Notes on Matter, Atomic Structure, Bonds, Water, and pH

2.1 The Organization of Matter

An element is a pure substance that cannot be broken down into simpler substances by ordinary chemical or physical techniques.
All matter (anything that occupies space and has mass) is composed of elements and combinations of elements.
Ninety-two different elements occur naturally on Earth.
Other artificial elements have been synthesized in the laboratory.
Key elements in living organisms:
- Four elements (carbon, hydrogen, oxygen, and nitrogen) make up more than 96% of the weight of living organisms.
- Seven elements (calcium, phosphorus, potassium, sulfur, sodium, chlorine, and magnesium) make up nearly 4%.
- Several trace elements vital for normal biological functions (such as iodine) occur in quantities less than 0.01%.

2.2 Atoms and Molecules

Elements are composed of atoms – the smallest units that retain the chemical and physical properties of an element.
Atoms combine chemically in fixed numbers and ratios to form the molecules of living and nonliving matter.
Molecular names are written as a chemical formula, using standard symbols for elements and subscripts for the number of atoms of each element (e.g., $CO_2$ ).

2.3 Compounds

Molecules whose component atoms are different (such as carbon dioxide) are called compounds.
The chemical and physical properties of compounds are typically distinct from those of their atoms or elements.
Example: Hydrogen (H) and Oxygen (O) – the elements that form liquid water (H extsubscript{2}O) – are both highly flammable gases.

2.2 Atomic Structure

Each element consists of one type of atom.
Each atom consists of an atomic nucleus surrounded by fast-moving, negatively-charged electrons.
Atomic nuclei contain positively charged protons and uncharged neutrons.
The number of protons (atomic number, Z) identifies an element.

Protons and Neutrons

Atoms are identified by their atomic number; the number of protons in the nucleus does not vary.
All atoms except hydrogen also contain at least one neutron.
A neutron and proton have almost the same mass, about $1.66 \times 10^{-24}$ g (one dalton, Da).
Electrons (mass ≈ $1/1800$ Da) contribute no significant mass.
An atom’s mass number is the total number of protons and neutrons in the nucleus.

2.3 Isotopes

Isotopes are distinct forms of atoms of an element with the same number of protons but different numbers of neutrons.
Organisms can use any hydrogen or carbon isotope, for example, without a change in their chemical reactions.

2.3 Mass and Weight

Mass is the amount of matter in an object.
Weight is a measure of the pull of gravity on an object.
In outer space, the mass of an object remains the same; however, it has no weight.

2.4 Electron Orbitals

In an atom, the number of electrons equals the number of protons in the nucleus, making the atom electrically neutral.
The speed of electrons in motion around the nucleus approaches the speed of light.
The region of space where an electron is found most frequently is its orbital.

Electron Orbitals (cont’d.)

Orbitals take different shapes depending on distance from the nucleus and degree of repulsion by electrons in other orbitals.
The most stable and balanced condition occurs when an orbital contains a pair of electrons.
Under certain conditions, electrons may pass from one orbital to another within an atom, or pass completely to another atom.
The ability of electrons to move between orbitals is the basis of chemical reactions that combine atoms into molecules.

Electron Shells

Within an atom, electrons are found in regions of space called energy levels or shells.
The shell nearest the nucleus may be occupied by a maximum of two electrons in a single, spherical orbital (1s).
Electrons at the second energy level may occupy one spherical orbital (2s) and three dumbbell-shaped 2p orbitals.

Electron Orbitals (illustrations)

1s, 2s, 2p orbitals illustrate how electrons occupy shells.

Elements and Electron Shells

The outermost energy level typically contains one to eight electrons occupying a maximum of four orbitals.
Number of electrons in energy levels follows a pattern tied to atomic number (protons).
For example, Hydrogen (H) has 1 proton and 1 electron (outer level with 1 electron); Helium (He) has 2 electrons in the first shell; Lithium (Li) begins filling the second shell, and so on.
The diagrammatic representation shows how energy levels fill (Energy level 1, Energy level 2, Energy level 3).

2.3 (continued) Electrons Determine Chemical Activity

The electrons in an atom’s outermost energy level are its valence electrons.
Atoms with a completely filled outermost energy level are nonreactive, or inert (e.g., Helium, Neon, Argon).
Atoms in which the outermost energy level is not completely filled with electrons are chemically reactive (e.g., Hydrogen).
The tendency of an atom to gain, lose, or share valence electrons underlies chemical bonds and forces that hold molecules together.
Atoms that differ from the stable configuration by more than one or two electrons tend to share electrons in joint orbitals with other atoms (e.g., Oxygen).

