The Nature of Science

• Science is a human and social activity with unique characteristics and a long history.
• Science is a way of learning about the universe based on curiosity, creativity, imagination, intuition, exploration, observation, replication of experiments, interpretation of evidence, and debate.
• Scientific activity provides a conceptual and theoretical base for predicting, interpreting, and explaining natural and human-made phenomena.
• Many historians, sociologists, and philosophers argue that there is no set procedure for conducting a scientific investigation

Driving Forces of Science

• Science is driven by a combination of theories, knowledge, experimentation, and processes anchored in the physical world.
• Theories are continually tested, modified, and improved as new knowledge and theories emerge.
• Scientific debate involves many participants with diverse backgrounds, fueled by theoretical discussions, experimentation, and social, cultural, economic, and political influences.

Basic Chemistry Review

• All matter is composed of atoms, which are collections of subatomic particles.

Atomic Structure

• Atoms have a nucleus at their core surrounded by electrons.
• The nucleus contains protons (positively charged) and neutrons (no charge).
• Electrons are much smaller and negatively charged.
• The atomic number, which defines the type of atom, is the number of protons in the nucleus.
• The mass number is the total number of protons and neutrons in the nucleus.
• A neutral atom has the same number of electrons surrounding the nucleus as there are protons.
• Isotopes are different atoms that have different numbers of neutrons.

Isotopic Mass Number

• The isotopic mass number is the weighted average of isotopic masses for a collection of atoms.

Isotopes

• Isotopes are a particular element with the same number of protons (atomic number) but a different atomic mass due to a change in the neutron number.

Carbon-12

• Carbon-12 accounts for 99% of the carbon atoms in nature and most isotopes are stable but decay over time.

Radioisotopes

• Radioactive isotopes, also known as radioisotopes, have unstable nuclei that break apart.
• Radioactive isotopes release subatomic particles and some radiation when they break apart.
• Carbon-14 is a radioactive isotope that decays into Nitrogen-14.

Common Uses for Radioisotopes

• Radioisotopes have various applications:
• Carbon-14 is used for dating and tracing cancer
• Calcium-45 measures bone formation rate
• Potassium-40 is used for fossil dating
• Radium-226 is used in cancer treatment

Valence Electrons

• Valence electrons are involved in bonding between atoms, possessing the highest energy
• Valence electrons are the 'outermost' electrons of an atom.

Core Electrons

• Core electrons are chemically inert under normal reaction conditions
• Core electrons are located in filled energy levels of an atom

Chemical Bonding

• Intramolecular bonds are formed between atoms through the interaction of valence electrons.
• Intramolecular bonds can be ionic, covalent, or polar covalent.

Ionic Bonds

• Ionic bonds involve the exchange or transfer of electrons from one atom to another
• Bonding occurs to make atoms more stable, affecting the substance's properties.
• Ionic bonds typically form between a metal and a nonmetal.
• Atoms that lose electrons become positively charged (cations), and atoms that gain electrons become negatively charged (anions), opposite charges attract each other
• H+ and Na+ are both critical ions for the body

Covalent Bonds

• Covalent bonds involve the sharing of electrons between atoms, usually two non-metals
• Covalent bonds aim to achieve a stable electron configuration (octet)

Electronegativity

• Electronegativity indicates the relative ability of bonding atoms to attract a shared electron pair in a covalent bond
• If two atoms of the same element are bound, their electronegativities are the same, resulting in a pure non-polar covalent bond

Electron Sharing

• Unequal electron sharing between atoms results in regions of partial positive (δ+) and partial negative (δ-) charge

Bond Classification

• If the electronegativity difference between two atoms is between 0 and 1.7 (but less than 1.7), the bond is polar covalent.
• If the electronegativity difference is 1.7 or greater, the bond is considered ionic.

Intermolecular Bonds/Forces (Van Der Waals Forces)

• Interactions between molecules that hold two or more molecules together.
• Intermolecular bonds are much weaker than intramolecular bonds and depend on the state of the substance.

