Chem Exam 2 Flashcards
Solution
Homogenous mixture that consists of one or more solutes uniformly dispersed at molecular or ionic level throughout a medium known as the solvent
- Not always liquid
Ex: Air: solution of nitrogen, oxygen, other gases
Former dental fillings: silver amalgams (mix of silver and mercury)
Homogenous mixture
Not possible to discern boundaries between components of the mixture. Ex: Normal saline (uniform throughout sample)
Phase Boundary
Separates regions of a mixture where the chemical or physical properties of the mixture change
Solute
Material that got dissolved; component kf solution present in smaller quantity. Ex: Sodium chloride
Solvent
Material that does dissolving, usually water
Molarity (molar concentration)
Moles of solute per liter of solution
M = mol of solute/L of solution (mol/L)
- Molar concentrations are conversion factors btw moles of material and liters of solution
- Depends on temperature of solution
Molar concentration will decrease at temperature increases
Molality (molal concentration)
Express concentration in terms of moles of solute per kilogram of solvent
m=mol of solute/kg of solute
- Used in physical chemistry -> quantities of solute and solvent are considered separately
- Mass is not temperature dependent, molality is not temperature dependent
- Less convenient in analysis bc quantities of a solution measured out by volume or mass in lab include both solute and solvent
- When doing stoichiometry, molality requires additional calculation to take into
Molarity vs Molality
Molality and molarity are never equal, but difference is is smaller as solutions become more dilute
Convert btw molarity and molality, need to know density of solution
molality of solution = x mol of solute/y kg of solvent
Percent by weight to volume
%w/v
The percent of concentration you encounter in a clinical setting when measuring out a volume of medicine in syringe
Grams of solute per 100mL of solution
%w/v = g of solute/100mL of solution
OR
%w/v= g of solute/mL of solution (100%)
First equation is useful as a conversion factor btw grams of solute and milliliters of solution
Second equation is useful tp calculate concentration of a solution
- To relate to percent by weight to percent weight to volume, use the density of solution
Percent by volume (%v/v): never used in analytical lab bc values are not additive
Equivalents (Eq)
Analogous to a mole
Normality
Analogous to molarity
Normality and equivalents
1 equivalent of a substance contains 1 mole of chemical reactivity
N= equivalents of solute/L of solution
Normality is equal to equivalents of solute per liter of solution
Unless context of chemistry is specified, normality is ambiguous
Parts per million
Concentration of extremely dilute solutions is sometimes expressed as part per million
A ppm concentration is analogous to a percent concentration, except you are comparing amount of solute to a million grams of solute instead of 100 gram.
ppm = g of solute/1x10^6g of solution =
g of solute/g of solution (1x10^6)ppm
Solubility
Some solutes are more soluble in given solute than others
The solubility of solute is amount of solute that will dissolve in a given amount of solvent at a given temp
Saturated solution contains maximum amount of a solute, as defined by its solubility. No more solute will dissolve in a solution saturated with solute. if solution is not saturated, more solute will dissolve in solution.
Supersaturated solution
Supersaturated solution contains more solute than allowed by solubility of solute.
- NOT a stable system, bc there is more solute dissolved in sample than solvent can accommodate
- Excess solute will come out of solution, crystallizing as a solid, separating as a liquid, or bubbling out as a gas
Ex: when blood or urine in kidneys become supersaturated with calcium oxalate or calcium phos, kidney stone can form - If solute is a gas in liquid solvent, will see bubbles forming in solution (Fizz when opening bottle of beer or soda)
Miscibility
Two liquids are miscible if they are soluble I each other in all proportions
Ex: Alcohol and water are miscible with each other.
