AL

MATTER & ENERGY - 100 Flashcards

Matter & Energy

  • Energy is the exertion of force (kinetic) or capacity (potential) to do work
    • Unit: J (joule)
  • Matter is tangible composition that may be solid, liquid, gas, or plasma
    • Solids: resist changes in shape and volume
    • Liquids: fluids with minimal to no compressibility; may change volume with changes in pressure and temperature
    • Gases: fluids that are compressible and easily change volume with changes in pressure and temperature
    • Plasma: mixture of ionized gas & free-floating electrons

Atom vs. Molecule

  • Atom
    • Smallest unit of an element that retains its properties
    • Can exist independently; singular entity
  • Molecule
    • Group of two or more atoms bonded together; smallest unit of a compound
    • Formed when atoms bond
    • Can consist of the same type of atoms (e.g., O2) or different types (e.g., H2O)

Atomic Structure

  • Atom: smallest particle of an element that retains characteristics; smallest unit that can enter chemical reactions
  • Subatomic particles:
    • Protons (positive charge)
    • Neutrons (electrically neutral)
    • Electrons (negative charge)

Properties of Subatomic Particles

  • Electrons: negative; almost massless; primarily determine the size of the atom
  • Protons: positive; located in the nucleus; ~1000× more massive than electrons
  • Neutrons: neutral; located in the nucleus; similar in mass to protons
  • Mass and Charge
    • Protons + Neutrons ≈ most of the atom's mass
    • Electrons ≈ charge and size

Atomic Structure (continued)

  • Protons & neutrons cluster in the nucleus
  • Electrons move around the nucleus in shells (Orbital Theory)
    • Negative charge of electrons attracted to positive nucleus
  • Outer-shell electrons are called valence electrons

Chemical Symbols

  • Chemical elements and electronic configuration are designated by chemical symbols
  • About 26 elements are commonly found in living cells

Molecular Structure: Naming

  • Systematic naming rules:
    • Prefix multiplier indicates how many of each element are present
    • Suffix “-ide” to the last element name
  • Examples:
    • N_2O
      ightarrow ext{Dinitrogen monoxide (nitrous oxide)}
    • NO
      ightarrow ext{Nitrogen monoxide (nitric oxide)}
    • CCl_4
      ightarrow ext{Carbon tetrachloride}
  • Some compounds are non-systematic:
    • H_2O
      ightarrow ext{Water}
    • NH_3
      ightarrow ext{Ammonia}
    • CH_4
      ightarrow ext{Methane}
    • C3H8
      ightarrow ext{Propane}

Molecular Bonding

  • Chemical concepts:
    1) Types of bonds formed between atoms
    2) Polarity: molecule acts like a magnet; uneven electron distribution creates partial charges
    3) Spin of electrons/protons: axis alignment can contribute to magnetic interactions

Comparing Bonds

  • Covalent Bonds
    • Atoms share electrons
    • Strong, stable bonds
    • Directional: electrons shared between specific atoms
  • Electrostatic (Ionic) Bonds
    • Atoms transfer electrons
    • Very strong attraction, but can be disrupted in water
    • Non-directional: attraction occurs around the ion

Covalent Bonds (details)

  • Atoms form bonds by sharing electrons
    • Nonpolar covalent bonds: equal sharing of valence electrons
    • Polar covalent bonds: unequal sharing (one atom pulls harder)
  • Bond types by electron sharing:
    • Single bond: one pair of electrons shared
    • Double bond: two pairs shared
    • Triple bond: three pairs shared

Types of Electrostatic Bonds

  • Ion–Ion (Ionic) bonds
  • Ion–Dipole bonds
  • Dipole–Dipole bonds

Electrostatic Bonds: Ionic

  • Attraction of electrons between atoms
  • Opposites attract (negative to positive)
  • Ionic bonds (ion–ion) are strong
  • Not directional; occur anywhere along outer electron shell
  • Typically have high melting and boiling points

