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263 Cards in this Set
- Front
- Back
matter exists in what 3 phases
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solid liquid and gas
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the state in which a compound exists depends on what
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the relative strength of the attractive forces between particles compared to the kinetic energy of the particles
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the attractive forces between particles in gases are
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very weak compared to their kinetic energy so particles move about more freely, are farther apart and have almost no influence on one another
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the attractive forces between particles in liquids are
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stronger, pulling particles close together but still allowing them considerable freedom to move about
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the attractive forces between particles in solids are
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much stronger that their kinetic energy of the particles, so the atoms/molecules/ions are held in a specific arrangement and can only wiggle around in one place
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changes of state (phase change)
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the change or transformation of a substance from one state of matter to another
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every change of state is ___ and characterized by what
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reversible and characterized by a free energy change, ∆G
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a change of state that is spontaneous in one direction (exergonic, negative ∆G) is
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non spontaneous in the other direction (endergonic, positive ∆G)
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formula for free energy change ∆G
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∆G= ∆H (enthalpy change) - T(in Kelvins) ∆S (entropy change)
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enthalpy change ∆H
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a measure of heat absorbed or released during a given change of state.
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when heat is absorbed, ∆H is
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positive- endothermic
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when heat is released, ∆H is
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negative- exothermic
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entropy change ∆S
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a measure of the change in molecular disorder or freedom that occurs during a process
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what happens to ∆S in the melting of a solid to a liquid
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disorder increases, ∆S is positive
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what happens to ∆S in the freezing of a liquid to a solid
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disorder decreases, ∆S is negative
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melting point (mp)
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the exact temperature at which solid and liquid are in equilibrium
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boiling point (bp)
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the exact temperature at which liquid and gas are in equilibrium
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sublimation
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when a solid changes directly to a gas without going through the liquid state. ex: dry ice
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intermolecular force
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a force that acts between molecules and holds molecules close to one another
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intermolecular forces in gases
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they are negligible (small and unimportant) so gas molecules act independently of one another
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intermolecular forces in liquids and solids
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they are strong enough to hold the molecules in close contact
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the stronger the intermolecular forces in a substance...
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the more difficult it is to separate the molecules and the higher the melting and boiling points of the substance
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3 major types of intermolecular forces
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dipole-dipole, london dispersion, and hydrogen bonding
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dipole-dipole force
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the attractive force between positive and negative ends of polar molecules
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dipole-dipole forces have what kind of strength
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they are weak with strengths around 1 kcal/mol
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london-dispersion force
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the short-lived attractive force due to the constant motion of electrons within molecules
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all molecules experience what type of intermolecular force
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london-dispersion force
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strength of london-dispersion force
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they are weak with strengths around 0.5-2.5 kcal/mol but increase with molecular weight and amount of surface area available for interaction between molecules
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what intermolecular force only happens between polar molecules
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dipole-dipole forces
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hydrogen bonds
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the attraction between a hydrogen atom bonded to an electronegative O, N, or F atom and another nearby O, N or F atom
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O-H, N-H, and F-H bonds are highly
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polar
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strength of hydrogen bonds
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they are strong with strengths around 10 kcal/mol
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gases have what kind of density, and what happens to volume when placed under pressure
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low density and are easily compressed to a smaller volume when placed under pressure
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gases undergo a larger_____ when their temp is changed than do liquids or solids
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they undergo a larger expansion or contraction
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kinetic-molecular theory of gas
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a group of assumptions that explain the behavior of gases
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kinetic-molecular theory of gas about their particles
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a gas consists of many particles, either atoms or molecules, moving about at random with no attractive forces between them- this explains when different gases mix together quickly
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kinetic-molecular theory of gas about their kinetic energy
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the average kinetic energy of gas particles is proportional to the Kelvin temperature- as temp increases, gas particles have more kinetic energy and move faster
