Source: Chapter 1, Dr Uyeda lecture slides
Tags: functional group, alkane, alkene, alkyne, alcohol, ether, amine, aldehyde, ketone, carboxylic acid, ester, amide, hybridisation, sp3, sp2, sp, sigma bond, pi bond, molecular orbital
Difficulty: Intermediate Prerequisites: Parts 1 and 2 of these Chapter 1 notes (electron configuration, Lewis structures, VSEPR). You should be able to draw Lewis structures and predict molecular geometry before tackling hybridisation.
Functional groups are the reactive parts of organic molecules. Learning to recognise them is like learning the alphabet before reading: every reaction you study later in the course will involve one or more of these groups. Hybridisation and molecular orbital theory explain why carbon can form four bonds, why double bonds are planar, and why triple bonds are linear. This section bridges the Lewis/VSEPR picture (a useful simplification) with the orbital picture (a more accurate description of where electrons are).
Functional groups are specific atom arrangements that determine a molecule's reactivity and properties; the suffix of an organic molecule's name usually tells you which functional group it carries. Carbon atoms hybridise their s and p orbitals to form sp³ (tetrahedral, 109.5°), sp² (trigonal planar, 120°), or sp (linear, 180°) hybrid orbitals. Sigma (σ) bonds form from head-on orbital overlap; pi (π) bonds form from side-on overlap of unhybridised p orbitals and are generally weaker than σ bonds.
Functional group
A specific arrangement of atoms within a molecule that determines how the molecule reacts and what properties it has. Think of it as the part of the molecule where the action happens.
Alkane
A hydrocarbon containing only carbon-carbon single bonds and the maximum number of C–H bonds. The basic framework of all organic molecules. Can be acyclic (a chain, e.g. butane) or cyclic (a ring, e.g. cyclohexane).
Alkene
A hydrocarbon containing at least one C=C double bond. Named with the suffix "-ene" (e.g. ethylene).
Alkyne
A hydrocarbon containing at least one C≡C triple bond. Named with the suffix "-yne" (e.g. ethyne, also called acetylene).
Alcohol
A molecule containing an O–H (hydroxyl) group bonded to carbon. Named with the suffix "-ol" (e.g. ethanol).
Ether
A molecule with an oxygen atom bonded to two carbon atoms by single bonds (C–O–C). Named with "ether" (e.g. diethyl ether).
Amine
A molecule containing a nitrogen atom bonded to one, two, or three carbon or hydrogen substituents (NR₃). Named with the suffix "-amine" (e.g. methylamine).
Aldehyde
A molecule containing a carbonyl group (C=O) with at least one hydrogen substituent on the carbonyl carbon. Named with the suffix "-al" (e.g. ethanal, also called acetaldehyde).
Ketone
A molecule containing a carbonyl group (C=O) flanked by two carbon substituents. Named with the suffix "-one" (e.g. acetone).
Carboxylic acid
A molecule containing a carbonyl group bonded to a hydroxyl group (–COOH). Named with the suffix "-ic acid" (e.g. acetic acid). Loses a proton to form a carboxylate anion (–COO⁻).
Ester
A molecule containing a carbonyl group bonded to an –OR group (where R is carbon). Think of it as a carboxylic acid where the acidic H has been replaced by a carbon group. Esters often have pleasant, fruity smells.
Amide
A molecule containing a carbonyl group bonded to an –NR₂ group. Named with the suffix "-amide" (e.g. acetamide). The nitrogen in an amide is trigonal planar (not pyramidal) due to resonance.
Hybridisation
The mixing of atomic orbitals (s and p) on the same atom to form new, equivalent hybrid orbitals that are better suited for bonding. The number of starting orbitals always equals the number of hybrid orbitals produced.
sp³ hybridisation
One s orbital and three p orbitals combine to form four equivalent sp³ hybrid orbitals arranged tetrahedrally (109.5° apart). Carbon in methane and ethane is sp³.
sp² hybridisation
One s orbital and two p orbitals combine to form three equivalent sp² hybrid orbitals in a trigonal planar arrangement (120°), with one unhybridised p orbital remaining perpendicular to the plane. Carbon in ethylene is sp².
sp hybridisation
One s orbital and one p orbital combine to form two equivalent sp hybrid orbitals arranged linearly (180°), with two unhybridised p orbitals remaining. Carbon in acetylene is sp.
