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YSU / Chemistry / CHEM 1515 / What is the value of temperature and pressure at stp?

What is the value of temperature and pressure at stp?

What is the value of temperature and pressure at stp?


Study Guide 2 

What is the value of temperature and pressure at stp?

Concentration of Solutions 

∙     Solution: A homogeneous mixture consisting of a solute dissolved in a  solvent

o Solute: the substance being dissolved; the substance present in the  smaller amount  

o Solvent: the substance doing the dissolving; the substance present in  the lower amount

∙     Concentration: the amount of solute dissolved in a given amount of solvent  or solution If you want to learn more check out What are the terms pertaining to the different levels of ecological organization?

o Concentrated solutions have more dissolved solute particles than a  dilute solution

Expressing Concentration  

∙     Molarity: expresses the concentration as the number of moles of solute  particles dissolved in 1 liter of solution

What is the relationship between density and molar mass?

o Molarity=moles of solute

liter of solution

∙ Weight percent: Another way of expressing concentration in a solution owt =mass of solute

mass of solution×100


∙ Dilution: the process of preparing a less concentrated solution o in a dilution, the volume and concentration change, although, the  number of moles remains constant

∙ Serial Dilutions: a method used to obtain a series of increasingly dilute  solutions  


∙     Electrolyte: A substance that conducts an electric current when dissolved in  water

What are reactants and products in a reaction?

If you want to learn more check out What are anorexia nervosa and bulimia nervosa?

∙     Nonelectrolyte: A substance that does not conduct electricity

∙     Strong electrolyte: a compound that dissociates completely into ions in an  aqueous solution  

o All water-soluble ionic compounds are strong electrolytes

o Strong acids are also strong electrolytes

∙     Weak electrolytes: compound that partially dissociate in water Acids and Bases 

∙ Arrhenius Acids and Bases

o Acid: A substance that produces H+ ions in an aqueous solution o Base: A substance that produces OH- ions in an aqueous solution ∙ Bronsted Acids and Bases

o Acid: a proton (H+) donor

o Base: a proton acceptor

∙ Monoprotic acids: each acid molecule has one proton to donate  ∙ Polyprotic acids: acids that donate more than one proton

o Diprotic and triprotic acids: two and three protons per acid molecule  Chemical Analysis  

∙ Quantitative analysis: experiments that measure the amount of a  substance present

∙ Titration: used in the quantitative analysis of acids and bases  o Standard solution: solution of known concentration We also discuss several other topics like What makes humans so different from other animals?


∙ Equivalence point: the point in the titration when the acid has been  completely neutralized by the base  

∙ Indicator: a substance that has a distinctly different colors in acidic and  basic media  We also discuss several other topics like What is the indicator that the task or objective of the group is well understood and accepted?

o Endpoint: the point in the titration when the indicator changes color  Used to indicate the equivalence point

 Not the same as equivalence point


∙ Most substances are liquid or solid at room temp (25 C) but the can exist as a gas under certain conditions Don't forget about the age old question of What are the reactions for alkali metals and halogen reactions?

Characteristics of Gases

∙ A sample of gas takes the shape and volume of its container o Particles have no fixed positions

o Gas flows like liquid

∙ Gases are compressible

o Gas particles have large distance between them

o In a smaller volume the particles become closer

∙ Gases are less dense than liquids and solids

o Densities of gases are highly dependent on pressure and temperature  Expressed as g/L

∙ Gases form homogeneous mixtures (solutions) with each other in  any proportion

o Gases can mix uniformly with other gases We also discuss several other topics like What are the four types of goods in an economy?

o Molecules are so far apart that there is very little interaction between  them

Gas Pressure 

∙ Gas pressure: the force applied, per unit area, against a surface o Gas pressure is a result of the force exerted by the collisions of the gas particles with the walls of the container

o Atmospheric pressure: the pressure exerted by Earth’s atmosphere

Gas Pressure: Units 

∙ Gas pressure is most commonly measured in millimeters of mercury (mm  Hg), atmospheres (atm), and torr.

o Standard atmospheric pressure = 1 atm

o Unit conversions you will be responsible for: 760 mmHg = 760 torr = 1 atm

 Other units of pressure:  

