Chemistry Fundamentals Quiz

Introduction to Chemistry

• This course, NSG 741, focuses on the interplay of genetics, chemistry, and physics within anesthesia.

Why Chemistry and Physics?

• Anesthesia practice relies on principles of chemical and physical science.
• Chemistry studies matter's composition, properties, and structure (atomic and molecular levels), explaining how matter behaves in reactions.
• Physics describes matter's motion, mechanics, force, and energy, examining its behavior in space and time.
• These interconnected fields explain fundamental anesthesia processes.

General Chemistry

• The universe comprises matter and energy.
• Energy pertains to physics.
• Matter exists as solid, liquid, gas, or plasma.
• Solids maintain shape and volume.
• Liquids maintain volume but not shape; they exhibit minimal to no compressibility and volume changes with pressure and temperature.
• Gases assume no fixed shape or volume, easily changing volume with pressure and temperature variations.

What is Chemistry?

• Chemistry studies matter and its changes.
• Physical chemists develop models to understand chemical systems theoretically.
• Inorganic chemists study non-carbon-based substances.
• Organic chemists focus on carbon-based compounds.
• Biochemists study chemical processes in living systems.
• Chemical changes produce new substances, while physical changes alter state without changing substance identity.

States of Matter

• Three common states: solids (definite shape and volume), liquids (definite volume, indefinite shape), and gases (indefinite shape and volume).
• Transitions between states: melting (solid to liquid), freezing (liquid to solid), vaporization (liquid to gas), condensation (gas to liquid), deposition (gas to solid), sublimation (solid to gas).

The Atom

• Atoms are basic building blocks of matter, composed of protons, neutrons, and electrons.
• Protons (positively charged) approximate 1 atomic mass unit (amu).
• Neutrons (no charge) approximate 1 amu.
• Electrons (negatively charged) have significantly smaller mass.
• Atomic number (Z): number of protons in the nucleus (identifies the element).
• Mass number (A): total protons and neutrons in the nucleus.
• Atomic weight: average atomic mass of naturally occurring isotopes, in amus.

Dalton's Atomic Theory

• Postulates of John Dalton's atomic theory encompass:

• Elements are composed of indivisible atoms, all identical within an element.

• Compounds form when different atoms combine in fixed proportions.

• Chemical reactions involve atom rearrangement, not creation or destruction.

• Dalton's atomic theory rests on two significant laws:

• Law of Conservation of Mass (total mass doesn't change in a chemical reaction).

• Law of Definite Proportions (pure compounds always contain the same elements in the same ratio by mass).

• Dalton theorized multiple proportions that relate composition variation in chemical compounds.

### The Periodic Table of the Elements

• Modern periodic table organizes elements based on recurring chemical and physical properties.
• Each box shows the element's symbol, atomic number, and average atomic mass.
• Elements are ordered by increasing atomic number (number of protons).
• Periods (rows) correspond to electron additions to quantum energy levels.
• Vertical columns (groups/families) define similar chemical and physical properties.
• Group 8A (18) represents the noble gases, characterized by inertness.

Average Atomic Weights & Classifications

• Atomic weights on the periodic table are weighted averages of the masses of naturally occurring isotopes.
• Elements classified as representative, transition, or inner transition.
• Representative elements have numbered group A designations.
• Transition elements are shown with group B designations.
• Inner transition elements are positioned at the table's bottom.

Metals vs. Nonmetals

• Metals are typically solids, ductile, malleable, and conduct heat and electricity, readily ionizing.
• Nonmetals can be solids, liquids, or gases and are typically poor conductors, with a tendency to gain electrons.

Ions & Elements

• Ions are charged atoms (or groups of atoms) formed by adding or removing electrons.
• Cations (positively charged) are formed by electron loss.
• Anions (negatively charged) are formed by electron gain.
• Elements contain only a single type of atom.
• Ionization energy is the minimal energy to remove the least tightly bound electron.

Compounds

• Compounds contain two or more different types of atoms bonded together.
• Two major bonding types: covalent (molecular) and electrostatic (ionic).
• Covalent bonds involve shared electrons; compounds exist as discrete molecules.
• Electrostatic bonds involve oppositely charged ions; compounds exist as a continuous lattice.

Mass Spectrometer

• Mass spectrometry gauges an atom or molecule's mass.
• The sample is introduced, vaporized, and ionized (loss of one electron) to measure by the magnetic field strength and determine atomic or molecular mass.

Moles

• A mole represents exactly 6.02 × 10²³ particles (atoms, molecules, or ions). Known as Avogadro's number.
• Moles are used to relate mass and number of particles.

