons are shared between 2 atoms to
get an octet, denoted with "≡"
Carbon can have a maximum of 4 bonds Hydrogen will always have 1 bond
When drawing Lewis structures it may be necessary to give + or - charges
to atoms in a molecule or ion, these are called Formal Charges Atoms that are part of stable compounds
have more or less than 8 electrons around it
Exceptions to the Octet Rule
Formal Charge (FC) = Group number of atom - number of lone pair electrons
- half the number of bonding electrons Example: Boron often has 6 Boron's empty orbital makes
valence electrons - BH3 it useful in organic reactions
FC = 4 (group number) - 0 (number of
lone pairs) - half (6) (bonding electrons
Example: Determine the FC on the C marked with a * in this structure Oxygen has 6 Hydrogen has 1 valence So, there are 8 valence
valence electrons electron (2 H atoms in water) electrons between them
= 4-3 = +1 = FC on the C*
When the hydrogens share an electron with the oxygen it has a full outer shell
Formal Charges Example: A water molecule
When the 8 electrons surround
the oxygen it has a full outer shell
The Lewis Structure of water
Redraw the structure showing the
lone pair (non-bonding electrons): Lone Pairs / Non-Bonding Pairs: The unshared electrons (the dots on the O)
Example: Determine the FC on N
in the structure below:
These are complex ions where 1 of the ions
Organic compounds may exist as ionic compounds
has atoms covalently bonded
FC = 5 - 2 - 1/2(6) = 3-3 = 0
STRUCTURE AND Example: CH3ONa - Sodium Methoxide - used as a
The FC on the N is 0 BONDING 2 Ionic Structures reagent in organic reactions
Half of 2 shared Ionic bond between positively charged Na and
electrons is 1 electron negatively charged O - atoms in the methoxide
(OCH3) are covalently bonded
Each H atom has 1 bonding
1 valence electron is what H
pair of electrons (2 shared
needs to become neutral
electrons)
H atoms with 1 bond are Example: Determine the FC on This is a neutral compound where the individual atoms are formally charged - B and N have 8/2 = 4 electrons
formally neutral (FC=1-0-1=0) each atom in methane (CH4) the Lewis structure shows that N and B have 4 shared bonding pairs of electrons contributing to their charges
Example: Determine the FC on the
hydronium ion (H30+)
Half of 8 shared Example: Determine N needs 5 valence electrons to be
FC = 5 - 0 - 1/2(8) = +1
electrons is 4 electrons the FC on H3N-BH3 neutral so it has a FC of +1
In the Lewis structure there are 8 B needs 3 valence electrons to be
4 electrons are what C The C atom has 4 bonding electrons: 6 from the O and 2 from the Hs FC = 3 - 0 - 1/2(8) = -1
needs to become neutral pairs of electrons (8 electrons) neutral so it has a FC of -1
- minus 1 because the ion has a + charge
C is formally neutral whenever it
Each H has 1 bond
has 4 bonds (FC=4-0-1/2(8) = 0) C and N have 4 shared
and is formally neutral
pairs of bonding electrons
O is surrounded by an octet
C needs 4 valence
with 6 bonding electrons and
electrons to be neutral
2 non-bonding electrons
Practice: Draw Lewis structures for the following Example: Determine the FC on
compounds and ions, showing formal charges {H2CNH2}+ N needs 5 valence electrons to be
Half the bonding electrons (3) plus FC = 5 - 0 - 4 = +1
neutral so it has a +1 formal charge
the non-bonding electrons (2) = 5
O needs 6 valence electrons to The compound can be drawn
with this Lewis structure: The C atom has 3 bonds 6/2 = 3 electrons,
be neutral, so it has a FC of +1 FC = 4 - 0 - 1/2(6) = +1
with 6 bonding electrons so C needs 1 more
Practice: Draw Lewis structures for:
FC= 6 - 2 -1/2(6) = +1 N has 6 bonding electrons
6/2 + 2 = 5, so the N is uncharged FC = 5 - 2 - 1/2(6) = 0
and 2 non-bonding electrons
2
, Notation to show the direction of movement of electrons during bond
formation or breakage which involve the movement of 1 or 2 electrons
Example: Use arrows to
Answer: show how the bonds are
broken or formed in the Half-headed = movement of 1 electron
The Curved-Arrow Formalism
Each half of the arrow-head represents a single electron
following conversion:
Double-headed = movement of 2 electrons
2 elements (non-metals) of different electronegativity react to form a bond
A curved line represents movement of electron(s)
Each atom contributes one electron to the bond pair
Forward reaction movement of electrons shown with half-headed arrows that meet between the 2 atoms