1 atm = 760 mm Hg

= 760 torr

Gas Laws: Boyle’s Law 

= 101,325 pascals = 28.96 in. Hg

∙ Boyle’s Law: For a fixed mass of gas at a constant temperature, the volume is inversely proportional to the pressure.

o At constant temperature, the product of pressure and volume will  remain constant

Charles’s Law 

∙ Charles’s Law: For a fixed mass of gas at a constant pressure, the volume is  directly proportional to the temperature in kelvin (K).

o V ∝ T

Gay Lussac’s Law 

∙ Gay-Lussac’s Law: For a fixed mass of gas at constant volume, the pressure is directly proportional to the temperature in kelvin (K).

o Like Charles’s Law:

 P ∝ T

Avogadro’s Law 

∙ Avogadro’s law: Equal volumes of gas at the same temperature and  pressure contain the same number of gas molecules.

o Equal volumes of gas at the same temperature and pressure contain  the same number of moles gas.

o The number of gas particles is directly proportional to the volume of  the container

 V ∝ n

Combined Gas Law 

∙ Boyle’s law, Charles’s law, Gay-Lussac’s law, and Avogadro’s Law can be  combined into one law called the combined gas law.

o When dealing with a fixed quantity of gas (constant mass), the  equation reduces to the more common form of the combined gas law.  PV/T = a constant

Standard Temp and Pressure 

∙ The actual temperature and pressure at which we compare two or more  gases does not matter.

o For convenience in making comparisons, chemists have selected one  pressure as a standard pressure, and one temperature as a standard  temperature.

o The standard temperature and pressure (STP) selected are 0°C (273 K)  and 1 atm pressure.

∙ All gases at STP, or any other combination of pressure and temperature,  contain the same number of molecules in the same volume

o One mole contains 6.022 x 1023 formula units; what volume of gas at  STP contains this many molecule

o This quantity has been measured and found to be 22.4 L

o Thus, one mole of any gas at STP occupies 22.4 L 

Ideal Gases 

∙ Ideal gas: a hypothetical sample of gas whose pressure-volume-temperature behavior is predicted accurately by the ideal gas equation.

o The behavior of real gases differs from what is predicted, however,  only slightly  

o The proportionality constant (R): Derived from the fact that1.00  mol of any gas at STP occupies 22.4 L.

Density and Molar Mass

∙ The density of a gas can be determined if the molar mass of a gas (g/L) is  known

o The ideal gas equation can be rearranged to solve for (mol/L) o Multiplying by the molar mass (M) gives the density of the gas   Therefore…

d=P M

RT or MM=dRT


Reactions with Gaseous Reactants and Products 

∙ Avogadro’s Law: The volume of a gas at a given temperature and pressure  is proportional to the number of moles

o 2CO (g) + O2 (g) → 2CO2 (g)

o 2 moles CO to 1 mole O2 or 2 liters CO to 1 liter O2 or 2 mL CO to 1 mL O2

o These ratios can be used to determine stoichiometric amounts of  reactants or products

∙ Determining the stoichiometric amounts of reactants or products must be  calculated differently in equations where only one of the reactants is a gas  o 2Na (s) + Cl2 (g) → 2 NaCl (s)

∙ The ideal gas equation must be used along with the balanced chemical  equation (mole ratios)

oPV=nRT so n=PV


Gas Mixtures 

Daltons Law of Partial Pressure

∙ When two or more gaseous substances occupy the same container, each  behaves independently

∙ Daltons Law of Partial Pressures: The total pressure, P, of each individual  gas:


PT = P1 + P2 + P3 + . . .

Mole Fractions 

∙ Mole fractions specify the relative amounts of the components of a gas  mixture  

∙ The mole fraction of a component in a mixture is the moles of that  component divided by the total number of moles in the mixture



∙ Three things to remember about mole fractions:

o The mole fraction of a mixture component is always less than 1 o The sum of the mole fractions for all the components in a mixture is  always 1

o Mole fractions are dimensionless (no units)

∙ Avogadro’s Law: the number of moles and the pressure of a gas are  proportional at constant temperature and pressure (n ∝ P).

o The mole fraction may also be determined by dividing the partial  pressure of a component divided by the total pressure.