Electrostatic Bonds: Ion–Dipole

  • An ion interacts with a molecule that has a partial charge
  • Weaker than ion–ion bonds
  • Example: water molecule bonding

Electrostatic Bonds: Dipole–Dipole

  • Form between molecules with dipolar (partial) charges
  • Creates weak attractions; responsible for properties like surface tension of water
  • Water is a prime example due to polarity
  • Molecules with uneven electron distribution can form induced dipoles
  • London dispersion forces: weakest type of molecular bond; temporary, weak bonds
  • These forces permit gases (e.g., O₂, N₂) to liquefy at very low temperatures

Bond Breaking

  • Energy to break or form a bond
  • Energy released when a bond is formed; energy consumed when a bond is broken
  • Involved in chemical reactions; amount released equals amount consumed
  • Covalent bonds create greater bond energies than electrostatic bonds
  • Bond energies are measured as an enthalpy change: riangle H

Enthalpy

  • Definition: total amount of energy possessed by a system (kinetic + potential)
  • Very difficult to measure directly; we discuss change in enthalpy: riangle H
  • Heat evolved (exothermic) or absorbed (endothermic) equals the change in enthalpy
  • Examples:
    • Food metabolism → energy available for immediate or later use
    • Fat breakdown → releases stored energy as fuel
    • ATP → ADP + P → consumes energy in the process

Functional Group Overview

  • Hydrocarbons: compounds consisting of only hydrogen and carbon
    • Saturated hydrocarbons: single bonds in a chain
    • Example: Alkanes (e.g., Halothane is a halogenated alkane)
    • Cyclic hydrocarbons: saturated in a ring
    • Examples: Cycloalkane
    • Unsaturated hydrocarbons: double or triple bonds in chains or rings
    • Alkenes (double bonds)
    • Alkynes (triple bonds)
  • Carbonyl compounds: esters and amides

Organic Compounds: Hydrocarbons

  • Hydrocarbons: ext{H}3 ext{C} - ext{CH}2 ext{ - CH}2 - ext{CH}3
  • Structure: straight chains with or without branches
  • Saturated hydrocarbons (alkanes): all remaining carbon bonds are single, each carbon bonded to hydrogens
  • Straight-chain molecule example

Organic Compounds: Hydrocarbons (branched and unsaturated)

  • Branched hydrocarbon example: ext{CH}3- ext{CH}( ext{CH}3)_2
  • A six-carbon hydrocarbon is called hexane
  • Unsaturated hydrocarbons have double or triple bonds
    • Alkenes contain double bonds (e.g., hexene)
    • Alkynes contain triple bonds

Organic Compounds: Cyclic hydrocarbons

  • Carbon chains in a ring structure
  • May contain single, double, or triple bonds
  • Examples: hexane (cyclic form) and benzene (aromatic)

Functional Groups: Amines/Amides

  • Amines: derivatives of ammonia ( ext{NH}3) with general formula ext{NR}3
    • With one or two with hydrogen substituents can be present
    • All amines have a lone pair of electrons on the nitrogen
  • Amides: (carboxamide group) ext{RCONH}2, ext{RCONHR}, ext{RCONR}2
  • Note: Further details referenced from additional lectures

Functional Groups: Alcohol

  • General formula: ext{ROH}
    • R = alkyl group (carbon and hydrogen)
    • ext{OH} = hydroxyl group
  • Hydroxyl group is very polar (binds with ext{H}^+)
  • Hydrophilic; alcohols dissolve many polar molecules

Functional Groups: Phenols

  • Similar to alcohol: ext{ROH}, but R is an aryl group (aromatic ring, e.g., benzene)
  • OH group is polar
  • Simple phenols are polar; complex phenols may be lipid soluble (e.g., propofol)

