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BIOL 2100 CH. 2 NOTES

Concept 2.1: Matter and Chemical Elements

Matter

  • Definition: Anything that takes up space and has mass.

  • Forms: Includes rocks, metals, oils, gases, and living organisms.

Elements and Compounds

  • Elements:

    • Cannot be broken down by chemical reactions.

    • Examples: Gold (Au), Copper (Cu), Carbon (C), Oxygen (O).

    • Each element has a unique symbol (e.g., Sodium = Na from Latin "natrium").

  • Compounds:

    • Substances made of two or more different elements in a fixed ratio.

    • Example:

      • Table salt (sodium chloride, NaCl) = Sodium (Na) + Chlorine (Cl).

      • Water (H₂O) = 2 Hydrogen (H) + 1 Oxygen (O).

    • Emergent Properties: Compounds exhibit characteristics different from their constituent elements.

The Elements of Life

  • Essential Elements:

    • Approximately 25 elements are essential for life.

    • Variation exists among organisms (e.g., humans need 25 elements; plants need fewer).

  • Major Elements in Living Matter:

    • Oxygen (O): 65%

    • Carbon (C): 18.5%

    • Hydrogen (H): 9.5%

    • Nitrogen (N): 3.3%

    • Other significant elements: Calcium (Ca), Phosphorus (P), Potassium (K), Sulfur (S).

  • Trace Elements: Required in minute quantities (e.g., Iron (Fe), Iodine (I)).

    • Example: Iodine is essential for thyroid hormone production; deficiency leads to goiter.

  • Major Elements:

    • Oxygen (O): 65%

    • Carbon (C): 18.5%

    • Hydrogen (H): 9.5%

    • Nitrogen (N): 3.3%

    • Others (Ca, P, K, S, Na, Cl, Mg): ~3.7%

  • Trace Elements: Less than 0.01% of body mass (e.g., Boron (B), Chromium (Cr), Iron (Fe), Zinc (Zn)).

Concept 2.2: An Element’s Properties Depend on the Structure of Its Atoms

Atoms and Elements

  • Atoms: Smallest unit of matter retaining element properties.

  • Element: Composed of a specific type of atom, symbolized by the element's abbreviation (e.g., C for carbon).

Subatomic Particles

  • Types:

    • Protons: Positive charge, located in the nucleus.

    • Neutrons: Neutral charge, also in the nucleus.

    • Electrons: Negative charge, form a cloud around the nucleus.

  • Nucleus: Dense core of protons and neutrons; electrons are attracted to it due to opposite charges.

  • In an uncharged atom the number of protons is equal to the number of electrons, it’s electrically neutral

  • Atomic Number and Mass Number

    • Atomic Number (Z): Number of protons in the nucleus; indicates the element.

    • Mass Number (A): Total number of protons and neutrons.

      • Example: Helium (²He) has 2 protons and 2 neutrons.

  • Calculation of Neutrons: Neutrons = Mass Number - Atomic Number.

Isotopes

  • Definition: Atoms of the same element with different numbers of neutrons.

  • Example: Carbon has three isotopes (¹²C, ¹³C, ¹⁴C).

  • Stability: Stable isotopes do not decay; radioactive isotopes do, transforming into different elements.

Radioactive Isotopes

  • Uses:

    • Medical diagnostics (e.g., PET scans).

    • Tracers in biological research.

  • Radiometric Dating: Measures radioactive decay to date fossils and rocks.

Energy Levels of Electrons

  • Electron Shells: Electrons exist at fixed energy levels; further from the nucleus = higher potential energy.

  • Movement: Electrons can move between shells by absorbing or releasing energy.

  • Electron Distribution and Chemical Properties

    • Valence Electrons: Electrons in the outermost shell determine chemical behavior.

    • Reactivity: Atoms with unpaired valence electrons are more reactive.

Electron Orbitals

  • Orbitals: Regions where electrons are likely found; different shapes (s, p, etc.) and orientations.