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kinetic-molecular theory of gas about space
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the amount of space occupied by the gas particles themselves is much smaller than the amount of space between particles- most of the volume taken up by gases is empty space
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kinetic-molecular theory of gas about collision of gas particles
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collisions of gas particles, either with another particle or with the wall of the container, are elastic- the total kinetic energy of the particles is constant
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ideal gas
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a gas that obeys all of the assumptions of the kinetic-molecular theory- no such thing as a perfectly ideal gas because all gases behave differently
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when do most real gases display nearly ideal behavior
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under normal conditions
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Pressure (P)
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the force (F) per unit are (A) pushing against a surface
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formula for pressure (P)
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P= F/A
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why is the atmospheric pressure not constant
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because it varies slightly from day to day depending on weather, and varies with altitude
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where is the density of air greatest in the atmosphere
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at the earth's surface and decreases with increasing altitude
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what is the most commonly used unit of pressure
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the millimeter of mercury (mmHg)- often called a torr
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unit of pressure given in the SI system
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the pascal (Pa)
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1 pascal (Pa)= how many mmHg
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0.007500 mmHg
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1 atm (atmosphere) = how many mmHg, psi and Pa
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1 atm= 760 mmHg= 14.7 psi= 101,325 Pa
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gas laws
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a series of laws that predict the influence of pressure (P), volume (v) and temperature (T) on any gas or mixture of gases
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which gas law is the relation between volume and pressure
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boyle's law
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boyle's law
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the volume of a gas is inversely proportional to its pressure for a fixed amount of gas at a constant temperature- volume and pressure change in opposite directions
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according to boyle's law, what happens to volume as the pressure goes up
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the volume goes down
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equations for boyle's law
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PV=K (a constant value), P1V1=P2V2, P2= (P1V1)/V2, V2= (P1V1)/ P2
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which gas law is the relation between volume and temperature
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charles law
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charles law
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the volume of a gas is directly proportional to it's kelvin temperature for a fixed amount of gas at a constant pressure
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according to charle's law, is the kelvin temperature of a gas is doubled, what happens to its volume?
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it's doubled
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equations for charles law
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V/T=K, (V1/T1)=(V2/T2)
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which gas law is the relation between pressure and temperature
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Gay-lussac's law
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Gay-lussac's law
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the pressure of a gas is directly proportional to its kelvin temperature for a fixed amount of gas at a constant volume
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according to gay-lussac's law, as temperature of a gas goes up, the pressure will
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go up
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equations for gay-lussac's law
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P/T= K, (P1/T1)=(P2/T2)
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combined gas law
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when the relationships of PV, V/T, and P/T merge together when the amount of gas is fixed
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equations for the combined gas law
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PV/T=K, (P1V1)/T1= (P2V2)/T2
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which gas law is the relation between volume and molar amount
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avogadro's law
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avogadro's law
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the volume of a gas is directly proportional to its molar amount at a constant pressure and temperarture
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according to avogadro's law, a sample that contains twice the molar amount will have ____ the volume
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twice
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equations for avogadro's law
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V/n=K, (V1/n1)=(V2/n2)
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standard temperature and pressure (STP)
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temperature of 0° C (273.15 K) and a pressure of 1 atm (760 mmHg)
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standard molar volume of any ideal gas at STP
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at STP, 1 mol of any gas (6.02 x 10^23 particles) has a volume of 22.4 L- 22.4 L/mol
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the ideal gas law
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combination of pressure (p) volume (v) temperature (T) and number of moles (n)
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equations for ideal gas law
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PV/nT=R, PV=nRT
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gas constant (R)
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the constant R in the ideal gas law, PV=nRT.
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what does the gas constant (R) depend on? what are the values of R
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depends on the units of pressure. for pressure in atm, R= 0.0821 L*atm/mol*K. for pressure in mmHg, R= 62.4 L*mmHg/mol*K
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mixtures of gases vs pure gases- behavior
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mixture of gases behave the same as pure gases and obey the same laws
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partial pressure
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the contribution of a given gas in a mixture to the total pressure
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daltons law
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the total pressure exerted by a gas mixture (P-total) is the sum of the partial pressures of the components in the mixture- Ptotal= Pgas1 + Pgas2 + Pgas3....