Functional Groups: Ethers

  • General formula: ext{R-O-R'}
    • R and R' are alkyl groups attached by oxygen
  • Ethers are relatively inert and highly flammable
  • Straight ethers are not used much anymore; halogen substitution can alter blood solubility, potency, and flammability
    • Fluorine (Sevoflurane) substitutions; Chlorine (Isoflurane) substitutions; replace Hydrogen

Functional Groups: Other general formulas

  • Aldehydes: ext{RCHO}
  • Esters: ext{RCOOR}
  • Ketones: ext{RCOR'}
  • Carboxylic acids: ext{RCOOH}
    • Weak acids; form carbonates and bicarbonates (buffers) in pH systems

Vapor Pressure

  • Volatile agents in a vaporizer; liquids turned into vapor; vaporization occurs
  • In a closed container: molecules escape liquid phase and become vapor, striking container walls
  • Vapor pressure is directly correlated with temperature
  • Increasing temperature raises the ratio of gas to liquid molecules, increasing vapor pressure

Vapor Pressure: Boiling Point

  • Boiling point: temperature at which ext{Vapor pressure} = ext{Atmospheric pressure}
  • Atmospheric pressure at sea level: P_{ ext{atm}} = 760 ext{ mmHg}
  • Boiling points for anesthetic agents at standard conditions:
    • Sevoflurane: 58.5^ ext{°C}
    • Desflurane: 22.8^ ext{°C}
    • Isoflurane: 48.5^ ext{°C}
    • Enflurane: 56.5^ ext{°C}
    • Halothane: 50.2^ ext{°C}

Vapor Pressure: 20°C Values

  • At 20°C, vapor pressures (mmHg):
    • Sevoflurane: 170
    • Enflurane: 172
    • Isoflurane: 240
    • Halothane: 244
    • Desflurane: 669

Vapor Pressure: Temperature Dependence and Clinically Relevant Concepts

  • Vapor pressure is a function of temperature
  • Henry’s Law context for gases in liquids; partial pressure effects in closed systems
  • Example problem framework: if a gas is added to a container with air (760 mmHg total), the gas fraction is determined from its vapor pressure relative to total pressure
  • In practice: calculate gas percentages from vapor pressure (VP) and atmospheric pressure (760 mmHg)

Latent Heat of Vaporization

  • Definition: energy needed to convert 1 g of liquid to vapor at constant temperature
  • Effect on volatile agents: vaporization cools the liquid and decreases vapor pressure
  • All common anesthetics have similar latent heats
  • Desflurane exception: less potent (requires more molecules for same depth) → greater temperature loss → larger drop in vapor pressure; may require specially heated vaporizers (TEC) to offset cooling

Desflurane: Special Vaporizer Considerations

  • Properties: high volatility + moderate potency
  • Standard variable-bypass vaporizers are not suitable
  • Special design: heated to 39^ ext{°C}, raising vapor pressure to about 1300 ext{ mmHg}
  • Delivery mechanism: injected directly into fresh gas flow to ensure accurate, safe delivery

Vapor Pressure: Is VP a function of volume, temperature, or pressure?

  • Answer: Temperature governs VP in practice; at constant temperature, VP is a property of the liquid/gas pair
  • Henry’s Law context: at constant T, amount dissolved in liquid is proportional to partial pressure of gas above the liquid
  • Clinical implications for volatile anesthetics in closed containers and Dalton’s Law of partial pressures

Dalton's Law of Partial Pressures

  • Statement: Total pressure of a gas mixture equals the sum of the partial pressures of each gas
  • Example components: ext{O}2, ext{N}2, ext{Ar}, ext{H}2 ext{O}, ext{CO}2
  • Typical air composition (atmosphere): ext{O}2 = 20.9 ext{%}, ext{N}2 = 78.1 ext{%}, ext{Ar} = 0.97 ext{%}, ext{H}2 ext{O} = 1.28 ext{%}, ext{CO}2 = 0.05 ext{%}
  • Total: P_{ ext{total}} = 101.3 ext{ kPa} (or 760 mmHg at sea level)
  • Concept: pressure exerted on container walls is the sum of partial pressures; depends on gas mole fractions and temperature

