Biology I - Chapter 2

Matter, Elements, and Compounds

• Matter is anything that takes up space and has mass.

• An element is a substance that cannot be broken down into other substances by chemical reactions ($92$ naturally occurring elements).

• A compound consists of two or more elements in a fixed ratio, demonstrating emergent properties distinct from its constituent elements.

The Elements of Life

• Approximately $20\text{--}25\%$ of natural elements are essential for life (humans require $25$, plants require $17$).

• Four elements make up $96\%$ (or $96.3\%$) of living matter: carbon ($\text{C}$), hydrogen ($\text{H}$), oxygen ($\text{O}$), and nitrogen ($\text{N}$).

• Calcium ($\text{Ca}$), phosphorus ($\text{P}$), potassium ($\text{K}$), sulfur ($\text{S}$), sodium ($\text{Na}$), chlorine ($\text{Cl}$), and magnesium ($\text{Mg}$) constitute most of the remaining $3.7\%$ of body mass.

• Trace elements are required in minute quantities (less than $0.01\%$ of mass); for example, human thyroid activity requires $0.15\,mg$ of iodine ($\text{I}$) daily.

Atomic Structure and Subatomic Particles

• An atom is the smallest unit of matter retaining the properties of an element.

• Key subatomic particles include:

  • Neutrons ($n$): Electrically neutral, mass of approx. $1.7 \times 10^{-24}\,g$ ($1\,dalton$).

  • Protons ($p^+$): Positive charge ($+1$), mass of approx. $1.7 \times 10^{-24}\,g$ ($1\,dalton$).

  • Electrons ($e^-$): Negative charge ($-1$), mass of approx. $\frac{1}{2000}$ of a proton or neutron.

• Protons and neutrons form the atomic nucleus, while electrons form a cloud of negative charge around it.

• Atomic number is the number of protons in an atom's nucleus (and electrons in a neutral atom).

• Mass number is the sum of protons plus neutrons; atomic mass is approximated by the mass number.

• Isotopes are atoms of the same element with different numbers of neutrons. Unstable or radioactive isotopes decay spontaneously, giving off particles and energy (used in fossil dating, metabolic tracing, and medical imaging).

Electron Configuration and Energy Levels

• Potential energy of an electron depends on its distance from the nucleus.

• Electrons occupy discrete electron shells (energy levels); energy is absorbed to move to outer shells and lost when falling to inner shells.

• Chemical reactivity is determined by valence electrons in the outermost shell (valence shell). Completed valence shells render atoms chemically inert.

• Electrons reside in three-dimensional orbitals (up to $2$ electrons per orbital).

Chemical Bonding and Molecular Structure

• Covalent bonds form when atoms share pairs of valence electrons.

  • Single bonds share $1$ pair of electrons; double bonds share $2$ pairs.

  • Nonpolar covalent bonds involve equal electron sharing due to similar electronegativities.

  • Polar covalent bonds involve unequal electron sharing due to differences in electronegativity, creating partial charges ($\delta+$ and $\delta-$).

• Ionic bonds form when one atom transfers electrons to a strongly electronegative partner, resulting in oppositely charged ions (cations and anions) that attract each other.

• Hydrogen bonds are weak attractions between a partially positive hydrogen atom covalently bonded to an electronegative atom (like $\text{O}$ or $\text{N}$) and another electronegative atom.

• van der Waals interactions are weak, fleeting attractions resulting from temporary asymmetric electron distribution.

• Molecular shape dictates biological function and allows specific recognition between biological molecules.

Chemical Reactions

• Chemical reactions make and break chemical bonds, transforming reactants into products while conserving matter.

• Photosynthesis equation:   $6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$

• All chemical reactions are reversible; chemical equilibrium is reached when the forward and reverse reaction rates become equal.

Matter, Elements, and Compounds

• Matter is anything that takes up space and has mass.

• An element is a substance that cannot be broken down into other substances by chemical reactions ($92$ naturally occurring elements).

  • Analogy: Think of elements like primary colors (red, yellow, blue)—you cannot break them down into simpler colors, but you can mix them to create everything else!

• A compound consists of two or more elements combined in a fixed ratio, demonstrating emergent properties distinct from its constituent elements.

  • Analogy: Table salt ($\text{NaCl}$) is made of sodium ($\text{Na}$), an explosive metal, and chlorine ($\text{Cl}_2$), a toxic gas. Combined in a $1:1$ ratio, they form safe, edible table salt!

  • Tip & Trick: Remember that emergent properties mean "the whole is greater than the sum of its parts."

The Elements of Life

• Approximately $20\text{--}25\%$ of natural elements are essential elements needed for an organism to live and reproduce (humans require $25$, plants require $17$).

• Four elements make up $96\%$ (or $96.3\%$) of living matter:

  • Acronym: CHON — Carbon ($\text{C}$), Hydrogen ($\text{H}$), Oxygen ($\text{O}$), and Nitrogen ($\text{N}$).