Oil and water are immiscible
Like dissolves like
- Polar solutes are more soluble in polar solvents
- Nonpolar solutes are more soluble in non polar solvents; insoluble in water
- Most organic molecules are nonpolar
- Most ionic compounds are polar -> dissolve in water but not soluble in organic compounds (ether, hexane)
Energy Changes
When a solute dissolves in a solvent, there is an associated energy change and temperature change to solution
The energy change when using heat and cold packs is called heat of solution or the enthalpy of solution:
Δhsoln
Defined as the energy change that accompanies dissolving exactly 1 mole of solute in a given solvent
Enthalpy H is equal to the heat Q as long as the pressure remains constant
The energy change may be endothermic or exothermic
Whether the heat of solution is endothermic or exothermic depends on the relative magnitudes of the lattice energy and the heat of solvation
* If tearing the ions apart requires more energy than is released by solvation, then “Δ”Hsoln is going to be positive (endothermic)
- If the energy released by solvation is greater than the energy required to tear the ions apart, “Δ”Hsoln is going to be exothermic
Exothermic
- If the solution process is exothermic, energy flows out of the system (solvent and solute) into the surroundings, resulting in a temperature increase in the solution
- Energy is lost
Endothermic
- If the solution process is endothermic, energy flows from the surroundings into the system, resulting in a temperature decrease in the solution
- Energy is gained
Effect of Pressure on Solubility
As pressure increases, the solubility of a gaseous solute in a liquid solvent increases
* With less pressure, solubility of gaseous solute decreases (CO2 escaping champagne bottle)
- Since solids and liquids are not very compressible, at least not compared to gases, pressure has very little effect on the solubility of solid and liquid solutes
The quantitative relationship between pressure and solubility is given by Henry’s law:
S = k(H)P(gas)
S = solubility kH = Henry’s law constant (0.042g/L/atm Pgas = partial pressure of the gas
Increased partial pressure of gas -> more gaseous molecules zipping around at near surface of the liquid
Effect of Temperature on Solubility
The solubility of solid and liquid solutes in liquid solvents generally increases with increasing temperature
- As temperature increases, vapor pressure of gaseous solutes increases to point that they escape the solvent into gas phase
Colligative Properties
Vapor Pressure
The vapor pressure of a solution Decreases with increasing solute concentration
The vapor pressure of a liquid results from the most energetic molecules near the surface of the liquid escaping into the gas phase
The most likely escape sites for the liquid molecules are at or near the surface of the liquid
As we begin to introduce solute molecules, some of these escape sites are occupied by the solute molecules, so fewer solvent molecules can escape into the gas phase
Therefore, the vapor pressure of the solution is less than the vapor pressure of the pure solvent
Raoult’s law states the vapor pressure of a volatile component of a solution (P) is equal to the vapor pressure of the pure substance (Po) times the mole fraction (ᵡ) of that substance
P = χPo
Colligative Properties
Boiling Point
The boiling point of a solution INCREASES with increasing solute concentration
Temperature at which the vapor pressure of the material is equal to the ambient pressure
The boiling point of a solution increases as the concentration of solute(s) increases
The change in boiling point is directly proportional to the molal concentration of the solute particles
ΔTbp = Tbp,solution – Tbp,solvent = kbp · mtotal
ΔTbp = the number of degrees by which the boiling point increases Tbp,solution = the boiling point of the solution, Tbp,solvent = the boiling point of the pure solvent kbp = a constant (called the ebullioscopic constant) that is characteristic of the solvent mtotal = the molal concentration of all solute particles
Activity is the effective concentration of a solute
Activity is always less than molality
Colligative Properties
Freezing Point
The freezing point of a solution DECREASES with increasing solute concentration
Temperature at which the liquid phase of the material is in equilibrium with the solid phase
In order to enter into the solid state, the molecules (or ions or atoms) of the sample need to settle into an orderly, crystalline lattice structure
The presence of solute particles interferes with this process by getting in the way
It is necessary to cool the sample to lower temperatures, thereby lowering the kinetic energy of the molecules even further, before they will settle into the solid phase
The relationship that quantifies the degrees of freezing point depression has an identical form to boiling point elevation
ΔTfp = Tfp,solution − Tfp,solvent = kfp · mtotal
ΔTfp = the number of degrees by which the freezing point decreases Tfp,solution = the freezing point of the solution, Tfp,solvent = the freezing point of the pure solvent kfp = a constant (called the cryoscopic constant) that is characteristic of the solvent mtotal = the molal concentration of all solute particles
Colligative Properties
Osmotic Pressure
The osmotic pressure of a solution INCREASES with increasing solute concentration
Osmosis is diffusion of water through a semipermeable membrane
The relative concentration of solutes in osmotic systems is called the tonicity
Two solutions are isotonic if they contain equal concentrations of particles
Hypertonic: higher concentration of solute
Hypotonic: lower concentration of solute
Osmosis always spontaneously occurs from the hypotonic solution to the hypertonic solution
Diffusion always spontaneously occurs in the direction from an area of high concentration to an area of low concentration
Entropy demands that osmosis occur between two solutions of unequal tonicity until the concentrations of the two solutions are equal
Osmotic pressure (symbolized as capital pi, Π) results from the potential drive for the concentration of water to equalize