Dalton’s Law Example: Volatile agent in oxygen flask

  • Scenario: isoflurane is added to a flask of oxygen (P_total = 760 mmHg)
  • Partial pressures: P{ ext{O}2} = 760 - P{ ext{isoflurane}} and P{ ext{isoflurane}} = ext{VP}_{ ext{isoflurane}}
  • If VP of isoflurane = 240 mmHg, then:
    • ext{% Iso} = rac{240}{760} imes 100 = 31.6 ext{%}
    • ext{% O}_2 = rac{760 - 240}{760} imes 100 = 68.4 ext{%}
  • Note: the two percentages sum to 100%

Dalton’s Law: Temperature Effects on Gas-Liquid Equilibria

  • Heating a gas-liquid system increases dissolved gas escape to the gaseous phase
  • Cooling increases dissolution of gas into the liquid
  • Clinical relevance: hypothermia slows emergence from volatile anesthetics due to increased gas solubility and slower release from tissues

Vapor Pressure: Practice Problems (Examples)

  • Enflurane added to a beaker of oxygen: determine % enflurane and % oxygen from VP values and total pressure
  • Desflurane added to a beaker of oxygen: determine % desflurane and % oxygen similarly

Clinical Applications: Vaporizer Misfill Scenarios

  • If halothane is placed in an enflurane or sevoflurane vaporizer, delivered concentration will be higher than the dial setting because halothane has a higher VP
    • VP(enflurane) ≈ VP(sevoflurane) around 170–172 mmHg; VP(halothane) ≈ 244 mmHg
  • If halothane is placed in an isoflurane vaporizer, delivery will be about the same as the dial setting (VP similar: 244 vs 240)
  • If enflurane or sevoflurane are placed in a halothane or isoflurane vaporizer, delivered concentration will be less than the dial setting (lower VP than the other agents)
  • If isoflurane is placed in an enflurane or sevoflurane vaporizer, delivered concentration will be higher than the dial setting (isoflurane VP is higher than enflurane/sevoflurane)
  • If isoflurane is placed in a halothane vaporizer, delivered concentration will be about the same (VPs are similar: 244 vs 240)
  • Practical takeaway: reason through VP differences to anticipate delivered concentrations
  • Mnemonics:
    • HLH: if a higher VP agent is placed in a vaporizer for a lower VP agent, delivered concentration will be higher than the dial setting
    • LHL: if a lower VP agent is placed in a vaporizer for a higher VP agent, delivered concentration will be lower than the dial setting

Solubility

  • Definition: ability of a solute (solid, liquid, gas) to dissolve in a solvent at a given temperature
  • Key factors:
    • Temperature: affects solids and gases
    • Pressure: affects only gases
    • Other influences: solute concentration, concentration gradient, molecular size (MW)
  • Examples:
    • Solutes: salt, sugar, alcohol, anesthetic gases (isoflurane, sevoflurane)
    • Solvents: water, oil, blood
  • Clinical note: Solubility varies widely among substances and directly impacts drug/anesthetic behavior

Solubility in Gases and Liquids

  • Solubility of a gas (anesthetic gas) in a liquid (blood) affects induction and emergence
  • Process: anesthetic vapor → alveoli → dissolve in blood → carried to brain for mechanism of action

Diffusion

  • Movement of molecules from high to low concentration across a membrane
  • Requirements: concentration gradient must exist; no gradient → no diffusion
  • Outcome: molecules distribute evenly until equilibrium

Diffusion: Graham’s Law

  • Rate of effusion of a gas through an orifice is inversely related to molecular weight
  • Smaller molecules move faster
  • Limits to the law exist beyond an orifice

Diffusion: Five Key Factors

  • Directly related (promote diffusion):
    • Concentration gradient
    • Tissue area
    • Fluid/tissue solubility
  • Indirectly related (limit diffusion):
    • Membrane thickness
    • Molecular weight (MW)

