Week 8: Solutions, Concentrations, Body Fluids & Properties of Gases and Respiration

Characteristics of Solutions

• A solution is a homogeneous mixture where each substance retains its chemical identity.
• It contains two components:
  • Solute
  • Solvent

Solvents and Solutes

• Solvent: Present in the greatest amount.
• Solute: Present in a smaller amount relative to the solvent.

Properties of Solutions

• Contain a solvent and one or more solutes.
• Have variable composition (ratio of solute to solvent can vary).
• Dissolved solutes are present as individual substances (molecules, atoms, or ions).
• Solutes remain uniformly distributed.

Examples of Solutions

• Sports Drinks:
  • Solvent: Water
  • Solutes: Sugars, salts (ions), and vitamins
  • Isotonic drinks
• Isotonic Saline Drip Bags:
  • Solvent: Water
  • Solutes: NaCl and glucose

Types of Gaseous Solutions

• Gas dissolved in gas: Dry air (oxygen and other gases in nitrogen).
• Liquid dissolved in gas: Wet air (water vapor in air).
• Solid dissolved in gas: Moth repellent sublimed in air.

Liquid Solutions

• Gas dissolved in liquid: Carbonated beverage (carbon dioxide in water).
• Liquid dissolved in liquid: Cordial in water.
• Solid dissolved in liquid: Salt dissolved in water.

Solid Solutions

• Gas dissolved in solid: Hydrogen in platinum.
• Liquid dissolved in solid: Dental filling (mercury in silver).
• Solid dissolved in solid: Sterling silver (copper in silver).

Formation of Solutions

• Requires attraction between solute and solvent for uniform dispersion.
• Without attraction, solute particles stay together, not mixing with the solvent.

Water Solutions

• Water is the universal solvent because it is polar.
• Forms solutions with ionic and/or polar solutes.
• Ionic Compounds: Water molecules surround and hydrate ions (e.g., NaCl).

Covalent Compounds (Molecules) in Water

• Polar molecules (e.g., sucrose, ethanol, methanol) dissolve well in water.
• Nonpolar molecules (e.g., fat, oil, grease, iodine) do not dissolve in water.

Polarity and Solutions

• Polar solvents will not form a solution with a nonpolar solute.

Solubility

• "Like Dissolves Like": Polar molecules dissolve in polar solvents, nonpolar in nonpolar.
  • Sucrose in water (polar).
  • Grease in kerosene (nonpolar).
• Solubility: Maximum amount of solute that dissolves in a given amount of solvent, usually in grams per 100g of solvent, and is temperature-dependent.

Saturated Solutions

• Unsaturated solution: Less solute than the maximum amount is dissolved.
• Saturated solution: Contains the maximum amount of solute that can be dissolved under the given conditions.

Saturated Solutions in the Body

• Gout:
  • Uric acid concentration exceeds its solubility (7 mg/100mL) in blood plasma at 37°C.
  • Uric acid crystals form in cartilage, tendons, and soft tissues, causing painful attacks.
• Kidney Stones:
  • Excessive mineral ingestion and insufficient water intake cause mineral salts (e.g., calcium phosphate, calcium oxalate) to exceed their solubility.
  • Kidney stones cause pain and discomfort when passing through the urinary tract.

Concentration of Solutions

• Specifying solution composition involves specifying solute concentrations.
• Concentration: Amount of solute in a specified amount of solution.
• Concentrated solution: Contains a large amount of solute.
• Dilute solution: Contains a small amount of solute.

Percent Concentration

• Amount of solute expressed as a percentage.
• Example: 5% dextrose solution = 5g dextrose/100mL solution.
• Formula: \text{Percent conc. (%)} = \frac{\text{Mass of solute (g)}}{\text{Volume of solution (mL)}} \times 100

Percent Concentration Problem

• Problem 1: What is the percent concentration of a Gatorade sports drink that contains 4.5g glucose in 250 mL?
  • Percent conc. (%) = $\frac{4.5g}{250 mL} \times 100 = 1.8\%$

Percent Concentration Problem 2

• Problem 2: How many grams of salt (NaCl) are in 1000 mL of a 0.9% intravenous saline (salt) solution?
• $0.9\% = \frac{x}{1000 mL} \times 100$
• $x = \frac{0.9 \times 1000}{100} = 9.0 \text{grams of NaCl}$

Ratio Concentration

• Expressed as the ratio of solute to solution (e.g., 1:20).
• Example: A 1:1000 saline solution means 1g of saline in 1000 mL of water.
• Problem 3: Chlorhexidine is used as a 1:2000 solution for a general antiseptic; how many grams are required to make up 500 mL of this chlorhexidine solution?

Ratio Concentration

• Ratio concentration (strength) = 1:2000 ; We want to make up a solution of 500 mL how much do we need
• Need: xg in 500 mL
• Given strength: 1g in 2000 mL
• Calculation:
  • $\frac{x g}{500 mL} = \frac{1 g}{2000 mL}$
  • $x g = \frac{500 mL \times 1g}{2000 mL}$
  • $x g = 0.25 g$

Molar Concentration

• States the number of moles of solute in 1 litre of solution.
• A 1.0 molar solution (1.0 M) has 1.0 mole of solute in 1.0 litre of solution.