Xi×ntotal=ni and Xi×Ptotal=Pi 

Collection Of A Gas Over Water 

∙ The volume of a gas produced by a chemical reaction can be measured with  an apparatus like the one shown here.

o For example, the decomposition of potassium chlorate, KClO3, produces oxygen gas:

 2KClO3 (s) → 2KCl (s) + 3O2 (g)

∙ The oxygen gas produced is collected over water

o the volume of water displaced is equal to the volume of gas produced o the water levels, both inside and outside the collection flask, must be  the same before the volume collected volume can be read

 ensures that the pressure inside the inside the collection flask is  the same as the atmospheric pressure

∙ The measured volume is the result of both oxygen gas and water vapor ∙ The pressure inside the collection flask is the sum of the partial pressures of  the oxygen gas and water vapor


o the total pressure is equal to the atmospheric pressure

o the partial pressure, or vapor pressure, of water can be determined  from a table of known values

 vapor pressure of water is dependent on temperature

Kinetic Molecular Theory  

∙ The gas laws allow us to predict the behavior of gases under varying  conditions.

o They do not explain why gases behave the way they do.

∙ 19th Century physicists including Ludwig Boltzmann and James Maxwell,  develop a theory that explains the behavior of gases:  

o Kinetic Molecular Theory (KMT): explains why gases behave the  way they do based on their molecular nature.

o Enables us to understand some of the properties and behaviors of  gases using a few basic assumptions

∙ A gas is composed of particles that are separated by relatively large  distances; the volume occupied by an individual gas molecule is  negligible.

o Explains the compressibility of gases

 Gases can be compressed by reducing their volume

 Decreasing the volume forces the molecules closer together ∙ Gas molecules are constantly in random motion, moving in straight  paths, colliding with the walls of their container and with one  another in perfectly elastic conditions.

o energy is transferred, not lost in an elastic collision.

∙ Explains Boyle’s Law:  

o the volume of a gas is inversely proportional to its pressure

V ∝1P

Boyles Law 

The pressure of a gas is a result of the collisions of the gas molecules with the  walls of the container

o The magnitude of the pressure depends on:

 the frequency of the collisions

 the speed of the molecules when they collide with the walls

∙ Decreasing the volume increases the frequency of the collisions between the  gas molecules and the walls of the container

o pressure increases

Kinetic Molecular Theory  

∙ Gas molecules do not exert attractive or repulsive forces on one  another.

∙ Explains Dalton’s Law of Partial Pressures:  

oPtotal=Σ Pi 

 Since gas molecules do not attract or repel one another, the  pressure of one gas does not affect the other

∙ The average kinetic energy, E´k , of gas molecules in a sample is  

proportional to the absolute temperature.

o kinetic energy is the energy associated with motion

∙ Explains Charles’s Law

oV ∝T

∙ Heating a sample of gas increases its average kinetic energy o gas molecules move faster

o faster-moving molecules collide with the walls of the container more  frequently and with greater speed

o if the container can expand, the volume will increase

 decreasing the frequency of the collisions until the pressure  inside and outside of the container are equal

Mean Square Speed 

∙ The kinetic energy of an individual gas molecule is proportional to its mass  and it velocity squared


∙ The average kinetic energy of a group of gas molecules is determined using  the mean square speed, ´v2, the average speed squared for all the  

molecules in the sample.



 Where N is the number of molecules in the sample

Molecular Speed 


∙ The total kinetic energy of one mole of any gas is equal to  


∙ This, along with assumption 4, gives rise to an equation for the speed of a  molecule with the average kinetic energy in a gas sample; the root-mean square (rms) speed ( vrms¿ .

o vrms=√3 RT 



o where R=8.314


K ∙mol

o M = molar mass (in kg/mol)

∙ The rms speed is directly proportional to the square root of the absolute  temperature.