Osmosis

  • Movement of water across a semi-permeable membrane
  • Features: membrane allows water and small molecules to pass; large molecules are trapped
  • Result: water shifts to balance concentration difference; large molecules prevent complete equilibrium

Fick’s Law of Diffusion

  • Directly proportional (factors that increase diffusion):
    • Pressure gradient
    • Solubility
    • Membrane area
  • Inversely proportional (factors that hinder diffusion):
    • Membrane thickness
    • Molecular weight
  • Clinical relevance: explains gas exchange in normal breathing; key to induction and emergence from anesthesia
  • Standard mathematical form (conceptual): J ext{ (flux)} \propto \frac{D A \Delta C}{\Delta x}
    • Where flux increases with diffusion coefficient (solubility) and area, and with the concentration gradient; decreases with membrane thickness and distance

Henry’s Law

  • Definition: the amount of gas dissolved in a liquid is proportional to the partial pressure of that gas above the liquid
  • Clinical relevance: essential for understanding gas transport and anesthesia gas exchange
  • Key constant and expression:
    • C = k_H P or more generally, solubility is proportional to partial pressure
  • In the body (O2 and CO2):
    • Oxygen dissolution constant: 0.003\, \frac{mL}{100\,mL\, blood\, mmHg}
    • Carbon dioxide dissolution constant: 0.067\, \frac{mL}{100\,mL\, blood\, mmHg}

Henry’s Law: Numerical Examples

  • Oxygen dissolved in blood at PaO2 = 300 mmHg:
    • Dissolved O2 = 300 \times 0.003 = 0.9\,\text{mL}/100\,\text{mL}\,\text{blood}
  • Change from PaO2 100 to 500 mmHg:
    • Dissolved O2 at 100 mmHg = 100 \times 0.003 = 0.3\,\text{mL}/100\,\text{mL}
    • Dissolved O2 at 500 mmHg = 500 \times 0.003 = 1.5\,\text{mL}/100\,\text{mL}
    • Increase = 1.5 - 0.3 = 1.2\,\text{mL}/100\,\text{mL}
  • Carbon dioxide dissolved in arterial blood at PaCO2 = 50 mmHg:
    • Dissolved CO2 = 50 \times 0.067 = 3.35\,\text{mL}/100\,\text{mL}
  • Estimating PaO2 from inspired oxygen (FiO2) and flow rate (example): 5 L/min with FiO2 = 0.40
    • Inspired O2 partial pressure approximation: 200\text{ mmHg}
    • Dissolved O2 = 200 \times 0.003 = 0.6\,\text{mL}/100\,\text{mL}

Blood/Gas Solubility Coefficients (Ostwald Coefficients)

  • Definition: ratio of anesthetic in blood (liquid phase) to gas (gaseous phase) at equilibrium
  • Interpretation:
    • High blood solubility → more agent stays in blood, less in gas
    • Low blood solubility → more agent remains in gas phase
  • Numerical values (B/G Coefficients):
    • Halothane: 2.50
    • Isoflurane: 1.4
    • Desflurane: 0.42
    • Sevoflurane: 0.60
    • Nitrous oxide: 0.47
  • Interpretations and implications:
    • Based on Ostwald partition coefficients, which agent is most blood soluble? Halothane (2.50).
    • Which is least blood soluble? Desflurane (0.42).
    • Which agent will get to the brain the fastest? Those with the lowest blood solubility (desflurane fastest in uptake).
    • Which will be slower to induce? Those with higher solubility (halothane slowest).
    • Hypothermia effect: cooler temperatures generally increase gas solubility in blood, potentially slowing onset; thus, hypothermic patients may go to sleep slower with these agents.
  • Practical takeaway: Inhalation agents with low blood solubility leave the blood quickly and enter tissues, producing rapid anesthetic states; higher blood/gas solubility slows brain uptake