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vapor
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the gas molecules are in equilibrium with a liquid
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vapor pressure
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the partial pressure of vapor molecules in equilibrium with a liquid
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what does vapor pressure depend on
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the temperature and the chemical identity of a liquid
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vapor pressure rises with increasing temperature until
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it becomes equal to the pressure of the atmosphere
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normal boiling point
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the boiling point at a pressure of exactly 1 atm or 760 mmHg
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viscosity
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the measure of a liquids resistance to flow
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nonpolar molecules have relatively ____ viscosities, polar molecules have relativley___ viscosities
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Low (gasoline), High (glycerin)
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surface tension
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the resistance of a liquid to spreading out and increasing its surface area
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what substance has the highest specific heat of any liquid
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water- giving it the capacity to absorb a large quantity of heat while changing only slightly in temperature
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water has what degree of heat of vaporization and what does it mean
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high heat of vaporization (540 cal/g) which means that is carries away a large amount of heat when it evaporates
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what is the most fundamental distinction between solids
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some are crystalline and some are amorphous
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crystalline solid
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solid whose particles (atoms/ions/molecules) are ridgedly held in an ordered arrangement
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crystalline solids can be categorized as...
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ionic, molecular, covalent network or metallic
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ionic solids
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those like sodium chloride, whose constituent particles are ions held together by ionic bonds
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molecular solids
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those like sucrose or ice, whose constituent particles are molecules held together by intermolecular forces
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covalent network solids
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those like diamonds or quartz (SiO2) whose atoms are linked together by covalent bonds into a giant 3-dimensional array- one very large molecule
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metallic solids
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such as silver or ion, can be viewed as vast 3-dimensional arrays of metal cations immersed in a sea of electrons that are free to move about
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amorphous solid
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a solid whose particles do not have an orderly arrangement
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how are amorphous solids usually formed?
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when liquids cool before they can achieve internal order or when their molecules are large and tangled together- glass, tar, hard candies
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how do amorphous solids differ from crystalline solids
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by the softening over a wide temperature range rather than having sharp melting points and by shattering to gives pieces with curves rather than planar faces
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heat of fusion
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the quantity of heat required to completely melt one gram of a substance once it has reached its melting point
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heat of vaporization
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the quantity of heat needed to completely vaporize one gram of a liquid once it has reached its boiling point
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a liquid with a low heat of vaporization
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evaporates rapidly and is said to be volatile
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what happens if a volatile liquid spills on your skin
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you feel a cooling effect as it evaporates because it is absorbing heat from your body
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equations for the energy needed to complete a phase change (melting and boiling)
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for melting- heat (cal or J)= Mass (g) x heat of fusion (cal or J/ g). for boiling- heat (cal or J)= Mass (g) x heat of vaporization (cal or J/g)
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mixture
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an intimate combination of two or more substances, both of which retain their chemical identities
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mixtures can be classified as two things
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heterogeneous mixtures or homogeneous mixtures
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heterogeneous mixture
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a nonuniform mixture that has regions of different composition- rocky road ice cream, granite
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homogeneous mixture
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a uniform mixture that has the same composition throughout- seawater
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homogeneous mixtures can be classified as what two things. according to what
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either solutions or colloids- according to the size of their particles
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solutions
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the most important class of homogeneous mixtures- a mixture that contains particles the size of a typical ion or small molecule- roughly 0.1-2 nm in diameter- ex: air, seawater, gasoline, wine
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colloids
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a homogeneous mixture that contains particles that range in diameter from 2 to 500 nm- ex: milk, fog, butter, pearl
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characteristics of solutions
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transparent to light, does not separate on standing, nonfilterable
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characteristics of colloids
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often murky or opaque to light, does not separate on standing, nonfilterable
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characteristics of heterogeneous mixtures