Molar Mass (Molecular Weight)

• Molar mass equals the sum of atomic masses of elements and corresponds numerically to the element's atomic mass in g/mol.

What is a Mole?

• A mole is a unit denoting 6.02 × 10²³ particles.
• Carbon-12 has 1 mole of atoms in 12 grams.
• Moles establish a link between particle counts and substance mass.

Molar Mass and Atomic Mass

• Numerically, molar mass (g/mol) equals the element's atomic mass (amu).
• For molecules, molar mass is the total atomic mass in g/mol.

Molar Mass of Molecules

• Molar mass of molecules is calculated by summing individual atomic masses; it's expressed in g/mol.

Comparing Moles to Dozens

• Moles are analogous to dozens, but on a vastly larger scale.
• A mole of any substance contains 6.02 × 10²³ particles.

Moles as Conversion Factors

• Moles serve as conversion factors between number of particles (atoms, molecules) and mass (grams) of substances.

Isotopes

• Atoms with the same element but different neutron numbers.
• Same atomic number but different mass number.

Isomers

• Molecules with the same chemical formula but different structural arrangements.
• Includes structural isomers, differently arranged components but identical formulas, and stereoisomers, differing in the spatial arrangement of atoms.

Isomer Example

• Example of isomer structural differences in C₂H₆O (ethanol and dimethyl ether) with same molecular formula.

Enantiomers

• Mirror images of one another; cannot be superimposed, with similar but not identical properties.
• Optically active and can rotate polarized light; labeled + (dextro) or - (levo)
• Racemic mixtures contain equal amounts of both enantiomers, often encountered in pharmaceuticals.

Physical and Chemical Properties and Changes

• Physical changes alter a material's state without changing its fundamental chemical composition.
• Chemical changes transform a substance into a fundamentally different substance.
• Intensive properties are independent of the sample size (e.g., color).
• Extensive properties are dependent on sample size (e.g., mass).
• Chemical properties describe how a substance reacts or transforms.

Chemistry Concepts: Pure Substances and Mixtures

• Pure substances cannot be physically separated into simpler components, while mixtures are combinations of two or more pure substances.
• Homogeneous mixtures exhibit uniform chemical and physical properties (uniform throughout), whereas heterogeneous mixtures show distinct phase boundaries (different areas).

Intermolecular Forces

• Intermolecular forces affect macroscopic properties like evaporation tendencies.
• They are electrostatic in nature.
• Coulomb's Law describes how force increases as charges get closer together.

Lewis Structures

• These are nonmathematical models that qualitatively depict chemical bonding and associated properties.
• Atoms are represented using chemical symbols.
• Shared electron pairs (bonds) are depicted as lines.
• Remaining valence electrons are lone pairs (dots).

Balancing Chemical Equations

• Balancing chemical equations ensures equal numbers of each atom type on both sides of the equation.

Valence Shell Electron Pair Repulsion (VSEPR) Theory

• Repulsions between electron pairs determine molecular shapes (linear, trigonal planar, tetrahedral, bent, trigonal pyramidal, etc.)

Valence Bond Theory

• Covalent bonding results from electron sharing by overlapping valence electron clouds and encompassing both nuclei.

Bond Length and Bond Energy

• Ideal bond length minimizes energy and maximizes molecular stability.

Ionic Bonds

• Strong electrostatic attraction between oppositely charged ions, forming a crystal lattice. -Complete transfer of electrons between atoms results in ions.

Covalent Bonds

• Covalent bonds form when atoms share electrons, satisfying the octet rule to achieve stability.

Electronegativity and Polar Covalent Bonds

• Electronegativity measures an atom's tendency to attract electrons to itself, higher electronegativity for nonmetals.
• Polar covalent bonds occur in compounds with varying electronegativity, with shared electrons closer to the atom with higher electronegativity.

Polar Covalent Bonds

• Shared electrons in a polar covalent bond tend to be drawn toward the more electronegative atom, creating a dipole (polar molecule), uneven distribution of charges.
• Hydrophilic or hydrophobic tendencies of molecules are determined based on this property.

Electrostatic Bonds

• Electrostatic bonds form through the attraction of opposite charges between atoms, playing a role in ionic bonding and other interactions.

Ion-lon Bonding

• Strongest electrostatic bonds result from attraction between oppositely charged ions, not directional, and are typically solid at room temperature and pressure.

lon-Dipole Bonding

• Attractions between ions and polar molecules, allowing ionic solids to dissolve in water.

