Chapter 1-8 Chemistry Practice: Covalent Bonds, Polarity, and Ions
Covalent Bonding, Octets, and Diatomic Molecules
Key idea: covalent bonding = sharing electrons to satisfy the octet rule (eight electrons in the outermost shell).
Octet goal: have a total of 8 electrons in the outer shell (octet). When achieved, atoms resemble noble gases and are more stable.
Visual cue: dots around atoms in diagrams often represent valence electrons; the octet rule is the guiding idea for bonding.
Diatomic molecules: some atoms pair up and exist as molecules of two atoms without being bonded to another element (auto-stable basics). Notable example discussed: hydrogen forms H₂ (two hydrogen atoms bound together).
Clarification: molecular hydrogen (H₂) consists of two H atoms bonded; in elemental form on the periodic table you see a single H atom, but in a molecule it exists as H₂.
Implication: atom pairs can satisfy the octet by sharing electrons to form a bond with another atom of the same element.
Bond Types and Electron Sharing
Sharing can be unequal or equal:
If sharing is not 50/50, the bond is polar covalent (partial charges develop).
If sharing is roughly equal, the bond is nonpolar covalent.
Polarity arises from unequal sharing of electrons, creating partial charges (δ− on the more electronegative atom and δ+ on the other).
Example of polarity concept: polar covalent bonds result in electrostatic attractions between molecules (dipole–dipole interactions) and can influence solubility and boiling points.
How Many Bonds by Group (Valence) on the Periodic Table
The group number indicates the number of valence electrons.
To reach the octet (eight valence electrons), atoms in different groups require different numbers of electrons from bonding:
Group 5 (e.g., nitrogen, N): needs 3 more electrons to reach eight; can form up to 3 bonds (three singles, or one single + one double, or a triple bond).
Group 6 (e.g., oxygen, O): needs 2 more electrons; typically forms 2 bonds.
Group 7 (e.g., fluorine, F; chlorine, Cl): needs 1 more electron; forms 1 bond.
Examples mentioned:
Nitrogen (Group 5): can achieve an octet via a triple bond (e.g., N≡N in N₂) or other combinations (e.g., NH₃ has three single bonds around N).
Oxygen (Group 6): often forms two bonds (e.g., H₂O with two O–H bonds).
Hydrogen (Group 1): forms one bond; its shell can hold 2 electrons max.
Important caveats:
There are exceptions when moving further left on the periodic table (toward the nucleus) where rules like “eight in the outer shell” don’t apply in the same way (e.g., H, He, Li, Be, B behavior differs). The main practical rule for bonding in most chemistry contexts taught here is how many bonds are needed to reach eight, as dictated by group valence.
Left side tends to lose electrons to form cations; right side tends to gain electrons to form anions; this leads to ionic compounds when left and right combine (e.g., NaCl).
Ionic vs Covalent Compounds and Noble-Gas Electron Configurations
Ionic compounds form from electrostatic attraction between oppositely charged ions (cations and anions). Example discussed: sodium chloride (NaCl).
Cations are positively charged ions (loss of electrons); anions are negatively charged ions (gain of electrons).
Achieving an octet in ions:
Ions tend to end up with the outer shell that has 8 electrons, i.e., a noble gas electron configuration.
Example: sodium loses one electron to form Na⁺; its outer shell becomes the same electron configuration as neon: 1s^2\;2s^2\;2p^6 (i.e., 8 electrons in the outer shell of the resulting species).
Important conceptual point: protons do not change when forming ions via electron loss/gain; changing the number of electrons changes the element’s charge, but not its identity (the element’s symbol remains the same for the ion formed from that element).
Hydration and dissociation in solution:
In aqueous solutions, salts dissociate into their constituent ions (e.g., NaCl → Na⁺ + Cl⁻).
These ions retain their noble-gas-like electron configurations after gaining/losing electrons.
Common context: many substances we encounter (e.g., table salt) include ions as part of their composition; ionic compounds are typically salts.
Electronegativity, Polarity, and How to Judge Polarity
Electronegativitiy trend: increases as you move to the right across a period and up a group on the periodic table.
Higher electronegativity means an atom is more eager to attract electrons when bonding.
A practical polarity rule of thumb from the lecture:
If the electronegativity difference between two bonded atoms is Δχ ≥ 0.5, the bond is polar covalent.
If Δχ < 0.5, the bond is nonpolar covalent.
Common electronegativity examples:
Fluorine (F) is the most electronegative atom (far right and towards the top).
Oxygen (O) and nitrogen (N) are highly electronegative but less than fluorine.
Visualizing polarity using a right-to-left and up-to-down map:
Atoms further to the right and higher on the periodic table are more electronegative; thus, when bonded to carbon or another atom, the bond tends to be more polar if the other atom is less electronegative.
Practical implications of polarity:
Polar molecules tend to dissolve in polar solvents (e.g., water).