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murky or opaque to light, separates on standing, filterable, particle size is greater than 500 nm
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solute
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a substance that is dissolved in a solvent
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solvent
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the substance in which another substance (solute) is dissolved- ex: seawater- the dissolved salts are the solutes and the water is the solvent
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what determines whether a substance is soluble in a given liquid
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solubility depends on the strength of the attractions between solute and solvent particles relative to the strengths of the attractions within the pure substances
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rule of thumb for predicting solubility
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"like dissolves like"- substances with similar intermolecular forces form solutions with one another, whereas substances with different intermolecular forces do not
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solvation (or hydration specifically for water)
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the clustering of solvent molecules around a dissolved solute molecule or ion
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solid hydrates
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formed when ionic compounds attract water strongly enough to to hold on to water molecules even when crystalline
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in the solid hydrate, calcium sulfate hemihydrate→ CaSO₄ · ½H₂O, what does the dot mean
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the dot indicates that for every 2 CaSO₄ formula units in the crystal there is also one water molecule present
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hygroscopic
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compounds that have the ability to pull water molecules (vapor) from the surrounding atmosphere (humid air) so they can become hydrated
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miscible
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mutually soluble in all proportions- solute will continue to dissolve in solvent no matter how much is added
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saturated solution
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a solution that contains the maximum amount of dissolved solute at equilibrium
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solubility
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the maximum amount of a substance that will dissolve in a given amount of solvent at a specified temperature- usually expressed in grams per 100 mL (g/100mL)
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temperature and solubility on solids
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temperature has a dramatic effect on solubility- the effect is different for every substance and is usually unpredictable
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supersaturated solution
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a solution that contains more than the maximum amount of dissolved solute; a nonequilibrium situation
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unlike solids, the influence of temperature on the solubility of gases is
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predictable- addition of heat decreases the solubility of most gases
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the effect of pressure on the solubility of solids, liquids and gases
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pressure has no effect on the solubility of a solid or liquid, but has a strong effect on the solubility of a gas
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Henry's law
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the solubility (or concentration) of a gas is directly proportional to the partial pressure of the gas if the temperature is constant
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according to henry's law, if the partial pressure of a gas doubles, what else doubles
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solubility
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equations for henry's law
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C/Pgas=k, (C1/P1)=(C2/P2)
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percent concentrations express
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the amount of solute in one hundred units of solution- parts per hundred (pph)
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for solid solutions, percent concentrations are usually expressed as
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mass/mass percent concentrations (m/m)%
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mass/mass percent concentrations (m/m)%
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concentration expressed as the number of grams of solute per 100 grams of solution. (m/m)% concentration=mass of solute (g)/ mass of solution (g) x 100%
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the concentration of a solution made by dissolving one liquid in another is often expressed by
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volume/volume percent concentration (v/v)%
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volume/volume percent concentration (v/v)%
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concentration expressed as the number of milliliters of solute dissolved in 100 mL of solution. (v/v)% concentration= volume of solute (mL)/volume of solution (mL) x 100%
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mass/volume percent concentration (m/v)%
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concentration expressed as the number of grams of solute per 100 mL of solution- (m/v)% concentration= mass of solute (g)/ volume of solution (mL) x 100%
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when concentrations are very small, it is more convenient to use
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parts per million (ppm) or parts per billion (ppb)
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parts per million (ppm)
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number of parts per one million (10^6) parts- ppm= mass of solute/mass of solution x 10^6 or volume of solute/volume of solution x 10^6
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parts per billion (ppb)
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number of parts per one billion (10^9) parts- ppb= mass of solute/mass of solution x 10^9 or volume of solute/volume of solution x 10^9
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molarity (M)
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concentration expressed as the number of moles of solute per liter of solution- M=moles of solute/liters of solution (v)
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how would you find the moles of a solute
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molarity (moles per liter) x volume of solution (L)
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how would you find the volume of a solution
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moles of solute/molarity
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dilution
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adding additional solvent to lower the concentration- amount of solute remains constant, only volume is changed
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dilution factor
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the ratio of the initial and final solution volumes (Vc/Vd)