Osmotic pressure is a colligative property, and the osmotic pressure of a solution increases with increasing solute concentration
The relationship between osmotic pressure and concentration is given by:
ΠV = nRT
If the volume is in liters and we divide both sides of the relationship by volume, we get:
Π = MRT
where Π is the osmotic pressure, M is the molarity of the solute particles, R is the ideal gas constant, and T is the absolute temperature
Colloids
Similar to solutions in that they consist of one phase uniformly dispersed in a second phase
Examples: milk, blood, paint, and jelly
Not true solutions because the particles in the dispersed phase are not the size of molecules or ions
Particles in a colloid range in size from 10 nm to 200 nm
Colloidal particles cannot be filtered and do not settle out of solution
Colloids can be stable for years if they are stored under controlled conditions
Colloids exhibit the Tyndall effect, whereas solutions do not
Particles of a colloid are large enough to scatter light passing through
Particles in a true solution are too small to scatter visible light, so solutions do not exhibit the Tyndall effect
Chemical Equilibria
Le Châtelier’s Principle
States that equilibrium is a good thing, and nature strives to attain and/or maintain a state of equilibrium or homeostasis
Chemical Equilibria
Changing Concentration
If you add products, the equilibrium will shift toward reactants
If you remove products, the equilibrium will shift toward products
The system readjusts to counteract whatever change made
Chemical Equilibria
Changing Temperature
Increasing temperature favors endothermic processes
Chemical Equilibria
Changing Volume and Pressure
Significantly impacts equilibrium reactions only when at least one of the reactants or products is a gas bc liquids and solids are not compressible
Decreasing volume increases pressure -> smaller number of gas particles
The Equilibrium Constant
A system is in a state of equilibrium when there is a balance between reactants and products
This balance is defined by thermodynamic parameters, namely bond strengths and the intermolecular forces between all the molecules in the system
The equilibrium constant (K) is the numerical description of that balance
K is equal to the product of all of the molar concentrations of the products, each raised to the power of their stoichiometric coefficients, divided by the product of the molar concentrations of the reactants, each raised to the power of their stoichiometric coefficients
Subscripts
The equilibrium constant (K) is often appended with a subscript that denotes the type of equilibrium reaction
Keq is a generic equilibrium constant
Ka is an equilibrium constant governing the ionization of weak acids
Kb is an equilibrium constant governing the ionization of weak bases
Ksp is an equilibrium constant governing the solubility of sparingly soluble compounds
Meaning of K
As K increases, the reaction tends to increasingly favor products and the forward reaction becomes more favorable
As K decreases, the reaction tends to increasingly favor starting materials, and the reverse reaction becomes more favorable
Solids and Liquids
Pure solids or liquids comprise a different phase from where reactions in aqueous media occur
The concentration of pure solids and liquids is a constant value (or very nearly so), and these constant values for concentration are included with the equilibrium constant value
The concentration of solids, liquids, and water (as a solvent) do not appear explicitly in the equilibrium constant expression
Reversing a Reaction
When you reverse the equation for a chemical reaction, Kforward is the reciprocal of Kreverse
Kforward = 1/Kreverse
Definition of Acids and Bases
Arrhenius definition (most operational definition)
Acid–species that increases the hydronium ion (H3O+) concentration in an aqueous solution
Base–species that increases the hydroxide ion (OH–) concentration in an aqueous solution
Brønsted2 definition (most generally useful definition)
Acid–species that donates a hydrogen ion (H+) to a base
Base–species that accepts a hydrogen ion (H+) from an acid
Conjugate Acid-Base Pairs
The charge on the conjugate acid is always one greater than the charge on its conjugate base
When an acid gives away its hydrogen ion to a base, the acid is converted into its conjugate base
When a base accepts a proton from an acid, the base is converted into its conjugate acid
In generic form, we can express this process as:
HA + B → A− + BH+
Conjugate Base Ex: HNO3 -> NO3- (conjugate base)
Conjugate Acid Ex: H2O -> H3O+ (conjugate acid)
Amphiprotic Species
Can behave as either an acid or a base
Ex: H2O
Base w/ HCl -> H3O+ (becomes conjugate acid)
Acid w/ HCO3- -> OH- (becomes conjugate base)
Strong Acids
Acids donate protons
Very determined to foist their proton off onto a base
Essentially 100% ionized when dissolved in water, but this is not an equilibrating process, so all of the starting materials are converted into products
Strong acids are relatively rare, most acids are weak and only partially ionize in water
Ex: HCl, H2SO4, HI, HBr, HNO3
Strong Bases
Bases accept hydrogen ions
In water, the strongest possible base is the hydroxide ion, OH−
A strong base ionizes essentially 100% to produce the OH − ion, so a strong base is a soluble ionic hydroxide
Ex: LIOH, NAOH, KOH, Ba(OH)2
Weak Acids
Are able to donate hydrogen ions to bases, but are less determined to do so than strong acids
When a weak acid dissolves in water, it establishes a dynamic equilibrium between the molecular form of the acid and the ionized form
Ex: HC2H3O2, H2CO3, H3PO4, NH4+
Weak Bases
Are able to accept hydrogen ions from acids, but are less determined to do so than strong bases
Do not completely ionize in water to produce an equivalent concentration of the hydroxide ion, because when a weak base dissolves in water, it establishes a dynamic equilibrium between the molecular form and the ionized form
Ex: NH3, HCO3-, CO3^2-, HPO4^2
Polyprotic Acids
