AOL 1 : Chemistry of Life
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14 Terms
Describe the structure of an atom.
Overall organization
An atom consists of a dense, central nucleus surrounded by a cloud of electrons.
Most of the atom’s volume is empty space occupied by the electron cloud, while nearly all of its mass is in the nucleus.
Subatomic particles
Protons have a charge of +1 and a mass of about 1 atomic mass unit. They are found in the nucleus.
Neutrons have no charge and a mass of about 1 atomic mass unit. They are also found in the nucleus.
Electrons have a charge of −1 and a negligible mass. They move around the nucleus in regions called orbitals, which are organized into energy levels (shells).
Key numbers
The atomic number is the number of protons and defines the element.
The mass number is the number of protons plus neutrons.
In a neutral atom, the number of electrons equals the number of protons, so the charges cancel.
Variations
Isotopes are atoms of the same element with different numbers of neutrons.
Ions are atoms that have gained or lost electrons. Losing electrons gives a positive cation, and gaining electrons gives a negative anion.
Valence electrons
These are the electrons in the outermost shell.
They determine how an atom bonds, because atoms tend to gain, lose, or share electrons to fill their outer shell.
Example: carbon has 4 valence electrons, so it can form four covalent bonds.
Explain how nonpolar and polar covalent bonds are formed between atoms.
What a covalent bond is
A covalent bond forms when two atoms share a pair of electrons, which allows each atom to fill its outer (valence) shell.
The shared electrons are attracted to the positively charged nuclei of both atoms, which holds the atoms together.
Nonpolar covalent bonds
They form between two atoms with equal or very similar electronegativity.
The shared electrons are distributed evenly, so neither atom carries a partial charge.
Examples: C–C and C–H bonds, and the bonds in O₂ and N₂.
Polar covalent bonds
They form between two atoms with different electronegativities.
The more electronegative atom pulls the shared electrons closer to itself, so the electrons spend more time near that atom.
This creates a partial negative charge (δ−) on the more electronegative atom and a partial positive charge (δ+) on the less electronegative atom. The bond has a dipole.
Examples: O–H, N–H, and C=O bonds.
Continuum of bond types
Bonding falls along a spectrum. As the difference in electronegativity grows, a bond goes from nonpolar covalent to polar covalent to ionic, where one atom takes the electron(s) completely.
Explain how the potential energy of two atoms changes as they form or break a bond.
Forming a bond (atoms approaching each other)
When the atoms are far apart, there is essentially no interaction, and the potential energy is near zero.
As the atoms approach, the attraction between each nucleus and the other atom’s electrons lowers the potential energy.
The potential energy reaches a minimum at a specific distance called the bond length. This is the most stable arrangement.
If the atoms are pushed closer than the bond length, repulsion between the two nuclei and between the electron clouds makes the potential energy rise sharply.
Energy changes
Forming a bond releases energy (often as heat), because the atoms end up at a lower potential energy than when they were separate.
Breaking a bond requires an input of energy, equal to the depth of the energy well (the bond dissociation energy).
Graph interpretation
On a graph of potential energy versus distance between nuclei, the curve dips down to a minimum at the bond length.
The depth of the well shows how much energy is released upon bond formation and how much is needed to break the bond.
Describe each type of chemical bond and intermolecular force: nonpolar covalent bonds, polar covalent bonds, ionic bonds, hydrogen bonds, and van der Waals' forces.
Nonpolar covalent bond
Two atoms share electrons equally, because their electronegativities are the same or very similar.
Example: C–H and C–C bonds in hydrocarbons.
Polar covalent bond
Two atoms share electrons unequally, which produces partial charges (δ+ and δ−).
Example: the O–H bonds in water.
Ionic bond
One atom with much higher electronegativity takes an electron completely from another, forming a cation and an anion that attract each other by electrostatic attraction.
Example: Na⁺ and Cl⁻ in table salt.
Hydrogen bond
This is an attraction between molecules (or parts of molecules), not a bond that shares electrons.