∙ The rms speed is inversely proportional to the square root of M (molar mass).

o ∴ For any two samples of gas at the same temperature, the  gas with the larger molar mass will have the lower rms speed. Molecular Speed: Comparing Gases 

∙ Any two gas samples at the same temperature have the same kinetic energy. o Gives rise to the following equation is used to compare the vrms values of molecules in different gas samples at a given temperature. vrms(1)

vrms(2)=√M2 M1 

Diffusion and Effusion 

∙ Two phenomena that can be observed in gases, diffusion and effusion. o Diffusion: the mixing of gases as a result of the random motion and  frequent collisions of gas molecules.

o Effusion: the escape of gas molecules from a container to a region of  vacuum

∙ Graham’s Law: the rate of effusion or diffusion of a gas is inversely  proportional to the square root of its molar mass.

oRate∝1√ M

 ∴ lighter /gases effuse and diffuse more rapidly than heavy  gases (they move faster)

Energy and Energy Changes 

∙ Energy: The capacity to do work or transfer heat

o Energy can be Kinetic or Potential

∙ Chemistry is the study of matter and the changes that matter undergoes. ∙ Every change that matter undergoes is accompanied by an exchange of  energy

o Both chemical and physical changes involve energy exchanges ∙ Some changes require, or absorb, energy

o The melting of ice requires energy in the form of heat

 H2O (s) + Heat → H2O (l)

∙ Some changes release, or produce, energy

o The combustion of methane releases energy in the form of heat   CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g) + Heat

Kinetic Energy

Kinetic energy (KE) is the energy of motion

∙ KE increases as the objects velocity (speed) increases

o At the same velocity, a heavier object has greater KE

Ek=12m v2 

∙ Examples of kinetic energy:

o An airplane in flight

o Water falling from a waterfall

o A ball thrown by a pitcher

∙ Thermal energy: a form of kinetic energy associated with the random  motion of particles

Potential Energy  

Potential energy: stored energy; the energy an object has because of its position  or composition.

∙ Examples of potential energy include:

o A stretched rubber band

o Water at the top of a waterfall

o A coiled spring

o A ball at the top of a hill

∙ In chemistry, the most important forms of potential energy are chemical  energy and electrostatic energy 

∙ Chemical energy is stored within the structural units of chemical  substances as in, for example, foods such as carbohydrates and fats. o Energy is given off when these substances take part in chemical  reactions.

o The amount of chemical energy in a sample depends on the and type  of atoms and the way they are arranged

∙ Electrostatic energy is the energy that results from the interaction of  charged particles


o The law of conservation of energy:

o Energy can neither be created nor destroyed

∙ Energy can be transferred or transformed from one form to another o Kinetic and potential energy are interconvertible

 Examples:  

∙ Dropping an object converts potential energy to kinetic  


∙ Chemical reactions that give off heat convert chemical  

(potential) energy to thermal (kinetic energy)


∙ The specific part of the universe we are interested in when studying energy  changes is defined as the system

o This usually includes the substances involved in chemical and physical  changes

 Example: an acid base reaction in a beaker; the beaker is  

defined as the system

∙ The rest of the universe outside of the system is defined as the  surroundings

∙ There are three types of systems; open, closed, and isolated o Open systems  

 Exchange both heat and mass with the surroundings

o Closed systems

 Exchange heat with surroundings (not mass)

o Isolated systems

 Do not exchange heat or mass with the surroundings


∙ Heat is a form of energy.

o Heat is the transfer of thermal energy that occurs when two objects of  different temperature are brought into contact.

o Thermochemistry is the study of the heat (the transfer of thermal  energy) in chemical reactions

∙ Endothermic processes require (absorb) heat  

o Heat must be supplied to the system (the system absorbs heat) o Heat is a reactant

 Heat + 2HgO (s) → 2Hg (l) + O2 (g)

∙ Exothermic processes give off (produce) heat

o Heat is transferred from the system to the surroundings

o Heat is a product

 2H2 (g) + O2 (g) → 2H2O (l) + Heat

Units of Energy 

∙ The joule (J) is the SI unit for energy

o The amount of kinetic energy in a 2 kg object moving at a speed of 1  m/s

o The magnitude of a joule is so small that it is often expressed as  kilojoules (kJ)