A diprotic acid has two hydrogen ions to donate, so a diprotic acid can behave as an acid twice
A triprotic acid has three hydrogen ions to donate
The number of acidic protons is not necessarily the number of hydrogens in the molecular formula
The reason for the dichotomy of hydrogen lies in the molecular structure of acetic acid
The acidic hydrogen is bonded to a highly electronegative oxygen atom
The O—H bond is polarized toward the oxygen to a point that a base can snatch away the hydrogen as an H+ ion from the acetic acid molecule
The other three hydrogens are bonded to a carbon atom, and carbon and hydrogen have almost identical electronegativities, so those bonds are nearly nonpolar and the hydrogens have no tendencies to be removed as H+ ions
Acid and Base Strength: Ka and Kb
Some acids are stronger than others
A stronger acid is more determined to give its proton to some base
A stronger base is more determined to take a proton from some acid
When a strong acid dissolves in water, it dissociates completely
All other acids are weak acids
The relative strength of a weak acid is quantified by the equilibrium constant Ka governing the ionization of the weak acid
A larger value of Ka means a stronger acid
Acid/Base Strength of Conjugate Acid–Base Pairs
The stronger the acid, the weaker its conjugate base
The stronger the base, the weaker its conjugate acid
General guidelines
The conjugate base of a really strong acid has no base strength
Ex: HCl is a strong acid; Cl- phase no base strength
The conjugate base of a weak acid has base strength
Ex: Acetic acid is a weak acid; acetate acid has base strength
The conjugate acid of a weak base has acid strength
Ex: Ammonia is a weak base; ammonium ion behaves like an acid in water
Acid-base reactions
Involve a transfer of a hydrogen ion from the acid to the base
In order to predict the products of an acid–base reaction:
Identify which is the acid and which is the base
Move an H+ ion from the acid to the base, converting the acid into its conjugate base and the base into its conjugate acid
Any acid–base reaction has two acids and two bases:
One acid and one base on the reactant side
Conjugate acid and conjugate base on the product side
The base almost always has a lower (more negative) charge than the acid
Also, hydrogen is almost always the first atom listed in the formula of an acid
The reaction equilibrium always favors the formation of the weaker acid
Look up the Ka’s for the reacting acid and the conjugate acid of the reacting base
Measuring acidity
The p-Function
The p in pH is a mathematical operator that means the negative logarithm of, and the H in pH means hydrogen ion concentration, so the definition of pH is:
pH = −log[H+]
p-function operator used to express the concentrations of many ions, but the meaning of the p is the same in every case—take the logarithm of the concentration and then change the algebraic sign
A logarithm function is a way to map a vast range of values onto a much smaller set of values
To apply the p-function, take the logarithm of the [H+] concentration and then change the sign
The value of the pH will be the same as the absolute value of the exponent when the first part of the scientific notation is exactly 1
pH values have no intrinsic units—logarithms represent “pure numbers”
Each change of 1 pH unit means the hydrogen ion concentration is changing by a factor of 10, so small changes in pH correspond to much larger changes in acidity level
Removing the p-function is accomplished by changing the sign of pH and then taking the antilog
Self-Ionization of Water
Water has some very weak acid–base properties, and therefore sets some limits on the parameters of the pH scale
A tiny fraction of water molecules dissociates or ionizes into a hydrogen ion and a hydroxide ion
H2O ⇄ H+ + OH−
The equilibrium constant that governs this equilibrium is called Kw
At 25°C, Kw = 1.0 × 10−14 so pKw is 14.00
Since water is a pure liquid, the concentration of water does not appear in the equilibrium constant
Kw = [H+] [OH−] = 1.0 × 10 −14 at 25°C
In pure water, the concentrations of the H+ and OH− ions are equal
[H+] = [OH−] in pure water
The pH of pure water is 7.00 because that is the negative logarithm of the hydrogen ion concentration that is generated by the self-ionization of water
Relationship Between pH and pOH
Because pH and pOH are derived from the ionization of water, there is a fixed relationship between them
The pH plus the pOH of any aqueous solution (at 25°C) always adds up to 14.00:
14.00 = pH + pOH
pKa and pKb
The equilibriums constants for acids and bases are commonly presented as p-functions
The pKa and pKb of a conjugate acid–base pair sum to give 14.00:
pKa + pKb = 14.00
The Key Relationships
To solve pH calculation problems, you need to be conversant with six key relationships:
pH = −log [H+]
pOH = −log [OH−]
[H+][OH−] = Kw = 1.00 × 10−14
pH + pOH = pKw = 14.00
pKa + pKb = 14.00
KaKb = Kw
Calculating the pH of a Strong Acid Solution
Strong acids completely ionize in water, which means every mole of strong acid falls apart into an equal number of moles of hydrogen atoms
[H+] is equal to the formal concentration of the acid
Calculating the pH of a Strong Base Solution
Strong bases completely ionize in water, which means every mole of strong base falls apart into an equal number of moles of hydroxide ions, except for Ba(OH)2, which gives 2 moles of hydroxide ions per mole
[OH−] is equal to the formal concentration of the base
Be careful when calculating the pH of a basic solution because the first number that comes out of your calculator is going to be pOH, not pH
pOH= -logOH pH = 14 - 0.82 = pOH
Calculating the pH of a Weak Acid Solution
Weak acids and bases do not ionize completely, so we must use the equilibrium constant expression to find the pH of weak acid solutions
Calculating the pH of a Weak Base Solution
Involves exactly the same process as calculating the pH of a weak acid
pH of Salt Solutions
Acids and bases react to give salts and (usually) water
The pH of a salt solution depends on the acid/base strength of the acid or base from which it was derived