  • Shell Capacity:

    • 1st shell: 2 electrons (1s)

    • 2nd shell: 8 electrons (2s) (2 in each orbital)

Concept 2.3: Formation and Function of Molecules and Ionic Compounds

Chemical Bonding

  • Chemical Bonds: Atoms combine to form molecules and ionic compounds through interactions that complete their valence shells.

  • Types of Bonds:

    • Covalent Bonds: Sharing of valence electrons (e.g., H₂).

    • Ionic Bonds: Transfer of electrons resulting in oppositely charged ions (e.g., NaCl).

Covalent Bonds

  • Definition: A covalent bond involves the sharing of a pair of valence electrons between two atoms.

    • Will always be represented with a solid line between atoms

    • Equally shared electrons are non-polar covalent

    • Unequally shared electrons are polar covalent

      Space-filling model of the water molecule, with a large oxygen atom bonded to two smaller hydrogens.

      Examples:

      • Hydrogen (H₂): Two H atoms share one pair of electrons.

      • Oxygen (O₂): Two O atoms share two pairs of electrons (double bond).

      • Water (H₂O): One O atom shares electrons with two H atoms.

      • Methane (CH₄): One C atom shares electrons with four H atoms.

  • Table 2.10, covalent bonding in four molecules.

Representations of Covalent Bonds

  • Molecular Formula: Indicates the number of atoms (e.g., H₂).

  • Electron Distribution Diagram: Shows shared electrons.

  • Lewis Dot Structure: Dots represent valence electrons.

  • Structural Formula: Lines represent shared pairs of electrons.

  • Space-Filling Model: Represents the actual shape of the molecule.

Ionic Bonds

  • Definition: Formed when one atom transfers an electron to another, creating cations and anions.

    • Will be represented by full oppsoite charges, not lines

      • Example: Sodium (Na) transfers an electron to Chlorine (Cl) to form Na⁺ and Cl⁻, resulting in NaCl.'

    Electron transfer and ionic bonding.

  • Properties: Ionic compounds form crystals and do not consist of individual molecules; their formula indicates the ratio of ions.

Weak Chemical Interactions

  • Importance: Weak interactions (e.g., hydrogen bonds, van der Waals interactions) are crucial for biological functions.

  • Hydrogen Bonds: Formed between a hydrogen atom covalently bonded to an electronegative atom and another electronegative atom.

    • Will always be represented by a dotted line btw. atoms

  • Van der Waals Interactions: Occur due to transient charges in molecules, allowing temporary adhesion. Can be very weak individually but together with multiple reactions it can be very powerful

Molecular Shape and Function

  • Shape Determination: Molecular shape is influenced by the arrangement of orbitals during covalent bonding.

    • It determines how biological molecules recognize and respond to one another

    Molecular shapes due to hybrid orbitals.

  • The single s and three p orbitals form four new hybrid orbitals shaped like identical teardrops extending from the region of the atomic nucleus

    • If we connect the larger ends of the teardrops with lines, we have the outline of a geometric shape called a tetrahedron,

  • For water molecules (H2O), two of the hybrid orbitals in the oxygen’s valence shell are shared with hydrogens.

    • The result is a molecule shaped roughly like a V, with its two covalent bonds at an angle of 104.5°.

  • The methane molecule (CH4) has the shape of a completed tetrahedron because all four hybrid orbitals of the carbon atom are shared with hydrogen atoms

Concept 2.4: Chemical Reactions and Bond Dynamics

Overview

  • Chemical Reactions: Processes that make and break chemical bonds, altering the composition of matter.

  • Example: Reaction of hydrogen (H₂) and oxygen (O₂) to form water (H₂O).

    • Bond Dynamics:

      • Breaks covalent bonds in H₂ and O₂.

      • Forms new bonds in H₂O.

        2 H 2 and O 2 are reactants that go through a chemical reaction to yield 2 H 2 O products.

Chemical Equations

  • Reactants: Starting materials (e.g., H₂ and O₂).

  • Products: Resulting materials (e.g., H₂O).