• Calcium ($\text{Ca}$), phosphorus ($\text{P}$), potassium ($\text{K}$), sulfur ($\text{S}$), sodium ($\text{Na}$), chlorine ($\text{Cl}$), and magnesium ($\text{Mg}$) constitute most of the remaining $3.7\%$ of body mass.

• Trace elements are required in minute quantities (less than $0.01\%$ of mass); for example, human thyroid activity requires $0.15\,mg$ of iodine ($\text{I}$) daily.

Atomic Structure and Subatomic Particles

• An atom is the smallest unit of matter retaining the properties of an element.

• Key subatomic particles include:

  • Neutrons ($n^0$): Electrically neutral, mass of approx. $1.7 \times 10^{-24}\,g$ ($1\,dalton$).

  • Protons ($p^+$): Positive charge ($+1$), mass of approx. $1.7 \times 10^{-24}\,g$ ($1\,dalton$).

  • Electrons ($e^-$): Negative charge ($-1$), mass of approx. $\frac{1}{2000}$ of a proton or neutron (negligible mass).

  • Acronym / Memory Trick:

  • Proton = Positive

  • Neutron = Neutral

  • Electron = Electric / Negative charge

• Protons and neutrons form the central atomic nucleus, while electrons form a cloud of negative charge around it.

  • Analogy: If the atom were the size of Yankee Stadium, the nucleus would be the size of a pencil eraser on the pitcher's mound, and the electrons would be like two gnats buzzing around the stadium. Atoms are mostly empty space!

• Atomic number is the number of protons in an atom's nucleus (and electrons in a neutral atom).

• Mass number is the sum of protons plus neutrons ($\text{Mass Number} = \text{Protons} + \text{Neutrons}$); atomic mass is approximated by the mass number.

• Isotopes are atoms of the same element with different numbers of neutrons.

  • Radioactive isotopes: Unstable isotopes that decay spontaneously, giving off particles and energy. Used in fossil dating, metabolic tracing, and medical imaging (PET scans).

Electron Configuration and Energy Levels

• Potential energy of an electron depends on its distance from the nucleus.

  • Analogy: A ball bouncing down a flight of stairs. It can rest on any step (electron shell) but cannot hover between steps. The higher the step, the higher the potential energy!

• Electrons occupy discrete electron shells (energy levels); energy is absorbed to move to outer shells and lost when falling to inner shells.

• Chemical reactivity is determined by valence electrons in the outermost shell (valence shell).

  • Completed valence shells render atoms chemically inert (nonreactive).

  • Tip & Trick: Group numbers on the periodic table help you quickly identify the number of valence electrons!

• Electrons reside in three-dimensional orbitals (up to $2$ electrons per orbital).

Chemical Bonding and Molecular Structure

• Covalent bonds form when atoms share pairs of valence electrons.

  • Single bonds share $1$ pair of electrons; double bonds share $2$ pairs.

  • Nonpolar covalent bonds: Equal electron sharing due to similar electronegativities.

  • Polar covalent bonds: Unequal electron sharing due to differences in electronegativity, creating partial charges ($\delta+$ and $\delta-$).

  • Analogy:

  • Nonpolar: Two equal-strength twins sharing a toy equally.

  • Polar: A big kid and a little kid sharing a blanket—the bigger, more electronegative kid pulls most of the blanket over to their side!

• Ionic bonds form when one highly electronegative atom strips an electron from another atom, creating oppositely charged ions (cations and anions) that attract each other.

  • Memory Trick:

  • Cation = "Paws-itive" (like a cat) or has a plus sign in the word (cat+ion).

  • Anion = A Negative Ion.

  • Analogy: One atom completely steals the electron cash, becoming negatively charged, leaving the victim positively charged—and now their opposite charges draw them together!

• Hydrogen bonds: Weak attractions between a partially positive hydrogen atom covalently bonded to an electronegative atom (like $\text{O}$ or $\text{N}$) and another electronegative atom.

• van der Waals interactions: Fleeting attractions resulting from temporary asymmetric electron distribution.

  • Analogy / Real World Example: Millions of van der Waals interactions between microscopic hairs on a gecko's toes allow it to walk straight up a smooth glass wall!

• Molecular shape dictates biological function and allows specific recognition between biological molecules (e.g., endorphins and morphine binding to the same brain receptors).

Chemical Reactions

• Chemical reactions make and break chemical bonds, transforming reactants into products while conserving matter.

• Photosynthesis equation: $6\text{CO}_2 + 6\text{H}_2\text{O} \rightarrow \text{C}_6\text{H}_{12}\text{O}_6 + 6\text{O}_2$

• Chemical Equilibrium: Reached when the forward and reverse reaction rates become equal ($\rightleftharpoons$).

  • Tip & Trick: Chemical equilibrium does NOT mean equal amounts of reactants and products; it means the rates of reaction are equal, so product/reactant concentrations remain constant!