Salts of Strong Acids/Bases
Very strong acids have conjugate bases with no base strength, so these ions do not change the pH of a solution
Salts of Weak Acids
Salts of weak acids are bases, and most acids are weak acids, so the conjugate bases of these acids have some appreciable base strength
The vast majority of anions increase the pH of a solution
Salts of Weak Bases
Salts of weak bases are acids
Many medications are organic amines, or alkaloids, which are weakly basic organic compounds that are generally not very soluble in water
They are converted to their conjugate acid form by reacting with a strong acid (such as HCl)
Extreme Concentrations
Be wary when concentrations get below 10^−3 molar, or when the concentrations aren’t that different from the Ka values
Example: What is the pH of 1.0 × 10^−8 M HCL
If you use your calculator to determine the negative log of 1.0 × 10−8, you get an answer pH = 8.0, which is incorrect
You cannot be totally complacent about these kinds of calculations, especially when the concentrations become very dilute
Other acidic species
Nonmetal oxides dissolve in water to give acid solutions
The most physiologically important example is carbon dioxide
Buildup of carbon dioxide in the blood results in acidosis
In cellular tissue, where the carbon dioxide concentration is relatively high, the increased acidity slightly alters the structure of hemoglobin and facilitates the release of oxygen
Carbon dioxide is a nonmetal oxide because it is a compound composed of a nonmetal and oxygen
Nonmetal oxides are sometimes called acid anhydrides because they are produced by stripping water from an acid
When carbon dioxide dissolves in water, it combines with a water molecule to give carbonic acid
CO2 + H2O ⇄ H2CO3
When the carbonic acid forms, it dissociates according to its acid strength
H2CO3 ⇄ H+ HCO−3
So, when CO2 dissolves in water, the pH drops
Buffers
A pH buffer is a solution that resists changes in pH
It contains a weak acid (HA) and its conjugate base (A−) or a weak base and its conjugate acid
Buffer solutions resist change in pH
If a strong base is added to a buffered solution, the weak acid in the buffer HA reacts with the hydroxide ion to give water and the weak base A−
HA + OH− ⇄ H2O + A−
This results in converting a strong base OH− into a weak base A−
The pH increases, but not by much
If a strong acid is added to a buffered solution, the weak base in the buffer (A− ) reacts with the H+ ion to give HA
A− + H+ ⇄ HA
This results in converting a strong acid H+ into a weak acid HA
The pH decreases, but not by much
Calculating the pH of a Buffer
The pH of a buffer is determined by the acid strength of HA
As the acid strength of HA increases, the pH range maintained by a buffer system based on HA decreases
Equation for the acid ionization of HA:
HA ⇄ H+ + A−
The equilibrium constant expression for this equilibrium is:
Ka = [H+] [A-] / [HA]
The equilibrium state can be achieved from an infinite number of starting points
The Henderson–Hasselbach equation, or buffer equation is:
pH = pKa + log [A-] / [HA]
This equation can be used to calculate the pH of a buffer or to determine the ratio of weak acid to conjugate base at a given pH
Dilution of Buffer Solutions
The pH depends on the ratio of weak base to its conjugate acid, not the numerical value of each, so a pH buffer can be diluted with water without changing the pH of the solution
Buffer Capacity
To be an effective buffer, the pH must be within one pH unit of the pKa of the weak acid
pHeffective = pKa ± 1
To prepare an effective buffer, we need to identify an acid with a pKa as close as possible to the desired pH
Alpha Plots
An alpha plot for a buffer (or an amino acid) shows the percentage of each component in a buffer system as a function of pH
Alpha for a given species is the percent of all the material that is present in that form
When alpha for the acid form is 75%, the ratio of acid to conjugate base is 75:25
The shape of the curves in an alpha plot is determined by the buffer equation
At low pH, most of the material is present as its acid
As the pH increases, the acid form converts to the conjugate base form
At high pH, most of the material is present in the conjugate base form
Partial Pressure
Total Pressure = 1 atm = 760mmHg
50% of gas present -> 1/2 760mmHg
If proportion of molecules goes up, partial pressure goes up
As partial pressure goes up, more gas molecules enter liquid phase
Partial pressure determines likelihood of going into liquid
Factors that affect movement of gas into or out of surface layer of liquid
-solute
-solvent
-temperature (more likely to leave liquid phase if temp is high)
KH- tells likelihood of going into liquid
Concentration = Partial pressure/KH Concentration = P/KH -> Henry's law
Colligative Properties
Vapor pressure lowers
Boiling point elevation
Freezing point depression
Presence of solute interferes w/ solvent at interface btw 2 phases
Liquid has particular vapor pressure bc particles at surface of liquid can spontaneously enter gas phase
Number of gaseous particles present determines vapor pressure
Solute particles at interface occupy some of surface area, inhibiting some solvent from evaporating
Presence of solvent causes vapor pressure to decrease
When solution is raised to temp of solvent’s boiling point, solute particles block solvent molecules from going into gas phase -> increased temp
Presence of solute interferes with lattice formation -> decrease in temp
Salt added to snow, prevents ice formation because it interferes w/ lattice formation -> lower temp
Concentration
measure of how much solute present per volume of solvent
Diffusion
Movement of fluid from area of higher concentration to a lower concentration
Hypertonic
More concentrated
Water molecules move out of cell and cause cell contraction or plasmolysis
Hypotonic
Less concentrated
Water molecules move into cell and cause cell expansion or cytolysis (cell bursting)
Isotonic
Equal amounts of water molecules in and out of cell. Cell neither expands nor contracts..