Sigma (σ) bond
A bond formed by head-on (end-to-end) overlap of orbitals along the internuclear axis. All single bonds are σ bonds. The first bond in any double or triple bond is also a σ bond.
Pi (π) bond
A bond formed by side-on (lateral) overlap of parallel p orbitals above and below the internuclear axis. Pi bonds are weaker than σ bonds because the overlap is less efficient. A double bond = 1 σ + 1 π; a triple bond = 1 σ + 2 π.
Bonding molecular orbital
A molecular orbital formed when two atomic orbitals combine constructively (in phase). Lower in energy than the original atomic orbitals. Electrons in bonding orbitals stabilise the molecule.
Anti-bonding molecular orbital
A molecular orbital formed when two atomic orbitals combine destructively (out of phase). Higher in energy than the original atomic orbitals, with a node between the nuclei. Electrons here destabilise the molecule. Anti-bonding destabilisation is greater than bonding stabilisation, which is why He₂ does not form.
Functional groups determine how molecules interact with each other and how they react. The carbon-hydrogen skeleton (the alkane framework) is relatively inert; it is the functional groups attached to that skeleton that give each molecule its characteristic chemistry. The suffix in an organic compound's name usually identifies its functional group.
Functional group | Defining feature | Naming suffix | Example |
|---|---|---|---|
Alkane | C–C and C–H single bonds only | -ane | Butane |
Alkene | C=C double bond | -ene | Ethylene |
Alkyne | C≡C triple bond | -yne | Ethyne (acetylene) |
Alcohol | O–H group on carbon | -ol | Ethanol |
Ether | C–O–C linkage | ether | Diethyl ether |
Amine | N bonded to H and/or C | -amine | Methylamine |
Aldehyde | C=O with at least one H on carbonyl C | -al | Formaldehyde |
Ketone | C=O flanked by two C groups | -one | Acetone |
Carboxylic acid | C=O bonded to O–H | -ic acid | Acetic acid |
Ester | C=O bonded to O–R | -ate (as salt/ester) | Ethyl acetate |
Amide | C=O bonded to N | -amide | Formamide |
Acyclic molecules are open chains (e.g. butane: four carbons in a row).
Cyclic molecules form rings (e.g. cyclohexane: six carbons in a ring).
In line drawings, each vertex represents a carbon atom, and hydrogens on carbon are not drawn (you work them out by counting: carbon wants 4 bonds total).
In a simple amine (e.g. methylamine), nitrogen has a lone pair and three bonds, giving trigonal pyramidal geometry. In an amide, the lone pair on nitrogen participates in resonance with the adjacent C=O, delocalising into a π system. This gives nitrogen sp² character and a trigonal planar geometry. The resonance structure shows a C–N double bond, an N with a positive formal charge, and an O with a negative formal charge.
Amide bonds (peptide bonds) connect amino acid residues in proteins. Each amino acid has an amine group and a carboxylic acid group; the amide linkage forms between these. Amino acid side chains carry various functional groups (alcohols, amines, carboxylates, amides, thiols, aromatics), which determine the protein's properties.
Drug molecules like simvastatin (Zocor) contain multiple functional groups within one structure: alcohols, esters, and alkenes. Being able to spot each group is a core skill. Drugs often work by mimicking the shape and functional groups of a natural substrate, fitting into the same binding site on a protein.
Carbon's ground-state electron configuration (1s² 2s² 2px¹ 2py¹) has only two unpaired electrons, which would suggest carbon should form only two bonds. In reality, carbon consistently forms four bonds. Hybridisation explains this: the 2s and 2p orbitals mix to create new, equivalent hybrid orbitals that each hold one electron, ready to bond.