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electrolyte
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a substance that produces ions and therefore conducts electricity when dissolved in water
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the ability of a solution to conduct electricity depends on
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the concentration of ions in solution
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strong electrolyte
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substance that ionizes completely when dissolved in water ex: NaCl
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weak electrolyte
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a substance that is only partly ionized in water ex: molecular substances like acetic acid CH₃CO₂H
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nonelectrolyte
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a substance that does not produce ions when dissolved in water= ex: molecular substances like glucose
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Equivalent (Eq)
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for ions, one Eq is equal to the number of ions that carry 1 mol of charge
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gram-equivalent (g-Eq)
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for ions, the molar mass of the ion (in grams) divided by the ionic charge: 1 gram-equivalent of ion= molar mass of ion(g) / charge on ion
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how do you find the number of equivalents of a given ion per liter of solution
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Molarity of ion (M) x charge on ion
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one milliequivalent (mEq) of an ion
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1 mEq of an ion is 1/1000 of an equivalent or 1 mEq= 0.001 Eq, 1 Eq= 1000 mEq
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anion gap
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the difference in concentration of positive ions and negative ions- difference is made up by the presence of negatively charged proteins and the anions of organic acids
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colligative property
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a property of a solution that depends only on the number of dissolved particles (concentration of dissolved solute), not on their chemical identity
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4 colligative properties of solutions
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1. vapor pressure is lower for a solution than for a pure solvent 2. boiling point is higher for a solution than for a pure solvent 3. freezing point is lower for a solution than for a pure solvent 4. osmosis occurs when a solution is separated from a pure solvent by a semipermeable membrane
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why is the vapor pressure lower for a solution than for a pure solvent
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if some of the liquid (solvent) molecules at surface are replaced by other (solute) particles that don't evaporate, then the rate of evaporation of solvent molecules decreases and the vapor pressure is lower
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why is the boiling point for a solution higher than for a pure solvent
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it is a consequence of the vapor pressure lowering for a solution- because vapor pressure of a solution is lower than that of the pure solvent at a given temperature, the solution must be heated to a higher temperature for its vapor pressure to reach atmospheric pressure
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for each mole of solute particles added, what happens to the boiling point (equation)
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for each mol of solute added, the boiling point of 1 kg of water is raised by 0.51 ° C or ∆Tboiling=(0.51°C kg water/mol particles)(mol particles/kg water)
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why is the freezing point of solutions lower than freezing points of pure solvents
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because solute molecules are dispersed between solvent molecules throughout the solution, making it more difficult for solvent molecules to come together and organize into ordered crystals- lowering the freezing point
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for each mole of nonvolatile solute particles added, what happens to the freezing point (equation)
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for each mol of solute particles , the freezing point of 1 kg of water is lowered by 1.86 °C: ∆Tfreezing= (-1.86°C kg water/mol particles)(mol particles/ kg water)
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semipermeable
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a material or membrane allowing certain substances like water and other small molecules to pass through it, but they block the passage of large solute molecules or ions
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osmosis
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the passage of solvent through a semipermeable membrane separating two solutions of different concentration
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osmotic pressure (π)
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the amount of external pressure that must be applied to a solution to prevent the net movement of solvent molecules across a semipermeable membrane: π=(n/V)RT- v= solution volume n= # of moles of particles in solution R= gas constant T= absolute temperature of the solution
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what is convenient to use to describe the concentration of particles in solution
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osmolarity (osmol)
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osmolarity (osmol)
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the sum of the molarities of all dissolved particles in a solution
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isotonic
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having the same osmolarity
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hypotonic
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having an osmolarity less than the surrounding blood plasma or cells
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what is hemolysis
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a process where when red blood cells are placed in a hypotonic solution, water passes through the membrane into the cells, causing the cell to swell up and burst
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hypertonic
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having an osmolarity greater than the surrounding blood plasma or cells
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what is crenation
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a process where when red blood cells are in a hypertonic solution, water passes out of the cells into the surrounding solution, causing the cells to shrivel
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dialysis
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similar to osmosis, except that the pores in a dialysis membrane are larger than those in an osmotic membrane so that both solvent molecules and small solute particles can pass through it, but large colloidal particles cannot pass
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according to the Arrhenius definition, what is an acid
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a substance that produces hydrogen ions, H⁺ when dissolved in water