∙ The calorie: one calorie is equal to the amount of heat energy required to  increase the temperature of 1 gram of water by 1oC.

o 1 cal = 4.18 J

o The kilocalorie (kcal) also known as the large Calorie is the calories  listed on food packages

State Functions 

∙ State functions: properties that are determined by the state of a system,  regardless of the how conditions are achieved

o When the state of a system changes, the magnitude of the change  depends only on the initial and final states of the system

∙ The state of a system is defined by all of the values of all relevant  macroscopic properties such as:

o composition  

o energy

o temperature

o pressure

o volume

∙ Energy, pressure, volume, and temperature are state functions o Depend only on the initial and final conditions

 Example: elevation depends on position

 Change in elevation depends only on initial and final position,  not on the path taken to get there

∙ State functions do not depend on the path by which they are achieved ∙ Heat and work are not state functions

o Work = force x distance

o Distance traveled depends on the path taken

First Law of Thermodynamics 

∙ The first law of thermodynamics: energy can be converted from one form  to another, but it can not be created or destroyed.

o Based on the law of conservation of energy

o The change in internal energy (∆U) of a system is given by:

 ∆U = Uf – Ui

o Where Uf and Ui are the internal energies of the system in their final  and initial states respectively

 The symbol ∆ indicates the change in something (final – initial) o ∆U is a state function (depends on initial and final states)

∙ The internal energy of a system consists of two parts: kinetic energy and  potential energy  

o It is impossible to accurately measure either of these accurately o Changes in internal energy, however, can be determined  


Changes in Internal Energy 

∙ Consider the following reaction: S (s) + O2 (g) → SO2 (g) ∙ The internal energy of the system (molecules of the reactants and products)  is unknown  

∙ The change in energy can be determined: ∆U = U (products) – U  (reactants) 

∙ This reaction gives off heat: ∆U < 0 

∙ Indicates that the energy of the reactants is higher than the energy of the  products

∙ There is a net release of energy (exothermic reaction)

∙ Some of the energy from the reactants has been converted to thermal energy ∙ This transfer of energy does not change the total energy of the universe; the  sum of the energy change is zero

o ∆U system + ∆U surroundings = 0

∙ Thermal energy lost by the system is absorbed by the surroundings o ∆U system = - ∆U surroundings

∙ Energy released in one place must be gained somewhere else o energy is in terms of the system

o ∆U usually refers to ∆U system

∙ Energy lost by the system is gained by the surroundings, -∆U ∙ The system may gain energy from the surroundings, +∆U o If the system undergoes an energy change, the surroundings must  experience an energy change that is equal in magnitude but opposite  sign

 ∆U system = - ∆U surroundings

Work and Heat 

∙ Energy: the capacity to do work or transfer heat.

∙ Changes in work and heat lead to changes in internal energy o The overall change in internal energy is given by: ∆U = q + w  q = heat (released or absorbed by the system)

 w = work (done on the system or done by the system)  

∙ Sign conventions of q and w

o (+) q = endothermic processes (heat gained by the system)

o (–) q = exothermic processes (heat lost by the system)

o (+) w = work done on the system (by the surroundings)

o (–) w = work done by the system (on the surroundings)

Heat and Work 

∙ Internal energy decreases if the system does work (on the surroundings) or  releases heat (to the surroundings)

o These processes are energy depleting

∙ Internal energy increases if heat is gained by the system (from the  surroundings) or work is done on the system (by the surroundings.

Pressure Volume Work 

∙ To calculate internal energy, U, we need to know the signs of both heat, q,  and work, w.

o ΔU = q + w

o q can be determined by measuring temperature changes.