The number of atomic orbitals you start with always equals the number of hybrid orbitals you produce. No orbitals are created or destroyed.
Mix 1 s orbital + 3 p orbitals → 4 sp³ hybrid orbitals.
The four sp³ orbitals point to the corners of a tetrahedron (109.5° apart).
Each sp³ orbital is asymmetric: one large lobe for bonding, one small back lobe.
Example: methane (CH₄). Each C–H bond is a σ bond formed by overlap of a C sp³ orbital with an H 1s orbital.
Example: ethane (C₂H₆). The C–C bond is a σ bond from overlap of two sp³ orbitals. The six C–H bonds each come from sp³ overlapping with H 1s.
Mix 1 s orbital + 2 p orbitals → 3 sp² hybrid orbitals + 1 unhybridised p orbital.
The three sp² orbitals lie in a plane at 120° to each other. The remaining p orbital is perpendicular to that plane.
Example: ethylene (H₂C=CH₂). The C–C σ bond forms from two sp² orbitals overlapping head-on. The C=C π bond forms from the two unhybridised p orbitals overlapping side-on, above and below the plane. Each C–H bond is sp²–1s.
Mix 1 s orbital + 1 p orbital → 2 sp hybrid orbitals + 2 unhybridised p orbitals.
The two sp orbitals point in opposite directions (180°). The two remaining p orbitals are perpendicular to each other and to the sp axis.
Example: acetylene (HC≡CH). The C–C σ bond forms from two sp orbitals. The two C–C π bonds form from the two pairs of unhybridised p orbitals. Each C–H bond is sp–1s.
Groups bonded | Hybridisation | Geometry | Bond angle | σ bonds | π bonds |
|---|---|---|---|---|---|
4 | sp³ | Tetrahedral | 109.5° | 4 | 0 |
3 | sp² | Trigonal planar | 120° | 3 | 1 |
2 | sp | Linear | 180° | 2 | 2 |
Count the number of groups (atoms + lone pairs) attached to the atom of interest. 4 groups = sp³; 3 groups = sp²; 2 groups = sp. A double bond counts as one group; a triple bond counts as one group.
Formed by head-on overlap of orbitals along the axis connecting the two nuclei.
Can form from s–s overlap (e.g. H₂: two 1s orbitals), sp³–s overlap (e.g. C–H in methane), sp³–sp³ overlap (e.g. C–C in ethane), sp²–sp² overlap (e.g. C–C σ component of ethylene), or sp–sp overlap (e.g. C–C σ component of acetylene).
σ bonds have strong overlap and are the strongest type of covalent bond.
Every single bond is a σ bond. In a double bond, one bond is σ and one is π. In a triple bond, one bond is σ and two are π.
Formed by side-on overlap of two parallel, unhybridised p orbitals.
Electron density sits above and below the internuclear axis, with a node on the axis itself.
π bonds are weaker than σ bonds because the lateral overlap is less efficient.
π bonds are present only in double and triple bonds.
When two H atoms (each with one 1s electron) combine, both electrons go into the bonding molecular orbital, stabilising the molecule: H₂ forms. When two He atoms (each with two 1s electrons) try to combine, two electrons fill the bonding orbital and two fill the anti-bonding orbital. Because anti-bonding destabilisation is slightly greater than bonding stabilisation, there is no net energy benefit. He₂ does not form a stable bond.
Bond strength correlates with the degree of orbital overlap. Head-on (σ) overlap is more effective than side-on (π) overlap, so σ bonds are stronger. Among σ bonds, sp orbitals (50% s character) overlap more effectively than sp³ orbitals (25% s character), making C–H bonds in acetylene stronger and shorter than C–H bonds in ethane.
Functional groups dictate a molecule's physical and biological properties. Ethanol (an alcohol) is water-soluble and drinkable; diethyl ether (an ether with the same atoms rearranged) is water-immiscible and was historically used as an anaesthetic. Esters give fruits their characteristic smells: ethyl acetate is the dominant scent compound in nail polish remover. Amide bonds hold proteins together; every enzyme, antibody, and structural protein in your body is a chain of amino acids linked by amide bonds. Drugs like simvastatin work by mimicking the shape and functional groups of a natural substrate (HMG-CoA), fitting into the active site of the target enzyme.