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according to the Arrhenius definition, what is a base
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a substance that produces hydroxide ions, OH⁻ when dissolved in water
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according to the Arrhenius definition, a neutralization reaction of an acid with a base yields
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water plus a salt
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what is a salt
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an ionic compound composed of the cation from the base and the anion from the acid
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why are swedish chemist, Svante Arrhenius's definition of acids and bases limited
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because they only refer to reactions that take place in aqueous solutions
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hydronium ion
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the H₃O⁺ ion, formed when an acid reacts with water
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acids generally have what kind of taste
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sour- nearly every sour food contains an acid
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bases are present in many
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household cleaning agents- perfumed bar soap, ammonia based window cleaners, substances you put down the drain to dissolve hair and grease
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what are the 9 most common acids and bases
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1. sulfuric acid (H₂So₄) 2. hydrochloric acid (HCl) 3. phosphoric acid (H₃PO₄) 4. nitric acid (HNO₃) 5. Acetic acid (CH₃CO₂H) 6. sodium hydroxide (NaOH) 7. calcium hydroxide (Ca(OH)₂) 8. magnesium hydroxide (Mg(OH)₂) 9. Ammonia (NH₃)
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what is sulfuric acid
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(H₂So₄)- most important raw material and most manufactured chemical. used in hundreds of industrial processes including preparation of phosphate fertilizers, Most common consumer use is as the acid found in automobile batteries
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what is hydrochloric acid
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(HCl)- historically known as muriatic acid- many industrial applications including its use in metal cleaning and in the manufacture of high-frutose corn syrup. Also present as stomach acid in the digestive systems of most mammals
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what is phosphoric acid
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(H₃PO₄)- used in the manufacture of phosphate fertilizers. Also used as an additive in foods and toothpastes- the tart taste of many soft drinks is due to the presence of phosphoric acid
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what is nitric acid
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(HNO₃)- strong oxidizing agent that is used for many purposes including the manufacture of ammonium fertilizer and military explosives. when spilled on skin, it leaves a characteristic yellow coloration because of its reaction with skin proteins
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what is acetic acid
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CH₃CO₂H- the primary organic constituent of vinegar. Also occurs in all living cells and is used in many industrial processes such as the preparation of solvents, lacquers and coatings
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what is sodium hydroxide
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NaOH- also called caustic soda or lye- most commonly used of all bases. Industrially it is used in production of aluminum from its ore, production of glass and in manufacture of soap from animal fats. Often found in drain cleaners because it reacts with the fats and proteins found in grease and hair
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what is calcium hydroxide
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Ca(OH)₂- or slaked lime, its made industrially by treating lime (CaO) with water. Many applications, including its use in mortars and cements
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what is magnesium hydroxide
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Mg(OH)₂- or milk of magnesia. an additive in foods, toothpastes and many over the counter medications.
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what is Ammonia
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NH₃- used primarily as a fertilizer, but it also has many other applications including the manufacture of pharmaceuticals and explosives. dilute solution is frequently used as a glass cleaner
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bronsted-lowry acid
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a substance that can donate a hydrogen ion H⁺ to another molecule or ion- also called a proton donor
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monoprotic acids
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acids with one proton to donate, such as HCl or HNO₃
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diprotic acids
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acids that have two protons to donate, such as H₂SO₄
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triprotic acids
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acids that have three protons to donate such as H₃PO₄
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bronsted-lowry base
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a substance that can accept H⁺ ions from an acid
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an acid-base reaction is one in which a ____ is transferred
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proton
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a bronsted-lowry base can be two things
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neutral or negatively charged
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if a bronsted-lowry base is neutral, then the product has a
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positive charge after H⁺ has been added
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if the bronsted-lowry base is negatively charged, then the product
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is neutral
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conjugate acid-base pairs
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two substances whose formulas differ only by a hydrogen ion H⁺- they are found on opposite sides of the chemical equation
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conjugate base
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the substance formed by the loss of H⁺ from an acid
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conjugate acid
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the substance formed by the addition of H⁺ to a base
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the number of protons in a conjugate acid-base pair is always
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one greater than the number of protons in the base of the pair
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what acids and bases are highly corrosive (3) and how do they react
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sulfuric acid (H₂SO₄), hydrochloric acid (HCl), sodium hydroxide (NaOH). They react readily and can cause serious burns to skins
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why are some acids and bases relatively safe while others must be handled with extreme cation?