A δ+ hydrogen that is covalently bonded to a very electronegative atom (O or N) is attracted to a lone pair on another electronegative atom (O or N).
Example: hydrogen bonds between water molecules and between the base pairs of DNA.
Van der Waals forces
These are weak, transient attractions between molecules or between nonpolar regions.
They arise because the electron cloud around an atom constantly fluctuates, creating temporary dipoles that induce opposite dipoles in neighboring atoms (London dispersion forces).
Example: the attractions that hold the fatty acid tails of a membrane together.
Summary of categories
Covalent bonds (nonpolar and polar) are intramolecular and share electrons.
Ionic bonds transfer electrons and attract by charge.
Hydrogen bonds and van der Waals forces are weaker intermolecular attractions.
Define and relate the terms electronegativity, polarity, hydrophilic, and hydrophobic.
Electronegativity
It is the tendency of an atom to attract shared electrons toward itself in a covalent bond.
Polarity
A bond is polar when the atoms have different electronegativities, which creates partial charges (δ+ and δ−) at opposite ends of the bond.
A molecule is polar when its polar bonds are arranged so that the partial charges do not cancel out.
A nonpolar bond or molecule has an even distribution of charge.
Hydrophilic (“water-loving”)
Polar or charged molecules and regions are hydrophilic.
They interact favorably with water, mainly by forming hydrogen bonds and ion–dipole interactions, so they dissolve in water.
Hydrophobic (“water-fearing”)
Nonpolar molecules and regions are hydrophobic.
They cannot form hydrogen bonds with water, so they tend to cluster together and are poorly soluble in water.
How the terms relate
Differences in electronegativity between atoms determine bond polarity.
Polar regions are hydrophilic, and nonpolar regions are hydrophobic.
Example: the polar –OH group on an alcohol is hydrophilic, while its long hydrocarbon chain is hydrophobic.
Identify the number of covalent bonds formed by H, C, O, and N.
The general rule
An atom forms the number of covalent bonds needed to complete its outer shell, which is usually 8 valence electrons (the octet rule), or 2 for hydrogen.
Number of bonds by element
Hydrogen (H) forms 1 covalent bond. It has 1 valence electron and needs 1 more to fill its shell.
Carbon (C) forms 4 covalent bonds. It has 4 valence electrons and needs 4 more. This lets it build long chains, rings, and branched structures.
Nitrogen (N) forms 3 covalent bonds. It has 5 valence electrons and needs 3 more.
Oxygen (O) forms 2 covalent bonds. It has 6 valence electrons and needs 2 more.
Counting bonds
A double bond counts as two bonds and a triple bond counts as three.
Example: in a carbonyl group (C=O), oxygen has one double bond, which gives it its total of 2.
Quick memory aid
The numbers of bonds, in order H, O, N, C, are 1, 2, 3, 4.
Compare and contrast the relative electronegativities of H, C, O, and N.
Relative order (highest to lowest)
O (3.5) > N (3.0) > C (2.5) ≥ H (2.1) on the Pauling scale.
Comparisons
Oxygen is the most electronegative of the four, followed by nitrogen. Both pull shared electrons strongly toward themselves.
Carbon and hydrogen have low and similar electronegativities. Carbon is only slightly higher than hydrogen.
What this means for bonds
C–H bonds are considered nonpolar, because the electronegativity difference is very small.
O–H and N–H bonds are polar, with δ− on O or N and δ+ on H.
C–O and C–N bonds are polar, with δ− on O or N and δ+ on C.
The C=O bond is very polar.
Biological significance
Because O and N are strongly electronegative, they are the atoms that make biological molecules polar and able to take part in hydrogen bonding.
Determine whether a region of a molecule is likely to be polar or nonpolar based on its chemical structure and the relative electronegativity of its atoms.
Step-by-step approach
Identify the bonds in the region.
Compare the electronegativities of the bonded atoms (O > N > C ≈ H).
Decide which bonds are polar. A bond is polar if it involves O or N bonded to C or H.