Nonpolar molecules tend to be immiscible with water and dissolve in nonpolar solvents; examples: methane (CH₄) and propane (C₃H₈) are nonpolar and gases at room temperature.
The OH group in alcohols makes them polar; water (H₂O) is highly polar.
The concept of polarity also leads to inter-molecular attractions beyond covalent bonds: partial charges on different molecules attract each other (electrostatic interactions).
A quick comparison exercise mentioned:
Between bonds like C=O (carbon–oxygen) and C–N (carbon–nitrogen), the bond with the atom further to the right/up (oxygen, in this case) is more polar because oxygen is more electronegative than nitrogen.
Carbon–carbon or carbon–hydrogen bonds are typically nonpolar (little to no electronegativity difference).
Solubility, Boiling Points, and the Role of Polarity in Life
Polar molecules (e.g., water, many alcohols) tend to have higher boiling points than nonpolar molecules of similar molar mass due to stronger intermolecular attractions.
Polar solvents like water dissolve polar solutes more readily; polarity helps determine solubility and dissolution rates.
Examples discussed:
Water (H₂O) is highly polar and serves as a reference point for temperature changes (boil at 100°C, freeze at 0°C).
Alcohols are polar; methane (CH₄) and propane (C₃H₈) are nonpolar and gaseous at relatively low temperatures due to small size and nonpolarity.
Hydrophobic effect (driving force in biology):
Water tends to reorganize around nonpolar molecules, reducing contact with water (minimizes surface area) and causing nonpolar regions (like hydrocarbon chains) to coalesce.
This effect underlies the initial folding and organization of proteins (hydrophobic amino acids clustering away from water) and guides the formation of secondary and tertiary structures without external energy input.
Terminology:
Hydrophobic effect: the driving force by which nonpolar substances aggregate in aqueous solutions to minimize disruption of water's hydrogen-bonding network.
Polar vs nonpolar distribution of electrons can be visualized as a heat-map-like distribution of electron density across bonds.
Practical Examples and Clarifications from the Lecture
Hydrogens and diatomic molecules:
H₂ is diatomic and forms a covalent bond between two hydrogen atoms.
When considering elemental hydrogen on the periodic table, you refer to the single H atom; in molecules, hydrogen pairs up to form H₂.
Bonding counts and what they imply:
Left-hand side elements (toward the alkali/alkaline earth metals) tend to lose electrons to form cations and participate in ionic bonding with right-hand side elements.
Right-hand side elements (nonmetals) tend to gain electrons to complete octets and form covalent bonds with other nonmetals.
Important caveats about periodic trends:
The simple octet rule is a practical guide; there are exceptions, especially for elements toward the left of the table or with empty d-orbitals in some cases (not deeply discussed here).
Common abbreviations and notations:
δ+ and δ− denote partial positive and partial negative charges in polar covalent bonds.
Na⁺ and Cl⁻ denote ions formed from sodium and chloride respectively; NaCl is an ionic compound (salt).
A cation is a positively charged ion; an anion is a negatively charged ion.
Quick decision-making tips for exams (as discussed in the lecture):
To predict polarity, compare the bonded atoms in terms of their positions on the periodic table, focusing on the direction toward the right/up.
If two bonds are compared and one involves an atom that is farther to the right/up, that bond is typically the more polar one.
Recap of key terms to memorize:
Octet, diatomic, covalent bond, polar covalent bond, nonpolar covalent bond, electronegativity, cation, anion, salt, hydrophobic effect.
Quick Concept Summary for Exam Prep
The octet rule drives most covalent bonding: achieve eight valence electrons in the outer shell to reach noble-gas stability.
Diatomic molecules (H₂, O₂, N₂, etc.) exist as two-atom units; H₂ is the classic example of a diatomic molecule.
Bonding by group follows valence electrons: Group 5 needs 3 more electrons (can form up to 3 bonds); Group 6 needs 2 more (2 bonds); Group 7 needs 1 (1 bond).
Ionic bonding arises from electron transfer on left-side elements, forming cations and anions; covalent bonding arises from sharing on the right side; salts like NaCl illustrate ionic compounds.
Electronegativity increases across a period and up a group; fluorine is the most electronegative element.
Polarity depends on electronegativity difference; Δχ ≥ 0.5 generally indicates a polar covalent bond; Δχ < 0.5 indicates nonpolar covalent.
Polar bonds lead to higher solubility in polar solvents (e.g., water) and influence boiling/melting behavior; nonpolar compounds tend to be less soluble in water.
The hydrophobic effect explains how water drives nonpolar substances together, critical in protein folding and membrane biology.
Ion formation preserves the number of protons; ions achieve noble-gas electron configurations by gaining or losing electrons; transactions in solution (e.g., NaCl → Na⁺ + Cl⁻) result in dissociation.
Real-world examples like water, alcohols, methane, and propane illustrate the practical outcomes of polarity on phase, solubility, and interactions in biological and chemical systems.