2.4 Chemical Bonds and Chemical Reactions

Atoms of reactive elements tend to combine into molecules by forming chemical bonds.
The four most important chemical linkages in biological molecules are ionic bonds, covalent bonds, hydrogen bonds, and van der Waals forces.
Chemical reactions occur when atoms or molecules interact to form new chemical bonds or break old ones.

Ionic Bonds

Ionic bonds result from electrical attractions between atoms that gain or lose valence electrons completely (ions).
A positively charged ion (lost an electron) is called a cation (e.g., Na extsuperscript{+}).
A negatively charged ion (gained an electron) is called an anion (e.g., Cl extsuperscript{−}).
When a hydrogen atom (H) loses its single electron to form a hydrogen ion (H extsuperscript{+}), it consists only of a proton (often called a proton).
Metallic atoms (e.g., Ca extsuperscript{2+}, Mg extsuperscript{2+}, Fe extsuperscript{2+/3+}) lose electrons to form cations that bond with anions.
Ionic bonds have three key features:
- They exert an attractive force over greater distances than any other chemical bond.
- Their attractive force extends in all directions.
- They vary in strength depending on the presence of other charged substances.

Ionic Bonds (cont’d.)

Example: Formation of NaCl from Na and Cl (Na atom loses an electron to form Na extsuperscript{+}; Cl atom gains an electron to form Cl extsuperscript{−}).

Covalent Bonds

Covalent bonds form when atoms share a pair of valence electrons rather than gaining or losing them.
In molecular diagrams, a covalent bond is designated by a pair of dots or a single line representing a pair of shared electrons.
Example: $ext{H}_2 ext{ is represented as } ext{H:H} ext{ or } ext{H-H}.$

Covalent Bonds (cont’d.)

Shared orbitals that form covalent bonds extend between atoms at discrete angles and directions, giving covalently bound molecules distinct three-dimensional forms.
Example: Carbon, with four unpaired outer electrons, forms tetrahedrons, branched and unbranched chains, and rings—with single, double, or triple bonds.
Methane (CH extsubscript{4}) is a tetrahedron with four covalent bonds fixed at an angle of $109.5^{\circ}$ .

Shared-Orbital Model of Methane

Visualization showing tetrahedral geometry causing methane’s shape.

Polarity

Electronegativity is the measure of an atom’s attraction for the electrons it shares in a chemical bond with another atom.
The more electronegative an atom is, the more strongly it attracts shared electrons.
Covalent bonds differ widely in degree of sharing of valence electrons, depending on the difference in electronegativity between the bonded atoms.
Nonpolar covalent bond: electrons are shared equally.
Polar covalent bond: electrons are shared unequally; the more electronegative atom carries a partial negative charge, δ−, while the other carries a partial positive charge, δ+.
The overall molecule is polar if there is an unequal distribution of charge.

Water, a Polar Molecule

In water, an oxygen atom forms polar covalent bonds with two hydrogen atoms.
Electrons are attracted more strongly to the oxygen nucleus than to the hydrogen nuclei.
The water molecule is asymmetric; the oxygen atom is δ− and the hydrogens are δ+.
This polar nature enables water to adhere to ions and weaken their attractions.

Polarity in Water (cont’d.)

Oxygen, nitrogen, and sulfur share electrons unequally with hydrogen; —OH, —NH, and —SH groups tend to be located asymmetrically in biological molecules, making these regions polar.
Carbon–hydrogen bonds tend to be arranged symmetrically (as in methane), so their partial charges cancel and the molecule is nonpolar.

Polar Molecules Associate and Exclude Nonpolar

Polar molecules attract other polar molecules and charged ions, forming polar associations that tend to exclude nonpolar molecules.
Polar molecules that associate readily with water are hydrophilic ("water-loving").
Excluded nonpolar molecules tend to clump together in nonpolar associations, reducing surface area exposed to the polar environment.
Nonpolar substances are hydrophobic ("water-fearing").