Decide whether the region as a whole is polar, taking the number and arrangement of those bonds into account.
Nonpolar regions
They contain only C–C and C–H bonds, as in hydrocarbon chains and rings.
Example: the long tail of a fatty acid.
Polar regions
They contain bonds such as O–H, N–H, C–O, C–N, and C=O.
Examples of polar functional groups: hydroxyl (–OH), carbonyl (C=O), carboxyl (–COOH), amino (–NH₂), and phosphate.
Charged regions
Groups that carry a full charge, such as COO⁻ and NH₃⁺, are very polar.
Mixed molecules
Many molecules have both. For example, a fatty acid has a nonpolar hydrocarbon tail and a polar carboxyl head.
Relate the potential energy change of an atomic interaction to its "strength".
Core relationship
The strength of an interaction is determined by how much the potential energy drops when the atoms come together.
A larger drop in potential energy (a deeper energy well) means a more stable, stronger interaction.
Why strength relates to energy to break
The energy released when the interaction forms equals the energy needed to break it.
A deeper well means more energy is required to separate the atoms.
A shallow well means the interaction is weak and can be broken with little energy.
Reading an energy diagram
The depth of the well on a potential energy versus distance graph shows the strength of the interaction.
Example: a covalent bond has a very deep well, while a van der Waals interaction has a very shallow one.
Connection to thermal energy
Molecules at body temperature have a small amount of thermal energy. Interactions with wells much deeper than that energy are stable, while very shallow ones constantly form and break.
Compare the relative strengths of bonds and electrostatic interactions and relate these strengths to the amount of energy required to break them.
Relative strengths (strongest to weakest)
Covalent bonds (nonpolar and polar) are the strongest, at roughly 80–110+ kcal/mol (about 350–460 kJ/mol). They require a large amount of energy to break and are stable under normal cellular conditions.
Ionic bonds / ionic interactions are strong in a dry crystal, but in water they are much weaker because water molecules surround the ions and shield their charges. In a biological setting they are often considered comparable to or somewhat stronger than hydrogen bonds.
Hydrogen bonds are weak, at roughly 3–5 kcal/mol (about 12–20 kJ/mol). They break easily but are very important in large numbers.
Van der Waals interactions are the weakest, at roughly 0.1–1 kcal/mol. Each one is very weak, but many add up.
Relating strength to energy
The stronger the interaction, the more energy is needed to break it.
Breaking a covalent bond takes far more energy than the thermal energy available at body temperature (about 0.6 kcal/mol), so covalent bonds do not break spontaneously. Weaker interactions are constantly broken and re-formed by thermal motion.
Biological significance
Covalent bonds hold molecules together (for example, the backbone of DNA and proteins).
Weak interactions are reversible, so they allow dynamic processes. Examples: DNA strands separating, proteins folding, and enzymes binding substrates.
Many weak interactions together can be very strong.
Determine whether a molecule is able to participate in hydrogen bonding with itself or another molecule based on its chemical structure and relative electronegativity of its atoms.
Requirements for a hydrogen bond
A hydrogen bond donor is a hydrogen covalently bonded to a highly electronegative atom (O or N). This hydrogen carries a δ+ charge.
A hydrogen bond acceptor is a highly electronegative atom (O or N) with a lone pair of electrons and a δ− charge.
Step-by-step approach
Look for O–H or N–H bonds, which are potential donors.
Look for O or N atoms with lone pairs, which are potential acceptors.
If the molecule has donors and acceptors, it can hydrogen bond with itself and with other molecules. If it has only acceptors, it can accept hydrogen bonds from partners such as water but not donate.
Important exception
C–H bonds do not form hydrogen bonds, because carbon is not electronegative enough to make the hydrogen sufficiently δ+.
Examples
Water (H₂O) is both a donor and an acceptor, so it hydrogen bonds with itself.
Alcohols (R–OH) and amines (R–NH₂) can act as donors and acceptors.
A ketone (C=O) can act as an acceptor only, since it has no O–H or N–H.