Hydrogen Bonding

• Attractions between hydrogen atoms in one molecule and highly electronegative atoms (like oxygen or nitrogen) in another molecule, often causing a high surface tension in substances that have hydrogen bonding.

Van der Waals Forces: London Dispersion Forces & Dipole-Dipole Attractions

• Weak intermolecular forces involving temporary charge imbalances within molecules.
• Common in molecules with nonpolar bonds.

Molecular Bond Strength

• Covalent bonds are stronger than ionic and polar covalent bonds, with hydrogen bonds and Van der Waals forces being the weakest.
• Biological systems depend significantly on hydrogen bonds, often the dominant intermolecular attraction in water-based solutions.

Bond Breaking

• Bond energy corresponds to the energy required to sever bonds.
• Energy is released during bond creation, and consumed during bond rupture.

Enthalpy

• Enthalpy encompasses the total energy within a system, including kinetic and potential energy.
• Involves energy in relation to height or position, as well as energy stored in bonds.
• Energy-related changes in chemical processes are calculated as change in enthalpy (ΔH).

Electrolytes

• Substances that form solutions capable of conducting electricity; typically ionic compounds that readily dissolve into separate ions when dissolved in water.

Hydrocarbons and Functional Groups

• Hydrocarbons are compounds made of carbon and hydrogen.
• Functional groups dictate a molecule's specific properties and often dictate chemical reaction behaviour.

Alkanes

• Simplest hydrocarbons with only single bonds.
• Classified as saturated because they have the maximum number of hydrogen atoms possible per carbon.

Alkyl Groups

• Alkyl groups result from removing hydrogen from alkanes.
• They are very reactive.

Cycloalkanes & Saturation

• Cycloalkanes consist of carbon atom rings, potentially with differing hydrogen atom counts per carbon in a ring structure.
• The total number of hydrogen atoms present is dependent on the rings presence or lack thereof and is less than in alkanes.

Alkenes

• Hydrocarbons containing carbon-carbon double bonds.
• The geometry around alkene carbons is trigonal planar.

Alkynes

• Hydrocarbons consisting of carbon-carbon triple bonds.
• Relatively rare in medical settings, often named by adding the suffix "yne" to the IUPAC parent name.

Aromatics

• Cyclic compounds showing unique stability characteristics from delocalized electrons; the benzene ring structure (and derivatives) is very common in organic chemistry, as well as biologically significant molecules.
• Often form parts of larger molecules or function as a group in their own right.

Organohalogen Compounds

• Organic compounds containing halogen substituents (F, Cl, Br, I).
• Nomenclature is derived by naming the parent alkane or benzene ring, and then designating where the halogen substituent is located.

Alcohols

• Organic compounds with hydroxyl (-OH) groups covalently bonded to a carbon atom.
• Relatively soluble in water when the carbon count is lower due to the hydroxyl groups making hydrogen bonds with water molecules.

Ethers

• Organic compounds with an oxygen atom bridging between alkyl groups.
• Often used as solvents or to protect other reactive groups as well.

Amines

• Organic nitrogen-based functional group derivatives of ammonia.
• Contain nitrogen atoms covalently bonded to one or more alkyl groups or other alkyl substituents or hydrogens

Carbonyl Functional Groups

• Carbonyl groups consist of a carbon atom doubly bonded to an oxygen atom (C=O).
• Representative of many important organic compounds.

Aldehydes

• Carbonyl groups bonded to one hydrogen and one alkyl group.
• Frequently encountered in the context of carbohydrates.

Ketones

• Carbonyl groups bonded to two alkyl groups.
• Plentiful in nature and many medical treatments.

Reactions of Aldehydes and Ketones

• Oxidation involves increasing oxygen bonds and/or decreasing the number of hydrogen bonds.
• Reduction is the reverse of oxidation, decreasing oxygen bonds and increasing hydrogen bonds.

Formation of Acetals and Ketals

• Aldehydes and ketones react with alcohols to form acetals and ketals.

Carboxylic Acids

• Carboxylic acid functional groups have a carbonyl group and hydroxyl groups on the same carbon, thus representing a polar group.
• They can act as acids (donate a proton).

Esters

• Condensation products between carboxylic acids and alcohols, often having pleasant fruity odor.

Amides

• Condensation products between carboxylic acids and amines, frequently found in proteins and other biological molecules.

Test your knowledge on essential chemistry concepts, including types of elements, chemical bonds, and unique characteristics of ionic and molecular compounds. This quiz covers various topics relevant to chemistry, ensuring a thorough understanding of foundational principles.