Hydrogen Bonds

Hydrogen bonds are attractions between partially positive hydrogen atoms (sharing electrons unequally with oxygen, nitrogen, or sulfur) and partially negative atoms in a different covalent bond.
Hydrogen bonds may be intramolecular (within the same molecule) or intermolecular (between molecules).
Individual hydrogen bonds are weak, but numerous hydrogen bonds can be collectively strong.
They stabilize the three-dimensional structure of large biological molecules such as proteins.
Hydrogen bonds between water molecules are responsible for many properties of water essential to life.
Hydrogen bonds begin to break extensively at temperatures above $45^{\circ}\text{C}.$

Van der Waals Forces

Van der Waals forces are weak forces that develop over very short distances between nonpolar molecules (or regions of molecules) as moving electrons accumulate by chance in one part of a molecule or another.
Temporary zones of positive and negative charge make the molecule polar briefly, attracting or repelling charged regions of other molecules.
Many such bonds can stabilize the shape of large molecules, such as proteins.

Molecular Geometry Determines Function

Molecular geometry is the three-dimensional arrangement of atoms in a molecule.
A molecule’s characteristic size and shape are determined by its atoms and their orbitals.
Molecular shape is crucial because it determines the function of a molecule.
A cell’s function depends on interactions between molecules, similar to a lock and key.
Drugs may mimic the shape of natural molecules to block or interfere with interactions of these natural molecules.

Chemical Equations

An arrow shows the direction of the chemical reaction; reactants on the left, products on the right.
Chemical reactions written in balanced form (the same number of each type of atom on both sides) are chemical equations.
Example: Overall reaction of photosynthesis
$6 \ CO2 + 6 H2O \rightarrow C6H{12}O6 + 6 O2$

2.4 Hydrogen Bonds and the Properties of Water

Hydrogen bonds between water molecules produce a water lattice that affects properties such as density, heat absorption, cohesion, and surface tension.
Polarity of water molecules contributes to the formation of distinct polar and nonpolar environments critical to cell organization.
Water is a solvent for charged or polar molecules.
Water molecules can dissociate into hydrogen and hydroxide ions.

A Lattice of Hydrogen Bonds

Liquid water forms a water lattice; each water molecule constantly breaks and reforms hydrogen bonds with its neighbors (average ~3.4 bonds).
Ice forms a rigid, crystalline lattice; each water molecule forms four hydrogen bonds, spacing molecules farther apart than in the liquid.
Ice is about 10% less dense than liquid water, an unusual property that makes ice float.
Water reaches its greatest density at a temperature of $4^{\circ}C$ .

Water and Temperature

The hydrogen-bond lattice of liquid water retards the escape of individual water molecules as water is heated.
Water remains liquid over a wide temperature range (0°C to 100°C).
A large amount of heat must be added to break enough hydrogen bonds to boil water.
Water has a relatively high specific heat; it can absorb or release relatively large quantities of heat energy without undergoing extreme changes in temperature.

Specific Heat and Calories

Specific heat: amount of heat energy required to increase the temperature of a given quantity of water.
Measured in calories.
Calorie (small c): heat energy required to raise 1 g of water by 1°C.
Calorie (capital C): a kilocalorie (kcal) or 1,000 calories.

Heat of Vaporization

A large amount of heat ( $586\ \text{cal/g}$ ) must be added to give water molecules enough energy to break loose from liquid water and form a gas.
This heat of vaporization enables cooling in many organisms.

Chemical Reactions in Cells Involve Aqueous Solutions

Concentration is the number of molecules or ions of a substance in a unit volume (e.g., mL or L).
The number of molecules can be calculated indirectly by using atomic/molecular weights and Avogadro’s number.
Example (carbon): mass number 12; mass per carbon atom = $12 \times (1.66 \times 10^{-24} \text{ g}) = 1.992 \times 10^{-23} \text{ g}.$

Concentration (cont’d)

Divide the total weight of a sample by the weight of a single atom to obtain the number of atoms in the sample.
The atomic weight of an element (mass number in grams) equals the molecular weight of a molecule for that element.
The molecular weight of a molecule is the sum of the atomic weights of all atoms present in the molecule.