Methane (CH₄) and hydrocarbons cannot form hydrogen bonds.
Determine whether a region of a molecule is likely to be hydrophilic or hydrophobic based on its ability to form hydrogen bonds with water.
Core rule
A region is hydrophilic if it can form hydrogen bonds (or favorable charge interactions) with water. A region is hydrophobic if it cannot.
Step-by-step approach
Examine the region for O–H or N–H groups, which can donate hydrogen bonds to water.
Look for O or N atoms with lone pairs (for example, in C=O or C–O–C), which can accept hydrogen bonds from water.
Check for charged groups (such as COO⁻, NH₃⁺, or phosphate), which interact strongly with water.
If the region has any of these, it is hydrophilic. If it contains only C–C and C–H bonds, it is hydrophobic.
Why hydrophobic regions avoid water
They cannot hydrogen bond with water. Water molecules around them must adopt more ordered arrangements, which is unfavorable, so nonpolar regions cluster together (the hydrophobic effect).
Examples
A hydroxyl group (–OH) is hydrophilic.
A hydrocarbon chain is hydrophobic.
A phospholipid has a hydrophilic head and hydrophobic tails.
Determine whether a given molecule could exist in nature, based on your knowledge of the number of electrons each element can share.
Core rule
A molecule is plausible only if every atom has the right number of covalent bonds to fill its valence shell.
Typical numbers: H = 1, O = 2, N = 3, C = 4.
Step-by-step approach
Count the bonds around each atom, counting a double bond as 2 and a triple bond as 3.
Compare each count with the expected number for that element.
If any atom has too few or too many bonds, the structure is not valid as a neutral molecule.
Examples
CH₄ is valid, because carbon has 4 bonds and each hydrogen has 1.
CH₅ is not valid, because carbon would have 5 bonds.
H₂O is valid, because oxygen has 2 bonds.
H–O–H–H is not valid, because one hydrogen would need 2 bonds.
CO₂ (O=C=O) is valid, because carbon has two double bonds (4 bonds) and each oxygen has one double bond (2 bonds).
NH₃ is valid, because nitrogen has 3 bonds.
Common mistakes to avoid
Forgetting that a double bond uses two of an atom’s bonds.
Leaving a carbon with only 3 bonds (which would leave an unfilled valence shell).
Note on charged species
Atoms can sometimes have an unusual number of bonds if the molecule carries a charge (for example, NH₄⁺ has 4 bonds to nitrogen and H₃O⁺ has 3 bonds to oxygen). For neutral molecules, use the standard numbers.
Differentiate between hydrogen bonds, covalent bonds, ionic bonds, and Van der Waals interactions.
Covalent bonds
How they form: two atoms share electrons.
Strength: very strong.
Where they act: within a molecule (intramolecular).
Example: the O–H bonds inside a water molecule.
Ionic bonds
How they form: one atom transfers electrons to another, producing oppositely charged ions that attract each other.
Strength: strong in a dry crystal, but much weaker in water.
Example: Na⁺ and Cl⁻.
Hydrogen bonds
How they form: a δ+ hydrogen bonded to O or N is attracted to a lone pair on another O or N.
Strength: weak (about 3–5 kcal/mol each).
Where they act: between molecules or between parts of the same large molecule.
Example: hydrogen bonds between water molecules, and between DNA base pairs.
Van der Waals interactions
How they form: temporary, induced dipoles from fluctuating electron clouds attract neighboring atoms.
Strength: the weakest (about 0.1–1 kcal/mol), but they add up when many are present.
Where they act: between any atoms in close contact, especially nonpolar regions.
Example: the attraction between the fatty acid tails in a membrane.
Summary of ordering
By strength: covalent > ionic (in water, comparable to or stronger than hydrogen bonds) > hydrogen bonds > van der Waals.
By mechanism: covalent bonds share electrons, ionic bonds transfer electrons, and hydrogen bonds and van der Waals forces involve partial or temporary charges without sharing electrons.