London Dispersion Forces

• Weak attractive forces between all atoms and molecules
• London Dispersion Forces are due to temporary unequal distribution of electrons as the electron cloud randomly moves about the nucleus
• London Dispersion Forces hold small nonpolar molecules together
• CH4 (methane) is a gas at room temperature due to weak London Forces
• C8H18 (octane) is liquid at room temperature due to the cumulative effect of many London forces

Dipole-Dipole Forces

• Dipole-Dipole Forces hold polar molecules together
• Dipole-Dipole Forces are stronger than London forces
• Dipole-Dipole Forces occur when the partially positive side of a polar molecule is attracted to the partially negative side of the molecule

Hydrogen Bonding

• A strong dipole-dipole force that forms between the electronegative H of one polar molecule and the electronegative N, O, or F of another
• Hydrogen bonding is seen in water molecules
• Hydrogen bonds are individually weaker than covalent or ionic bonds but become strong when many are added together
• Hydrogen bonds can have a strong cumulative effect such as, the specific heat of water, and the structure of DNA.

Properties of Water

• List three properties of water and explain each using examples:
• Cohesion: Water molecules form hydrogen bonds with one another leading to high surface tension
• Adhesion: Water molecules form hydrogen bonds with other polar materials, causing capillary action (e.g., transpiration)
• Temperature moderation: Hydrogen bonding causes water to absorb/lose large amounts of heat

Properties of Water

• High Specific Heat of Vaporization: Hydrogen bonding allows liquid water to absorb a large amount of heat to become vapour which allows evaporative cooling for organisms
• Highest Density at 4°C: Water molecules form a crystalline lattice below 0°C, H-bonds spread apart, decreasing density, so ice floats

Water as a Solvent

• What makes water a good solvent and why is it important for all of life?: Water's polarity allows it to dissolve substances easily, dissociating ionic solids into anions and cations (spheres of hydration)
• Polar molecules readily dissolve in water

Solubility

• Substances that can dissolve in water are soluble
• Substances that cannot dissolve in water are insoluble
• Miscible describes liquids that dissolve into one another
• Immiscible describes liquids that form separate layers

Aqueous solutions

• Aqueous solutions pertains to a molecule or ion dissolved in water.
• Solutions are made of a solute (lesser quantity) and a solvent (greater quantity)

Hydrophilic substances

• Polar substances will dissolve in water and form hydrogen bonds and are hydrophilic (e.g., salts)
• Nonpolar substances will not dissolve in water or form hydrogen bonds and are hydrophobic (e.g., fats and oils)

Ionization of Water and pH

• Aqueous solutions can be acidic, basic, or neutral.
• Acidic solutions have a high concentration of hydrogen ions (H+)
• Basic solutions have a high concentration of hydroxide ions (OH-)
• Neutral solutions have an equal amount of hydrogen and hydroxide ions

Autoionization

• Pure water undergoes autoionization, where two water molecules react to produce a hydronium ion (H3O+) and a hydroxide ion (OH-)

Acids and Bases

• Acids are substances that increase the concentration of H3O+ ions (aq) when dissolved in water
• Acids turn litmus paper red.
• Strong acids like HCl ionize completely.
• Weak acids, like vinegar (CH3COOH), only partially ionize

Bases

• Substances that increase the concentration of OH- (aq) ions when dissolved in solution
• Bases turn litmus paper blue.
• Strong bases like NaOH ionize completely
• Weak bases, like ammonia (NH3), only partially ionize.
• Acids react with bases to produce water and a salt in a neutralization reaction

Neutrality and pH scale

• Pure water has the same number of hydronium and hydroxide ions, so it is neutral
• The pH scale is defined as pH = -log10[H3O+]
• Pure water at 25°C has a [H3O+] of 1.0 x 10-7 mol/L, with a pH of 7

Conjugate Acid and Bases

• An acid is a proton donor
• A base is a proton acceptor
• The direction of the reaction determines what substance is acting as the acid or base.
• pH levels within a cell must be maintained within a narrow range for cellular activity

Buffers

• Chemical systems that resist changes in pH
• Donate H+ ions when required and remove H+ ions when there are too many
• Carbonic acid (acid) - bicarbonate (base) buffer in human blood (pH 7.4)

Main Classes of Macromolecules

• The main classes of macromolecules are:
• Carbohydrates
• Lipids
• Proteins
• Nucleic acids
• Minerals
• Vitamins