Forms between atoms of similar electronegativity
When bonded together as a molecule the atom with the higher electronegativity pulls the bond pair
electrons towards itself - making it electron-rich, and the other atom electron-poor Equal sharing means equal ownership at any given time
The bond becomes polarised with a positive and a negative pole Bond pair electrons are equally shared between the 2 atoms = non-polar
Diatomic molecules (H2, N2, O2, X2
(X=halogen)) are held together by covalent
The atom of high electronegativity gets a slightly negative charge - represented with "δ‾" bonds The bond-pair electrons belong equally to each atom e.g. in H2, we say
that each H has 2 atoms even though they are being shared
STRUCTURE AND BONDING
The atom of less electronegativity gets a slightly positive charge - represented with "δ+"
OF COVALENT BONDS Each element in the molecule has 8 electrons around itself
2 of the same halogen atoms (represented by X,
A curved arrow begins at an unshared pair of
where X can be: F, Cl, Br, I) react to form X2 When counting electrons around each element,
electrons or a covalent bond - the curved arrow
There are 2 electrons around the H (the moves in the direction where there is a need for each bond pair belongs to both atoms equally
bond pair electrons), and 8 electrons around electrons (e.g. to an atom that is electron deficient)
the Cl (3 pairs of non-bonding electrons and "-" indicates a bond of 2 electrons
1 bond pair of electrons) Formation and Breakage of
Polar Covalent Bonds
When a hydrogen or halogen molecule is formed, each atom has
an unpaired electron that they share to form a covalent bond
Formation and Breakage of Non-Polar
Covalent Bonds
Bond breaks: both bond pair electrons are taken away by Illustrate movement of these electrons during
Bond formation
the atom of high electronegativity - shown with a double-headed bond formation with half-headed curved
curved arrow to show the movement of 2 arrows moving from the unpaired electrons to
Both bond pair electrons are taken away electrons from the middle of the bond to the atom of high the area between the 2 atoms to form a bond
by the more electronegative atom = electronegativity (forward reaction)
heterolytic bond cleavage / heterolysis
Example: Use arrows to show how When a hydrogen or halogen molecule breaks / cleaves
the bonds are broken or formed in into 2 elements, each atom gains a single electron -
the following conversion: illustrate movement of these electrons with half-headed
Answer: curved arrows moving from the middle of the bond to each atom
Bond breakage
When the bond breaks the atom of higher electronegativity gains a pair of electrons and becomes Homolysis / homolytic bond cleavage: This type of bond cleavage
negatively charged (an anion); and the less electronegative atom becomes positively charged (a cation)
Iodine (I2) consists of 2 identical I atoms
When an anion and a cation of a non-metal react, the anion (more electron rich) attacks the cation
The bond formation and bond breakage can be shown in 1 equation
(more electron poor) to form a polar covalent bond - the anion donates an electron pair to the cation
because the formation is reversible
Examples and Practice: In the examples and practices below, draw the Lewis structures of the reactants and
products, determine which species are acting as electrophiles (acids) and which are acting as nucleophiles (bases), Example 2 Practice 1
and use the curved arrow formalism to show the movement of electron pairs in these reactions
A bond has formed between the O
of the CH3-O- group and the C=O
C atom
This is a proton transfer from HCl to the C=O to the group of acetaldehyde
Acetaldehyde acts as the base (proton acceptor) HCl acts as the acid (proton donor) The CH3-O- group (nucleophile) donates the electrons to form the Practice 2
new bond to acetaldehyde (electrophile)
1 arrow must show the movement of electrons from the electron pair
donor (nucleophile/base) to the electron pair acceptor (electrophile/acid)
Another curved arrow must show the electrons that once formed
the H-Cl bond leaving with Cl to make a chloride ion (Cl-) Practice 3
Example 1
3
The resonance forms of the product show
The + charge is delocalised over the C and
that a pair of electrons can be moved
O atoms with most of the + charge on the O
between the O atom and the C=O pi bond
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