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answer lies in how easily they produce the active ions for an acid (H⁺) or base (OH⁻)
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strong acid
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an acid that gives up H⁺ easily and is essentially 100% dissociated in water
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dissociation
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the splitting apart of an acid in water to give H⁺ and an anion
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weak acid
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an acid that gives up H⁺ with difficulty and is less than 100% dissociated in water
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weak base
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a base that has only a slight affinity for H⁺ and holds it weakly
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strong base
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a base that has a high affinity for H⁺ and holds it tightly
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the stronger the acid, the ____ its conjugate base; the weaker the acid, the _____ the conjugate base
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the stronger the acid, the weaker its conjugate base; the weaker the acid, the stronger the conjugate base
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a strong acid and a weak base cause the reverse reaction to
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occur at a lesser extent
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a weak acid and a strong base causes the reverse reaction to
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occur more readily
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an acid-base proton transfer equilibrium always favors the
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reaction of the stronger acid with the stronger base and formation of the weaker acid and base
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in a contest for protons, the _____ base always wins
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the stronger base
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acid dissociation constant (Ka)
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the equilibrium constant for the dissociation of an acid (HA), equal to [H⁺][A⁻]/[HA]
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important points for Ka values (4)
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1. strong acids have Ka values much greater than 1 because dissociation is favored 2. weak acids have Ka values much less than 1 because dissociation is not favored 3. donation of each successive H⁺ from a polyprotic acid is more difficult than the one before it, so Ka values become successively lower 4. most organic acids, which contain the CO₂H group, have Ka values near 10⁻⁵
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in bronsted-lowry definition, water can as
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both an acid and a base
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when in contact with a base, water reacts as
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a bronsted-lowry acid and donates a protein to the base
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when in contact with an acid, water reacts as
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a bronsted-lowry base and accepts H⁺ from the acid
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amphoteric
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a substance that can react as either an acid or a base, such as water
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what happens to water when no other acids or bases are present
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one water molecule acts as an acid while another water molecule acts as a base, reacting to form hydronium (H₃O⁺) and hydroxide (OH⁻) ions
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ion-product constant for water (Kw)
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the product of the H₃O⁺ and OH⁻ molar concentrations in water or any aqueous solution: Kw=K[H₂O][H₂O]= [H₃O⁺][OH⁻]= 1.00 x 10⁻¹⁴ at (25° C)
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solutions are identified as..... according to what
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acidic, neutral or basic (alkaline), according to their value of their H₃O⁺ and OH⁻ concentrations
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concentration of acidic solution
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[H₃O⁺]> 10⁻⁷ M and [OH⁻]< 10⁻⁷ M
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concentration of neutral solution
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[H₃O⁺]= 10⁻⁷ M and [OH⁻]= 10⁻⁷ M
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concentration of basic solution
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[H₃O⁺]< 10⁻⁷ M and [OH⁻]> 10⁻⁷ M
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the ______ of an aqueous solution is a number usually between 0 and 14 that indicates the H₃O⁺ concentration of a solution
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the pH
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pfunction
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mathematically defined as the negative common logarithm of some variable
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pH
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a measure of the acid strength of a solution; the negative common logarithm of the H₃O⁺ concentration
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pH for an acidic solution
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pH<7, [H₃O⁺]> 1 x 10⁻⁷ M
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pH for a neutral solution
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pH= 7, [H₃O⁺]= 1 x 10⁻⁷ M
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pH for a basic solution
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pH>7, [H₃O⁺]< 1 x 10⁻⁷
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converting from pH to [H₃O⁺] requires finding the
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antilogarithm of the negative pH
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converting from [H₃O⁺] to pH requires finding the
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logarithm
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significant figures for an antilogarithm
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contains the same number of sig figs as the original number has to the right of the decimal point
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significant figures for a logarithm
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contains the same number of digits to the right of the decimal point as the number of sig figs in the original number