Students often think a double bond is simply "two single bonds." A double bond consists of one σ bond and one π bond, which have different geometries and strengths. The π bond restricts rotation around the bond axis.
Students sometimes confuse hybridisation with the number of bonds. Hybridisation is determined by the number of groups (atoms or lone pairs) attached, not the total number of bonds. A double bond counts as one group.
Students frequently assume the nitrogen in an amide is pyramidal, like a simple amine. It is not. Resonance with the adjacent carbonyl gives the amide nitrogen trigonal planar geometry.
Students may think that π bonds are always between carbon atoms. Pi bonds can form between any two atoms that have parallel p orbitals, including C=O, C=N, and N=O.
⚠️ You will be given complex molecules and asked to identify every functional group present. Practise with real drug structures.
⚠️ Expect questions asking for the hybridisation of every carbon (and sometimes nitrogen or oxygen) in a molecule. The shortcut: count groups attached to the atom.
⚠️ Know how to count σ and π bonds in any molecule. A common exam question gives a structure and asks: "How many σ bonds? How many π bonds?"
⚠️ Be able to explain why the amide nitrogen is planar using resonance structures. This concept returns in peptide and protein chemistry.
⚠️ Understand why He₂ does not form from a molecular orbital perspective. This illustrates the principle that anti-bonding destabilisation outweighs bonding stabilisation.
Fill in the blank: an alcohol contains an ___ group bonded to carbon.
True or false: a ketone has at least one hydrogen directly on the carbonyl carbon.
Fill in the blank: an sp²-hybridised carbon has ___ unhybridised p orbital(s).
True or false: a triple bond contains two σ bonds and one π bond.
Fill in the blank: the bond angle around an sp-hybridised carbon is ___°.
Answers: 1. O–H (hydroxyl). 2. False (that describes an aldehyde; a ketone has two carbon substituents). 3. One. 4. False (a triple bond contains one σ bond and two π bonds). 5. 180.
Q: Identify all functional groups in ethyl acetate (CH₃COOCH₂CH₃).
A: Ethyl acetate contains one ester group (a carbonyl bonded to an –OR group). The rest of the molecule is an alkane framework.
Q: What is the hybridisation of each carbon in the molecule H–C≡C–CH=CH–CH₃?
A: Working left to right: C1 (triple bond) = sp. C2 (triple bond) = sp. C3 (double bond + single bonds, 3 groups) = sp². C4 (double bond + single bonds, 3 groups) = sp². C5 (four single bonds, 4 groups) = sp³.
Q: How many σ bonds and how many π bonds are in acetylene (HC≡CH)?
A: Three σ bonds (two C–H and one C–C) and two π bonds (both in the triple bond).
Q: Explain why the nitrogen atom in formamide (H₂NCHO) is trigonal planar rather than trigonal pyramidal.
A: In formamide, the lone pair on nitrogen can delocalise into the adjacent C=O through resonance, creating partial C=N double-bond character. This resonance gives nitrogen sp² hybridisation and trigonal planar geometry, rather than the sp³/pyramidal geometry of a simple amine.
Q: A molecule has the structure O=C=O. State the hybridisation of the carbon, the number of σ bonds, and the number of π bonds.
A: Carbon is sp (two groups attached). There are 2 σ bonds (one to each oxygen) and 2 π bonds (one in each C=O double bond).
Functional groups are the organising principle for the rest of the course: each subsequent chapter focuses on the reactions of one or two specific groups (alkene reactions, alcohol reactions, carbonyl chemistry, etc.). Hybridisation connects directly to how you understand reaction mechanisms, because the geometry around a reacting atom determines which face a reagent can attack. Pi bonds are the sites where most addition reactions occur in alkene and alkyne chemistry.
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