Dilutions

• Diluting a solution of known concentration (stock solution) to a lower concentration.
• Dilution: Adding more solvent to a stock solution to lower its concentration.
• The amount of solute remains constant.
• Volume increases, causing the concentration to decrease.

Dilution Formula

• $C<em>1 \times V</em>1 = C<em>2 \times V</em>2$
  • $C_1$ = concentration of stock solution
  • $V_1$ = volume of stock solution
  • $C_2$ = concentration of diluted solution
  • $V_2$ = volume of diluted solution

Dilution Formula - Important

• The concentration units for $C<em>1$ and $C</em>2$ must be the same.
• The volume units for $V<em>1$ and $V</em>2$ must be the same.

Dilution Calculations

• Problem 4: What is the final concentration if 80.0 mL of a 2.0% stock glucose solution is diluted to 400.0 mL?
  • $2.0\% \times 80.0 mL = C_2 \times 400.0 mL$
  • $C_2 = \frac{2.0\% \times 80.0 mL}{400.0 mL}$
  • $C_2 = 0.4\%$

Fluid Compartments of the Body

• Intracellular fluid volume = 25 L, 40% body weight
• Total body water volume = 40 L, 60% body weight
• Extracellular fluid volume = 15 L, 20% body weight
• Interstitial fluid volume 12 L, 80% of ECF
• Plasma volume = 3 L, 20% of ECF
• Fluid compartments of the body, are complex solutions transporting many solutes around the body, examples include;
  • Gases; e.g. $O<em>2$, $CO</em>2$
  • Ions; e.g. $Na^+$, $Cl^-$, $K^+$, $Ca^{2+}$, $HCO_3^-$
  • Nutrients; e.g. glucose
  • Proteins; e.g. hormones insulin
  • Wastes, etc.

Gases and the Respiratory System

• The major function of the respiratory system is to supply the body with oxygen ($O<em>2$) and to dispose of carbon dioxide ($CO</em>2$). As cells use oxygen, they produce carbon dioxide as a waste product.

Properties of Gases

• Gas particles are far apart.
• A gas has no definite shape or volume, filling any container.
• A gas is less dense and can be compressed.

Kinetic Theory of Gases

• A gas is composed of very small particles (molecules and atoms).
• The particles of a gas are very far apart.
• Gas particles move rapidly, colliding with each other and the container walls.
• Gas particles do not attract or repel one another.
• The kinetic energy of gas particles is related to the temperature; motion increases with temperature.

Properties of Gases

• Important properties related to gases:
  • Pressure (P)
  • Volume (V)
  • Temperature (T)

Pressure

• Pressure of a gas is the force created when gas particles hit the container wall.
• Typical pressure units:
  • Pascal (Pa) - the SI unit
  • Atmosphere (atm)
  • Millimeters of mercury (mmHg)

Volume

• The volume of a gas is equal to the volume of the container.
• Typical volume units:
  • Litre (L) - the metric unit
  • Millilitre (mL)

Temperature

• All calculations with gases use the Kelvin temperature scale.
• $K = °C + 273$
• Absolute zero (0 K) means particles would have no energy or motion.

Atmospheric Pressure

• The air covering the Earth's surface contains many gas molecules, exerting atmospheric pressure due to their mass and gravity.
• Atmospheric pressure can be measured using a barometer. At 1 atm, mercury in the glass tube would be 760 mm high.
• Standard atmospheric pressure is 1 atm or 760 mmHg.

Atmospheric Pressure Variation

• Atmospheric pressure varies at different heights and depths.
  • Sea level: 1 atm
  • 10 m below: 2 atm
  • 20 m below: 3 atm
  • 30 m below: 4 atm
  • 5 km high: 0.5 atm
  • Mount Everest (8.9 km high): 0.35 atm

Boyle’s Law

• Volume (V) of a gas changes inversely with the pressure (P), as long as temperature and amount of gas are constant.
• $P<em>1 \times V</em>1 = P<em>2 \times V</em>2$
  • Where:
    • P is the gas pressure
    • V is the gas volume
    • Subscript 1 is initial conditions
    • Subscript 2 is final conditions

Boyle's Law Example 1

• A balloon contains 500 mL of gas at 760 mmHg. It rises to an altitude where its pressure becomes 380 mmHg. What is its new volume? (Assume temperature is constant).
• $P<em>1 \times V</em>1 = P<em>2 \times V</em>2$
• $(760 mmHg \times 500 mL) = (380 mmHg \times V_2)$
• $V_2=\frac{(760 \times 500)}{380}$
• $V_2 = 1000 mL$
• (the volume has doubled)

The Mechanics of Breathing and Boyle’s Law

• During inspiration, lung volume increases, decreasing air pressure in the lungs. Air flows in from outside (high pressure) to equalize pressure.
• During expiration, lung volume decreases, increasing lung pressure. Air flows out to equalize pressure.