∙ To determine w we need to know whether the reaction occurs under constant volume or constant-pressure conditions.

o w = -PΔV

 p = external (opposing) pressure

 ΔV = change in volume

o an increase in volume gives –w (work done by the system)

o a decrease in volume gives +w (work done on the system)

o at constant volume PΔV = 0 (no work has been done)

∙ Constant Pressure:

o If the same reaction occurs in a vessel with a movable piston (allows  for changes in volume) the result is an increase in volume rather than  pressure.

 the reaction (system) has done work on the surroundings

 qp = ∆U + P∆V

 qp denotes heat at constant pressure

 changes in volume constitute changes in distance (work)

∙ Constant Volume:

o The internal energy of a system is given by:

 ΔU = q − PΔV

o For a reaction at constant volume no work is done

o ΔU = qv

 qv denotes heat at constant volume

 at constant volume, -PΔV = 0

 qv is a state function (q is not)

∙ ΔU is a state function


∙ Enthalpy (H): a thermodynamic function defined as the internal energy of a  system plus the pressure-volume work.

o H = U + PV

∙ Since both U and PV have energy units, enthalpy also has energy units (J) o U, P, and V are all state functions  

∙ Changes depend only on initial and final states

∙ Changes in enthalpy (∆H) also depend on only initial and final states (state  function)

o ∆H = ∆U + ∆(PV)

∙ At constant pressure the heat exchanged between the system and  surroundings is equal to the enthalpy change:

o ∆H = ∆U + P∆V or qp = ∆H  

∙ Enthalpy of Reaction:

∙ Enthalpy of reaction (∆H): the change in enthalpy for a reaction o Defined as the difference between the enthalpies of the products and  the enthalpies of the reactants.

 ∆H = H(products) – H(reactants)

o For an endothermic process (heat absorbed by the system)

 ∆H > 0

o For an exothermic process (heat is released by the system)

 ∆H < 0

Molar Enthalpy 

∙ At temperatures above 0 ˚C, 6.01 kJ of heat energy is absorbed for every one mole of ice converted to liquid water:  

o H20 (s) → H20 (l)  

∙ ∆H = +6.01 kJ/mol

∙ ∆H > 0; the reaction is endothermic

Thermochemical Equations 

∙ Thermochemical equations: chemical equations that show enthalpy  changes as well as the mass relationships.

∙ A balanced chemical equation is necessary

∙ Guidelines for using thermochemical equations:

∙ 1. The physical states of all reactants and products must be known o They help determine the actual energy change

∙ 2. If the chemical equation has been multiplied by a factor, n, ∆H must also  change by the same factor.



∙ 3. When a chemical equation is reversed, the roles of the reactants and  products change

o The magnitude of ∆H must also change

∙ Example: if a reaction absorbs heat from its surroundings (+∆H), then the  reverse reaction must release heat into the surroundings (- ∆H)



∙ Calorimetry: the measurement of heat changes that accompany chemical  and physical processes.

o Measured with a calorimeter; a closed, specially designed container ∙ Specific heat (s): the amount of heat required to raise the temperature of 1  gram of a substance by 1 ˚C.  

∙ Heat Capacity (C): the amount of heat needed to raise the temperature of  an object by 1 ˚C.

o The specific heat of a substance can be used to determine the heat  capacity of a specified amount of a substance

Heat Capacity and Specific Heat 

∙ The specific heat of a substance can be used to determine the heat  capacity of a specified amount of a substance

o For example, we can use the specific heat of water to determine  the heat capacity of a kilogram of water:

 heat capacity of 1 kg of water = 4 .184 J/1 g⋅°C × 1000 g  

= 4184  or 4.184 × 103 J/°C

∙ Specific heat can be used to determine the energy associated with a  temperature change

o q = sm∆T

o s = specific heat

o m = mass  

o ∆T = temperature change (Tf –Ti)

∙ Heat capacity can also be used:

o q = C∆T

 Where c is the heat capacity of the object

Constant Pressure Calorimetry 

∙ A constant pressure calorimeter made from two Styrofoam cups o the reaction mixture is insulated

o two solutions of known volume at the same temperature are mixed  together

∙ The heat absorbed or produced is determined by measuring the temperature chang

∙ The reactants and products constitute the system

∙ The water and the calorimeter are the surroundings

o qsys = -sm∆T

Hess’s Law 

∙ For many reactions, it is not possible to determine the enthalpy change  directly

∙ Hess’s law allows us to determine the enthalpy change for these reactions  o uses a series of chemical equations (steps)  

o the equations are arranged to sum the desired equation

∙ Equations may require manipulation

o multiply, divide, or reverse an equation

o follow guidelines for manipulating thermochemical equations (previous  lecture)