Equilibrium
Bronsted-Lowry Acid
HCl
Proton donor
Conjugate base -> Cl-
Bronsted-Lowry Base
H2O
Proton acceptor
Conjugate acid -> H3O+
Lewis acid
HCl
Electron pair acceptor
Electrophile
Lewis base
H2O
Electron pair donor
Neutrophile
pH meter
consists of pH probe (electrode) and a volt measurement device, when calibrated displays pH of given liquid
Buffer system:
Ph: 7.35-7.45
Ph < 7.35=acidosis (high H ion concentration)
Ph<7.45 = alkalosis (low H ion concentrate(ion)
CO2 in blood reacts to form carbonic acid (weak acid) -> dissociates to form bicarbonate + H ion
Most CO2 must be transformed to be transported out of the body -> form of equilibrium buffer reaction system
Too many H ions -> equilibrium reaction moves to left to form carbonic acid to prevent continued formation of H ion and prevent large change in ph
Too few H ions -> equilibrium reaction moves to right to form more CO2 converted to carbonic acid and convert to H ion to prevent drastic lowering of ph
Henderson-Hasselbalch equation:
Calculate ph of buffer solution.
Buffers consist of weak acid [HA) and its conjugate base [A-].
Ph= pka + log [A-]/[HA]
Find pka of weak acid, add log of concentration of conjugate base/concentration of weak acid
Particulate diagram: number of particles are counted for each molecule to determine concentration of molecules. Equal concentration results in log 1 (0). Concentration with more acid -> ph < pka value. Concentration with more conjugate base -> ph > pka value
Ph= pka + log [A-]/[HA] pH-pka = log [A-]/[HA] 10^(ph-pka) = [A-]/[HA]
Ka value less than 1 indicates a weak acid
To find pka, take -log of ka value Log 1 = 0 Log of # < 1 = negative number Log of # >1 = positive number [HA] = [A-] -> ph = pka [HA] > [A-] -> ph < pka [HA] < [A-] -> ph > pka
Buffer:
Contains aqueous solution, contains both weak acid [HA] and conjugate base [A-]
Ka is the equilibrium constant
Acid dissociation reaction of HA (reaction is in equilibrium) H20 + HA
Fluids
A fluid is any material that has the ability to flow
Both liquids and gases are considered fluids
Basic forces that cause fluids to flow
- gravity difference
- pressure difference
When fluids are placed in a container, they assume the shape of the container
Hydrostatics
The study of fluids that are not moving
Important properties: density and pressure
Hydrostatic pressure
Ptop = Patmosphere = Ftop/A
Force at bottom = force of the top + water
Pressure at the Same Depth
Assume we have a point particle suspended in a fluid with density ρ
Since it is a point particle, it occupies no space or volume
No matter where it is placed in the fluid, the fluid will act the same, exerting the same pressure in all directions
Pressure Versus Container Shape
The pressure is independent of the container shape
Pascal’s Principle
When an external pressure is applied to a confined liquid, it is transmitted unchanged to every point within the fluid
- Plugged syringe w/ needle -> increased pressure on syringe not needle. Pressure increases everywhere within fluid by the same amount
Buoyancy
All fluids exert a buoyant force on objects immersed in them
Upward force fluid exerts on object
Archimedes principle: buoyant force = weight of displaced fluid
Archimedes’ Principle
An object immersed either totally or partially in a fluid feels a buoyant force equal to the weight of the fluid displaced
If the density of an object is greater than the density of the fluid , the object will sink
P(object>P(fluid)
If the density of the object is less than the density of the fluid , the object will float
P(object)<p> P(object) = P(fluid)
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Hydrometers
A hydrometer is a simple device used to measure the specific gravity of liquids such as urine or milk
A typical hydrometer is calibrated and has a weighted end to keep it upright in the liquid of interest
When placed in a liquid, it will sink until it displaces an amount of fluid exactly equal to its weight
If the fluid is dense, it will displace only a small amount of fluid and not sink very deep
If the fluid is not very dense, the hydrometer will sink deeper
The user reads the specific gravity of the liquid from the calibrated scale in the neck of the hydrometer
Hydrodynamics
In general, there are two types of flow
- Laminar flow:
- Smooth flow
- Characterized by an unchanging flow pattern where adjacent layers of fluid smoothly slide past each other - Turbulent flow
- Flow that is not smooth
- Chaotic and abruptly changing
Probability of turbulent flow vs laminar flow determined by Reynolds number
Flow rate
Flow rate = area * velocity
The volume of fluid passing a particular point per unit time
Will have units of volume divided by time, such as gallons per minute or liters per hour
