2.2 Chemical Bonds

Chemical bonds hold atoms together, forming the molecules and compounds that make up everything in your body. In anatomy and physiology, understanding these bonds explains why salt dissolves in your blood, why water regulates your body temperature, and how proteins hold their shape.

Chemical Bonds and Molecules

Molecules, compounds and chemical bonds

Atoms are the fundamental units of all matter, but they rarely exist alone in the body. They combine through chemical bonds to create molecules or compounds with properties distinct from the individual atoms.

• Molecules consist of two or more atoms joined by chemical bonds. They can contain atoms of the same element (like $O_2$, oxygen gas) or different elements (like $H_2O$, water).
• Compounds are a specific type of molecule containing two or more different elements bonded together. Table salt ($NaCl$) is a compound, and its properties are completely different from pure sodium (a reactive metal) or pure chlorine (a toxic gas).
• Chemical bonds are the attractive forces holding atoms together. They form when atoms share or exchange electrons to reach a more stable electronic arrangement. The strength of a bond is measured by its bond energy, which is the amount of energy needed to break that bond.

Ions, cations and anions

Ions are atoms or molecules that carry an electric charge because they've gained or lost electrons. They're central to ionic bonding and play huge roles in physiology (think nerve impulses and muscle contraction).

• Cations are positively charged ions, formed when atoms lose one or more electrons. Losing a negative electron leaves the atom with more protons than electrons, so the net charge is positive. Examples: sodium ion ($Na^+$), calcium ion ($Ca^{2+}$), iron(III) ion ($Fe^{3+}$).
• Anions are negatively charged ions, formed when atoms gain one or more electrons. The extra electron(s) give the atom a net negative charge. Examples: chloride ion ($Cl^-$), oxide ion ($O^{2-}$), phosphate ion ($PO_4^{3-}$).

A quick way to remember: cations are positive (the "t" looks like a plus sign), and anions are negative.

Types of Chemical Bonds

Ionic vs covalent bonds

These are the two main types of chemical bonds, and they differ in what happens to the electrons.

Ionic bonds form between a metal and a nonmetal through electron transfer:

1. The metal atom donates one or more electrons to the nonmetal atom.
2. The metal becomes a cation (positive), and the nonmetal becomes an anion (negative).
3. The opposite charges attract, holding the ions together.

Ionic compounds like sodium chloride ($NaCl$) and potassium bromide ($KBr$) tend to have high melting and boiling points, are brittle in solid form, and conduct electricity when dissolved in water (because the ions separate and move freely).

Covalent bonds form between two nonmetals through electron sharing:

1. Two atoms share one or more pairs of electrons.
2. The shared electrons orbit both nuclei, holding the atoms together as a molecule.

Covalent compounds like water ($H_2O$) and methane ($CH_4$) generally have lower melting and boiling points, are often soft or flexible as solids, and typically do not conduct electricity. The distance between the nuclei of two bonded atoms is called the bond length.

Spectrum of covalent bonds

Not all covalent bonds share electrons equally. The key factor is electronegativity, which is how strongly an atom attracts shared electrons.

• Nonpolar covalent bonds form when electrons are shared equally. This happens between identical atoms or atoms with very similar electronegativities. Examples: $H_2$, $N_2$, and $O_2$.
• Polar covalent bonds form when electrons are shared unequally because one atom is more electronegative than the other. The more electronegative atom pulls the shared electrons closer, gaining a partial negative charge ($\delta-$), while the other atom gets a partial positive charge ($\delta+$). Examples: water ($H_2O$), hydrogen fluoride ($HF$), ammonia ($NH_3$).

This polarity matters a lot in the body:

• Polar molecules (like sugars and amino acids) are attracted to other polar molecules and dissolve well in water, a polar solvent. This is why so many biological reactions happen in water.
• Nonpolar molecules (like fats and oils) are attracted to other nonpolar molecules and don't dissolve in water. This is why lipids can form cell membranes that separate watery compartments.

Hydrogen bonds in water

Hydrogen bonds are relatively weak attractions that form between a hydrogen atom bonded to a highly electronegative atom (nitrogen, oxygen, or fluorine) and another electronegative atom nearby. They're not true chemical bonds like ionic or covalent bonds; they're intermolecular forces. But their collective effect is enormous.

In water, the partially positive hydrogen ($\delta+$) of one $H_2O$ molecule is attracted to the partially negative oxygen ($\delta-$) of a neighboring $H_2O$ molecule. This constant network of hydrogen bonds gives water properties that are essential for life:

1. High specific heat capacity means water absorbs a lot of heat before its temperature rises significantly. This helps your body maintain a stable internal temperature.
2. Cohesion is the tendency of water molecules to stick to each other. This helps with the transport of blood and other fluids through narrow vessels.
3. Adhesion is the attraction of water molecules to other polar surfaces, allowing water to "climb" up narrow tubes (like capillaries in your circulatory system).
4. High surface tension results from hydrogen bonds pulling surface water molecules inward, creating a strong surface film.

Beyond water, hydrogen bonds are critical for holding together the structure of proteins, DNA, and RNA. The double helix of DNA, for example, is held together by hydrogen bonds between complementary base pairs.

Electronic Structure and Bonding

Valence electrons are the electrons in an atom's outermost energy level, and they're the ones involved in bonding. The number of valence electrons determines how an atom will bond with others.

The octet rule states that atoms tend to gain, lose, or share electrons until they have eight valence electrons in their outer shell (or two, in the case of hydrogen and helium). This full outer shell is the most stable configuration, which is why atoms form bonds in the first place.

Lewis structures (also called Lewis dot diagrams) are a way to visualize bonding. They show the valence electrons around each atom, including shared pairs (bonds) and unshared lone pairs. Some molecules can be drawn with more than one valid Lewis structure, a phenomenon called resonance, where the actual electron distribution is an average of all possible arrangements.