∙ Recall that enthalpy is a state function

o Hess’s Law states that the energy change for a reaction is the same  whether the reaction takes place in one step, or in a series of steps. o Example: the combustion of methane gas to produce gaseous water  Imagine the reaction occurring in two steps  

 CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) ∆H = -890.4  kJ/mol

o The liquid water produced is converted to gaseous water in the second  step

o 2H2O (l) → 2H2O (g) ∆H = +88.0  kJ/mol

∙ Identical terms on opposite sides of the equation can be canceled out.  ∙ The sum of the remaining equations must result in the desired chemical  equation

∙ ∆H values for each reaction are combined as well to give the overall enthalpy  for the reaction

∙ remember that any operations performed on the chemical equation must  also be performed on the value for ∆H

o CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (l) ∆H = -890.4 kJ/mol o 2H2O (l) → 2H2O (g) ∆H = +88.0 kJ/mol

o CH4 (g) + 2O2 (g) → CO2 (g) + 2H2O (g) ∆H = -802.4 kJ/mol Standard Enthalpies of Formation 

∙ So far, we have determined ΔH for a reaction by measuring the heat  absorbed or released by a reaction (at constant pressure).

∙ Recall that ΔH is the difference between the enthalpies of the products and  reactants

o ΔH = H(products) – H(reactants)

o Using this method requires that we know the enthalpies of the  reactants and products

∙ Standard enthalpy of formation (ΔHf˚) is defined as the heat change that  results when mole of a compound is formed from its constituent elements in  their standard state.

o the degree symbol, ˚, denotes the standard state: the most stable  state form of an element under standard conditions  

o although standard state does not specify a temperature, the ΔHf˚ values we use are measured at 25 ˚C.

o ΔHf˚ values are relative to this standard state

∙ A list of ΔHf˚ values for a number of compounds can be found in appendix 2 of your textbook. (you do not need to memorize these!)

∙ Using these known values for ΔHf˚ we can calculate ΔHrxn

o ΔHrxn = ΣnΔHf˚ (products) - ΣmΔHf˚ (reactants)

o Σ means “the sum of”

o n and m are the stoichiometric coefficients for the reactants and the  products respectively

o units for ΔHrxn stay in kJ/mol (the coefficients are unitless numbers; the  moles do not cancel out)

ΔHf ˚;Hess’s Law 

∙ Many compounds cannot be directly synthesized from their elements safely o some reactions proceed to slowly or produce other substances through  side reactions

o ΔHf˚ must be determined indirectly using Hess’s law.

 using a series of reactions for which ΔHrxn˚ can be measured

 equations are arranged so that the sum of the reactions gives  the equation for the formation of the compound of interest

The Nature of Light 

∙ Visible light is only a small part of the electromagnetic spectrum (a range  of electromagnetic radiation)

∙ Electromagnetic radiation (EMR): the transmission of energy through  waves

∙ The electromagnetic spectrum consists of the following types of EMR: o Gamma rays

o X-rays

o Ultraviolet  

o Visible

o Infrared

o Microwaves

o Radio waves

∙ The speed of a wave depends on two factors:

o The type of wave

o The medium through which the wave is traveling

∙ The speed of light (in a vacuum), c:

 c = 2.9998 × 108 m/s

∙ The speed of light, frequency, and wavelength are related through the  following equation

o c = λν

o λ = wavelength in meters (m) ν = frequency in reciprocal seconds (s 1) , known as Hertz (Hz)

o *it is customary to measure wavelength in units that correspond to the  magnitude of their wavelengths; for example, visible light is usually  measured in nanometers

Properties of Waves 

∙ Waves are characterized by the following properties:

o Wavelength λ (lambda): the distance between identical points on  successive waves

o Frequency ν (nu): the number of waves that through a point in 1  second

o Amplitude: the vertical distance from the midline of a wave to the top of the peak or the bottom of the through

∙ Wavelength and frequency are inversely proportional longer wavelength =  lower frequency

o c = λν

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