The official SI units of flow rate are cubic meters per second (m3/s)
Speed and Diameter
As the diameter of a tube decreases, the speed of the fluid flowing through it increases
Ex: thumb over garden hose or blood flowing from a large vessel to smaller vessel
Equation of Continuity
Velocity of a fluid flowing through a pipe is inversely related to the area of the pipe
- If the flow rate is constant, fluid flows faster through a thinner pipe
The Bernoulli Equation
As the speed of a fluid increases, the pressure exerted by the fluid decreases
Must apply the Work-Energy Theorem
Wtotal = ΔKE = KEfinal − Keintial
Venturi Tube Flowmeter
Device used to measure fluid speed in a pipe
Consists of a middle section with a small diameter connected on both ends to larger diameter sections via smooth transitions in order to prevent turbulence
A U-tube containing a fluid of known density connects the large and small diameter tubes
The U-tube is a manometer, which can be used to measure differences in pressure
The Venturi tube flowmeter equation can be used to determine the speed of fluid in a pipe
Viscosity
Viscosity is a measure of a fluid’s resistance to flow
Fluids with high viscosity, such as honey, do not flow very readily
Fluids with low viscosity, such as water, flow more easily
The closer a fluid molecule is to a wall, the slower it moves
Adjacent regions of the fluid will have different speeds
*Required difference in flow to maintain flow is proportional to then length and average speed, and inversely proportional to the cross-sectional area
The faster regions will flow past the slower ones
Ways to increase flow rate
- increase pressure differential across the catheter or needle
- Using a larger diameter catheter or needle
- Using a shorter catheter or needle
Solid
intermolecular forces, Van Der Waals, keep molecules in a fixed position -> lattice formation.
Heating solid -> kinetic energy of molecules increases and range of movement increases -> volume increase and expansion
Matter is not easily compressible (density not likely to change)
Liquid
Heat added to solid -> range of movement of molecules is disrupted -> liquid formation
Matter is not easily compressible in this state (density not likely to change)
Gas
Heat added to liquid -> kinetic energy of molecules increases until intermolecular forces can’t keep molecules close to each other -> gas formation -> free movement of molecules
Matter is easily compressed
Reynolds number
R = vpdn
v= velocity p= density d= diameter n= viscosity
Transition btw laminar and turbulent flow occurs when Reynolds number is about 2000
R > 2000 = turbulent flow likely
R < 2000 = laminar flow likely
Viscosity
Units: Pa s
Pascal seconds
Friction btw fluid layers
Ability of fluid to resist flow
Factors affecting viscosity:
- strength of intermolecular forces
- greater attraction -> more viscous - size/shape of molecule
- larger molecule -> more viscous - temperature
- lower temperature -> more viscous
*Molecules moving from faster layer to slower layer increase speed of slower layer and vice versa.
Units of resistance
Pa s m^-3
Laminar flow
steady flow in one direction
occurs at lower fluid velocities
- more easily calculated
Turbulent flow
swirls in eddies
occurs at higher fluid velocities
turbulent flow decreases and laminar flow more likely w/ change from larger diameter to smaller diameter of tube
- not easily calculated
Volume flow formula
F = v(pi)d24
Volume density formula
vd = 4F(pi)d
Poiseuille equation
Describes laminar flow
If fluid has no viscosity there is no internal friction -> net force of 0
If net force is 0 -> pressure difference is 0
flow = pressure drop * (pi)r48nl
n = viscosity l= length r= radius
Small changes in diameter of a tube will affect the pressure drop noticeably
Flow through tube
Length is greater than diameter
Flow through orifice
Length is less than diameter
Equation: αd2√ ΔPρ
d= diameter ΔP= pressure difference across it p = density of fluid
- if flow reaches the speed of sound in fluid, rate of mass flow no longer depends on downstream pressure (it depends on upstream pressure only)
Bernoulli’s law
The higher a fluid’s velocity is through a pipe, the lower the pressure on the pipe’s walls
The lower a fluid’s velocity is through a pipe, the higher the pressure on the pipe’s walls
As a fluid flows through a pipe it will not gain or lose energy -> no matter where fluid is in pipe, the energy is the same at all points
Factors of energy moving fluid
- potential energy due to its pressure
- kinetic energy due to its movement
- potential energy due to force of gravity on it
If the fluid is not moving, v = 0, so the equation becomes:
P+ρgh=constant
- Equation indicates that pressure at base of column of fluid is proportional to its height and its density.