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what is the simplest but least accurate method of measuring the pH of a solution
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using an acid-base indicator
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acid-base indicator
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a dye that changes color depending on the pH of a solution
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universal indicator
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makes pH determination easy, test kits that contain a mixture of indicators to give approximate pH measurements in the range 2-10
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buffer
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a combination of substances that act together to prevent a drastic change in pH; usually a weak acid and its conjugate base
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Henderson-Hasselbalch equation
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the logarithmic form of the Ka equation for a weak acid, used in applications involving buffer solutions: pH=pKa - log([HA]/[A⁻]) or pH= pKa + log([A⁻]/[HA])
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what conditions do most effective buffers meet (3)
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1. the pKa for the weak acid should be close to the desired pH of the buffer solution 2. the ratio of [HA] to [A⁻] should be close to 1, so that neither additional acid nor additional base changes the pH of the solution dramitically 3. the molar amounts of [HA] and [A⁻] in the buffer should be approximately 10 times greater than the molar amounts of either acid or base you expect to add so that the ratio [A⁻]/[HA] does not undergo a large change
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the pH of body fluids is maintained by how many buffer systems
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3
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what two buffer systems depend on weak acid conjugate base interactions exactly like those of the acetate buffer system
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the carbon acid- bicarbonate system (H₂CO₃-HCO₃) and the dihydrogen phosphate-hydrogen phosphate system (H₂PO₄-HPO₄²⁻)
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equivalent of acid
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amounts of acid that contains 1 mole of H⁺ ions
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1 gram-equivalent of acid
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= molar mass of acid (grams)/ number of H⁺ ions per formula unit
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equivalent of base
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amount of base that contains 1 mole of OH⁻ ions
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1 gram-equivalent of base
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= molar mass of base (grams)/ number of OH⁻ ions per formula unit
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1 equivalent of any acid does what to any base
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1 equivalent of any acid neutralizes one equivalent of any base
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normality (N)
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a measure of acid (or base) concentration expressed as the number of acid (or base) equivalents per liter of solution: N= equivalents of acid or base/ liters of solution
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the values of molarity (M) and normality (N) are the same for _____ acids, but are not the same for _____ acids
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M and N are the same for monoprotic acids, but are not the same for diprotic or triprotic acids
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normality of acid
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= (molarity of acid) X (the number of H⁺ ions produced per formula unit)
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normality of base
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= (molarity of base) X (the number of OH⁻ ions produced per formula unit)
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what are the most common kinds of bronsted-lowry acid-base reactions (3)
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1. those of an acid with hydroxide ion 2. an acid with bicarbonate or carbonate ion 3. an acid with ammonia or a related nitrogen containing compound
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reaction of an acid with hydroxide ion
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1 equivalent of an acid reacts with 1 equivalent of a metal hydroxide to yield water and a salt in a neutralization reaction
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reaction of an acid with bicarbonate and carbonate ions
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bicarbonate reacts with acid by accepting H⁺ to yield carbonic acid, H₂CO₃. Carbonate accepts 2 protons in reaction with an acid
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reaction of an acid with ammonia
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acid reacts with ammonia to yield ammonium salts
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titration
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a procedure for determining the total acid or base concentration of a solution
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when titration involves a neutralization reaction in which 1 mole of acid reacts with 1 mole of base, then
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(M-acid X V-acid)= (M-base X V-base)
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when coefficients for the acid and base in a balanced neutralization reaction are not the same, we use
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equivalents of acid and base instead of moles and normality instead of molarity: (Eq)acid= (Eq)base or (N-acid X V-acid)= (N-base X V-base)
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salt solutions can be... depending on what?
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neutral, acidic or basic, depending on the ions present because some ions react with water to produce H₃O⁺ and some ions react with water to produce OH⁻
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general rule for predicting acidity or basicity of a salt solution
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the stronger partner from which the salt is formed dominates- a salt formed from a strong acid and weak base yields an acidic solution, and vise versa. a salt formed from a strong acid and a strong base yields a neutral solution
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