Avogadro’s Number and the Mole

Dividing atomic or molecular weight by the weight of one atom or molecule yields a constant number, $N_A = 6.022 \times 10^{23}$ , atoms or molecules per mole.
Example: dividing 12 g by the weight of one carbon atom yields the number of atoms:
$\dfrac{12\ \text{g}}{1.992 \times 10^{-23} \text{ g}} = 6.022 \times 10^{23}.$
When concentrations are described, the atomic weight of an element or the molecular weight of a compound is known as a mole (mol).
The amount of a substance that contains $6.022 \times 10^{23}$ atoms or molecules is a mole.
The number of moles of a substance dissolved in 1 L of solution is the molarity, $M = \dfrac{\text{moles of solute}}{\text{liters of solution}}.$
Two solutions with the same volume and molarity contain the same number of molecules.

2.5 Water Ionization and Acids, Bases, and Buffers

Water dissociates into positively charged hydrogen ions (H extsuperscript{+} or protons) and negative hydroxide ions (OH extsuperscript{−});
$\text{H}_2\text{O} \rightleftharpoons \text{H}^+ + \text{OH}^-.$
The reaction is reversible; in pure water, [H extsuperscript{+}] = [OH extsuperscript{−}].

Acids and Bases

Acids are proton donors that release H extsuperscript{+} (and anions) when dissolved in water; e.g., hydrochloric acid (HCl) dissociates as
$\text{HCl} \rightarrow \text{H}^+ + \text{Cl}^-.$ (in water: HCl ⇌ H extsuperscript{+} + Cl extsuperscript{−})
Bases are proton acceptors that reduce the H extsuperscript{+} concentration; most release an OH extsuperscript{−} and a cation.
Example: NaOH → Na extsuperscript{+} + OH extsuperscript{−}; OH extsuperscript{−} + H extsuperscript{+} → H ext{O}.
Some bases do not dissociate to produce hydroxide directly; e.g., NH extsubscript{3} + H extsubscript{2}O → NH extsubscript{4}^+ + OH extsuperscript{−}.

pH

The concentration of H extsuperscript{+} ions relative to OH extsuperscript{−} determines acidity; measured on the pH scale (0 to 14):
$\text{pH} = -\log_{10}[\text{H}^+].$
Pure water has pH 7 (neutral); acids have pH < 7; bases have pH > 7.
pH 0 corresponds to 1 M HCl; pH 14 corresponds to 1 M NaOH.
Each whole pH number represents a 10-fold change in [H extsuperscript{+}].
All living organisms maintain internal acidity near pH 7 (homeostasis).

pH Scale Illustrative Values

Examples on the scale include: HCl (pH ~0); lemon juice (pH ~2); pure water (pH 7); sea water (pH ~8); NaOH (pH ~14).

Buffers

Buffers absorb or release H extsuperscript{+} to resist pH changes.
Most buffers are weak acids or bases that dissociate reversibly in water to release or absorb H extsuperscript{+} or OH extsuperscript{−}.

Carbonic Acid – Bicarbonate Buffer System

A carbonic acid–bicarbonate buffer system buffers blood pH:
$\text{H}2\text{CO}3 \rightleftharpoons \text{HCO}_3^- + \text{H}^+.$
Each buffer has a specific range of greatest buffering capacity.
Normal blood pH (7.4) lies outside the region of greatest buffering capacity for this system.
Graphical representation shows the % buffer present as HCO extsubscript{3}^- vs H extsubscript{2}CO extsubscript{3} across pH values.

2.5 Water Ionization and Buffers (cont’d) – Lattice and Buffering

The carbonic acid–bicarbonate buffer maintains blood pH by shifting between H extsubscript{2}CO extsubscript{3} and HCO extsubscript{3}^- as needed.
Region of greatest buffering capacity depends on the ratio of carbonic acid to bicarbonate ion.

Summary of Key Concepts

Matter is composed of elements, which combine to form atoms and molecules.
Atoms have a nucleus of protons and neutrons and orbiting electrons; the atomic number identifies the element.
Isotopes differ in neutron number but share proton number; mass and weight differ accordingly.
Electron arrangement (shells and orbitals) determines chemical activity and bonding.
Chemical bonds include ionic, covalent, hydrogen bonds, and van der Waals forces; bond type influences molecule structure and function.
Water is a polar molecule that drives many biological processes; hydrogen bonding governs water’s properties, including density, heat capacity, and solvent behavior.
Acids, bases, and buffers regulate pH in biological systems; buffers minimize pH changes.
The mole and Avogadro’s number connect atomic/molecular scale to measurable quantities; molarity relates moles to solution volume.
Molecular geometry and polarity influence biological interactions, drug action, and enzyme specificity.