Basis for manometers
If fluid flows through tube w/ constriction in it, velocity increases as fluid passes constriction
Pressure falls as velocity increases
Venturi
A constriction w/ an entry and exit in which diameter changes gradually and maintains laminar flow
- fluid flows faster through constriction to make up for large volume flowing through the smaller area to maintain equal flow rate through pipe
- faster moving fluid = lower pressure
Bernoulli effect
Reduction in pressure as fluid passes through Venturi
When gas exits constriction into wider portion of tube, linear velocity of flow decreases and pressure increases
- The slower the fluid flows, the more pressure it puts on the pipe itself
Cohesive forces
Intermolecular forces btw molecules of a liquid
Molecules within liquid have equal amount of cohesive forces exhibited on them
Molecules on out portion of liquid have 1/2 as many cohesive forces exhibited on them bc they are on outer are of substance
Cause shapes to minimize and contract to strengthen intermolecular bonds of substance
Ex: water droplet
Adhesive forces
Interaction of liquid and solid surface. Describe ability of a liquid to adhere to surface
Ex: water droplet on plastic stays droplet bc non polar surface of plastic does not interact with polar water molecules
But spreads on glass bc polar surface of glass spreads out and polar molecules interact w/ each to maximize interactions
Ex: meniscus concave curve in glass beaker w/ water
Surface tension: energy required to increase surface area of liquid
Ex: paperclip floating on water
Capillary action: liquid flows through material bc of attractions btw liquid molecules and surface of material
Ex: dip paper in water and fluid move upward into material to maximize hydrogen binding
Mass flow rate
As water flows through pipe, it pushes along water in pipe. So mass flow rate at one part is equal to the mass flow rate at any other point in pipe
- Where the pipe is narrower, fluid will flow faster to compensate
- Fluid moving faster has less pressure than fluid moving more slowly
Ex: if one part of pipe has 1kg of fluid running through it, the narrow part of pipe also has 1kg of water running through it
Venturi mask
100% O2 from wall O2 passes through green portion from tubing quickly
High velocity O2 flow -> lower than atmospheric pressure -> air from atmosphere to flow into O2 % compartment
As air flows into compartment, it mixes w/ 100% O2 and pt receives % of O2 on setting
Size of trapezoid hole determine how much air from atmosphere dilutes 100% O2 from wall.
Larger trapezoid = less wall O2 -> less concentrated O2
Ex of Venturi effect
* not precise
Reynolds law
Predict if flow will be laminar or turbulent
Inertial forces/Viscous forces
Inertial forces: related to momentum of fluid; forces that cause fluid to move
Viscous forces: frictional shear forces that develop btw layers of fluid due to viscosity.
If viscous forces dominate flow -> laminar flow
Smaller number = laminar
Ex: flow of blood through vessels
Flow velocity at pipe wall - 0 (no slip condition). May velocity at center of tub -> parabolic profile is parabolic (center jets out from sides -> faster motion)
Lower pressure
If inertial forces dominate -> turbulent flow
Ex: flow of smoke from chimney
flow of air behind car traveling at high speed
Flatter velocity profile due to flow mixing
Higher pressure
Law of Laplace
Pressure inside spherical bubble is equal to 2x tension divided by radius -> increased pressure tension to pull water inward while pressure pushes outward
Larger radius = larger force requirement, larger wall stress, larger metabolic demand
Aneurysm -> increased radius -> more stress on tissue + more force required to keep blood flowing
Small radius of alveoli = large pressure -> collapse without surfactant
Large radius of alveoli = small pressure -> overinflation without surfactant
Surfactant reduces surface tension (which would cause lung to close in on itself without surfactant) -> prevents alveoli collapse, evenly distributes ventilation of alveoli, improvement of lung compliance
Wave motion
transfer energy through solids, liquids gases, or empty spaces (vacuums)
Frequency: # of complete waves passing a fixed point, in a given amount of time
Usually 1 second per complete wave
Measured in hertz: how many complete cycles per second
How often a wave happens
Period: time for 1 complete cycle Measured in seconds Time it takes for wave to occur If it occurs repeatedly, event is periodic Period of a day is 24hrs
Wavelength: distance btw point on one wave and same point on next wave symbol: lambda measured in meters x-rays are short radio waves are long
Amplitude: as waves travel they create disturbance. Distance btw maximum disturbance to undisturbed position
Flat sea and incoming sea -> amplitude is height of highest point on wave to flat sea
Transverse waves
Particles vibrate at 90 degrees to the direction that the energy is moving
Used to show amplitude and wavelength
Ex: waves of water (make water surface go up and down
Lightwaves
Longitudinal waves
Particles of energy vibrate parallel to the direction in which wave of energy is traveling
Waves moves horizontally, back and forward
Waves close together = compressions
Waves farther apart = rarefactions
Ex: sound waves
Doppler effect
Source emits sounds waves:
If source moves toward observer or observer moves toward source, frequency will increase
Frequency detected by observer is greater than the frequency of the source
Wavelengths decrease and frequency increases
If source is moving away from observer or the observer moves away from source, frequency will decrease
Wavelength increases and frequency decreases
Inverse Square law
Change in intensity due to change in distance
When X-ray beam is created at anode target, its created in isotropic way meaning is its creating equally in all directions
Xray beam is also divergent -> spreads apart and is less intense as it travels through space
When tube position changes:
- beam intensity changes
- patient dose changes
- receptor exposure changes (number of photons striking image receptor)
Ex: when distance from flashlight increases, light is less bright -> decrease in intensity
When distance from light decreases, light becomes more bright -> increase in intensity