  • Arrow (→): Indicates conversion from reactants to products.

  • Coefficients: Indicate the number of molecules (e.g., 2H₂ means two molecules of hydrogen).

  • Conservation of Matter: Atoms are neither created nor destroyed; they are rearranged.

Photosynthesis

  • Importance: Fundamental biological process for food and oxygen production.

  • Raw Materials: Carbon dioxide (CO₂) and water (H₂O).

  • Process:

    • Sunlight powers the conversion of CO₂ and H₂O into glucose (C₆H₁₂O₆) and oxygen (O₂).

    • By-product: Oxygen gas released by aquatic plants (e.g., Elodea).

  • Energy Input: Sunlight is essential for the rearrangement of matter, it powers the conversion of carbon dioxide and water into glucose and oxygen

    Reactants 6 C O 2 and 6 H 2 O go through a reaction to form C 6 H 12 O 6 + 6 O 2.

Chemical Reversibility

  • Reversible Reactions: Products can revert to reactants (e.g., ammonia formation and decomposition).

  • Dynamic Equilibrium:

    • Forward and reverse reactions occur at the same rate.

    • Concentrations of reactants and products stabilize at a specific ratio.

    • Equilibrium does not imply equal concentrations, just a stable ratio.

Factors Affecting Reaction Rates

  • Concentration of Reactants: Higher concentration increases collision frequency, enhancing reaction rates.

  • Accumulation of Products: Leads to increased reverse reaction frequency until equilibrium is reached.

Summary

  • Chemical reactions are essential for life, involving the rearrangement of atoms and the conservation of matter.

  • Understanding these processes lays the groundwork for studying the molecules vital to biological functions, such as water.

Chapter 2 Review Notes

Summary of Key Concepts

Concept 2.1: Matter and Chemical Elements

  • Elements: Cannot be broken down chemically; consist of pure forms.

  • Compounds: Combinations of two or more different elements in a fixed ratio.

  • Key Elements in Living Matter: Oxygen, carbon, hydrogen, nitrogen.

Concept 2.2: Atomic Structure and Properties

  • Atoms: Smallest unit of an element, composed of:

    • Nucleus: Contains protons (positive charge) and neutrons (neutral).

    • Electrons: Negative charge, form a cloud around the nucleus.

  • Atomic Number: Number of protons; defines the element.

  • Atomic Mass: Measured in daltons; roughly equal to the mass number (protons + neutrons).

  • Isotopes: Variants of an element with different neutron numbers; unstable isotopes exhibit radioactivity.

  • Electron Shells: Electrons occupy specific energy levels; incomplete outer shells indicate reactivity.

  • Orbitals: Three-dimensional spaces within shells that define electron distribution.

    An atom is diagrammed.

Concept 2.3: Chemical Bonding

  • Chemical Bonds: Form when atoms interact to complete valence shells.

    • Covalent Bonds: Formed by sharing electrons.

      • Single Covalent Bond: Sharing one pair of electrons (e.g., H₂).

      • Double Covalent Bond: Sharing two pairs of electrons (e.g., O₂).

  • Electronegativity: Attraction of an atom for electrons in a bond.

    • Nonpolar Covalent Bonds: Equal sharing of electrons.

    • Polar Covalent Bonds: Unequal sharing, leading to partial charges.

  • Ionic Bonds: Formed through the transfer of electrons, resulting in charged ions (e.g., NaCl).

  • Weak Interactions: Include hydrogen bonds and Van der Waals interactions, crucial for molecular shape and adhesion.

  • Molecular Shape: Determined by valence orbital positions; essential for biological recognition.

    The lone valence electron of a sodium atom is transferred to join the 7 valence electrons in the third shell of a chlorine atom to form Sodium chloride (N a C l).

  • Two reactions are given.

Concept 2.4: Chemical Reactions

  • Chemical Reactions: Transform reactants into products while conserving matter.

  • Reversibility: All reactions are theoretically reversible.

  • Chemical Equilibrium: Achieved when forward and reverse