Matter, Atoms, Bonds, and pH – Quick Reference

Matter is made of chemical elements.
Element: a pure type of atom; cannot be broken down into other elements with the same properties.
Nature contains 92 naturally occurring elements.

Each element has a symbol (e.g., H, C, O, N) and properties.
Human body composition (major elements):
- Oxygen: $65\%$
- Carbon: $18.5\%$
- Hydrogen: $9.5\%$
- Nitrogen: $3\%$
These four major elements make up about $96\%$ of body mass.
Trace elements are needed in smaller amounts (for survival) and include:
- Calcium (Ca) ~ $1.5\%$ , Phosphorus (P) ~ $1\%$ , Potassium (K) ~ $0.4\%$ , Sulfur (S) ~ $0.3\%$ , Sodium (Na) ~ $0.2\%$ , Chlorine (Cl) ~ $0.2\%$ , Magnesium (Mg) ~ $0.1\%$
Deficiency in trace elements can cause disease (e.g., iron for oxygen transport; iodine for thyroid hormones).

Element: pure form of an atom; cannot be broken down without changing properties.
Atom: smallest unit of an element that retains its properties.
Subatomic particles: protons (+), neutrons (neutral), electrons (−).
Atom structure: nucleus (protons and neutrons) + surrounding electron shells.

Proton: + charge; located in nucleus; number = atomic number $Z\,.$
Neutron: neutral; located in nucleus; contributes to mass.
Electron: − charge; located in electron shells; outer ones (valence electrons) participate in bonding.
In a neutral atom, Z = \text{#protons} = \text{#electrons}\,.
Mass number $A = Z + N\,$ where $N = A - Z$ is the number of neutrons.
Atomic mass is measured in Daltons (Da).

Atomic number $Z$ = number of protons (and equals number of electrons in neutral atom).
Mass number $A$ = total number of protons and neutrons.
Isotopes: same element (same $Z$ ) with different numbers of neutrons (e.g., $^{12}C, ^{13}C, ^{14}C$ ).
Isotopes can be radioactive and used as tracers in medicine (e.g., carbon-13/14 in imaging like PET scans).
Carbon standard forms: Carbon-12 (most common in organisms), Carbon-13, Carbon-14.

Electron shells around the nucleus; electrons occupy shells with limited capacity:
- 1st shell: up to $2$ electrons
- 2nd shell: up to $8$ electrons
- 3rd shell: up to $8$ electrons (and so on)
Stability: outer shell filled (typically 8 electrons) → less reactive.
Valence electrons: electrons in the outer shell involved in bonding.
Atoms bond to fill their outer shell.

Covalent bonds: sharing electrons between atoms; typically strong and not easily broken by water.
Ionic bonds: transfer of electrons creating ions; bond is relatively weak in water and easily dissociates.
Hydrogen bonds: weak attractions between molecules (not true bonds); important for 3D structure of DNA, proteins, and water network; can be broken by heat (denaturation).

Methane (CH$_4$): carbon shares 4 electrons with 4 hydrogens; covalent bonds, nonpolar due to similar electronegativities.
Water (H$_2$O): oxygen and hydrogen share electrons; polar covalent bond due to significant electronegativity difference.
Bond type depends on electronegativity difference; small difference -> nonpolar covalent; large difference -> polar covalent.
Electron distribution:
- C (electronegativity ≈ $2.55$ ) vs H (≈ $2.20$ ): similar -> nonpolar covalent.
- O (≈ $3.44$ ) vs H (≈ $2.20$ ): different -> polar covalent; electrons spend more time near O (δ⁻ on O, δ⁺ on H).

Electronegativity values guide bond type:
- Similar values → nonpolar covalent (equal sharing).
- Different values → polar covalent (unequal sharing).
Example: H–O in water is polar covalent due to electronegativity difference.

Example: Sodium chloride (NaCl).
Sodium tends to lose 1 electron to form Na⁺; chlorine tends to gain 1 electron to form Cl⁻.
Ionic bond forms NaCl; in water, the bond can break and ions separate.

Water can dissociate into H⁺ and OH⁻; pH measures hydrogen ion concentration.
pH scale: 0 to 14; 7 is neutral; below 7 is acidic; above 7 is basic.
Formula: $\text{pH} = -\log_{10}[\mathrm{H^+}]$
Each pH unit represents a 10-fold change in hydrogen ion concentration:
- e.g., pH 6 is ten times more acidic than pH 7; pH 5 is 100x more acidic than pH 7.
Strong acids/bases and safety:
- Stomach acid ~ pH ≈ 2.
- Strong bases (e.g., oven cleaners) can be highly caustic.

pH shifts can profoundly affect enzyme activity and protein stability through changes in hydrogen ion concentration and hydrogen-bond networks.

Reactants: substances at the left of the arrow; products: substances at the right.
Reversible (bidirectional) reactions can proceed in both directions (arrow both ways).
Example context: water can be formed or broken apart depending on conditions.

Iodine and thyroid health:
- Iodine is needed to synthesize thyroid hormones (T3/T4).
- Iodine deficiency can cause goiter.
- Iodized salt is a common dietary source of iodine.
Fluoride:
- Added to water and toothpaste to prevent cavities; high exposure is a health concern.
Fortified foods and nutrition labels:
- Foods may be fortified with minerals (e.g., calcium).
- Read nutrition labels to know ingredients and added nutrients.

Radioisotopes (e.g., carbon-13, carbon-14) are used as tracers in diagnostics (e.g., PET scans).
Excessive exposure to radioactivity can mutate DNA and be harmful; limit exposure and follow safety guidelines.

Define matter and major vs trace elements.
State the major elements of the human body and approximate percentages.
Write the mass number and neutron equations: $A = Z + N\,,\ N = A - Z\,.$
Describe electron shells capacities and the concept of valence electrons.
Distinguish covalent, polar covalent, and nonpolar covalent bonds with examples.
Explain ionic bond formation with Na and Cl and why it breaks in water.
Define hydrogen bonds and give two roles in biology.
Explain pH concept, the formula, and the 10-fold change per unit.
What is the significance of isotopes in medical imaging?