Chem Exam #2

Ch. 4, 5, 10

Characteristics of Gases

Low density, nonmetalic elements, low molar mass, expand to fill containers, highly compressible, form homogenous mixtures

Temperature (T)

Pressure (P)

Amount of (F) per area (A) P=(F/A)

Volume (V)

Amount of space occupied by moles

Atmospheric pressure

Weight of air per unit of area

SI Unit is kPa

Ways to measure pressure

Mercury barometer, manometer

Units of Pressure

Pascals, Bar, mm Hg, Torricelli, atomsphere (atm)

Standard Pressure

Normal atmospheric pressure at sea level

1 atm

Boyle’s Law

Volume and pressure are related

Volume of fixed quantity of gas at constant temperature is inversely proportional to the pressure

As volume decreases, pressure goes up, and vice versa

PV = constant

Charles’s Law

Volume and temperature are related

As temp increases, volume increases

V = Constant x T

Gay-Lussac’s Law of Combining Volumes

At a given temperature and pressure, the volume of gases that react with each other is in small whole numbers

At constant volume, pressure is directly proportional to temperature

Avogadro’s Law

The volume of a gas at constant temperature and pressure is directly proportional to the number of moles of the gas

Combined Gas Law

Combines Boyle’s Charles’s, and Gay Lussac’s laws

Ideal Gas Equation

PV = nRT

Ideal gas is a hypothetical gas whose pressure, volume, and temperature is constant

uses a constant proportionality and reorganize

R

The ideal gas constant makes the equation and equality, not only a proportion

PV = nRT

Density of Gases (d)

d = (M x P) / (R x T)

Mole Fraction

Ratio of moles of a substance to total moles

Kinetic-Molecular Theory

1) Gases consist of large numbers of molecules that are in continuous, random motion

2) Combined volume of all molecules of gas is negligible relative to the total volume in which the gas is contained

3) Attractive and repulsive forces between gas molecules are negligible

4) Avg kinetic energy of gas in a container does not change, but can individually transfer energy between each other

5) Avg kinetic energy of the molecules is proportional to the absolute temperature

U_mp

Most probable speed related to size of molecules

U_av

Avg speed of the molecules

U_rms

the speed of a molecule whose kinetic energy is equal to the average (mean) kinetic energy of all the molecules

Effusion

Escape of gas molecules through a tiny hole into an evacuated space

Diffusion

•Diffusion is the spread of one substance throughout a space or a second substance.

Graham’s Law Describes Diffusion and Effusion

•Graham’s law relates the molar mass of two gases to their rate of speed of travel.

The raete of effusion of a gas is inversely proportioned to the dquare root of its molecular weight

•The “lighter” gas always has a faster rate of speed.

Real Gases

•In the real world, the behavior of gases conforms to the ideal-gas equation only at relatively high temperature and low pressure.

•Even the same gas will show wildly different behavior under high pressure at different temperatures.

Van der Waals equation

•One equation corrected for real gases is known as the van der Waals equation.

Takes care of errors encountered in previous equations

•The pressure adjustment is due to the fact that molecules attract and repel each other.

•The volume adjustment is due to the fact that molecules occupy some space on their own.

thermodynamics

•This chapter is about thermodynamics, which is the study of energy and its transformations.

thermochemistry

•Specifically, thermochemistry is the study of chemical reactions and the energy changes that involve heat.

energy

•From Chapter 1, energy is the ability to do work or transfer heat.

Potential energy

•The most important form of potential energy in molecules is electrostatic

Joule

the unit of energy commonly used is the Joule:

•Electrostatic attraction is seen between _____ charged ions.

oppositely

When is energy released?

•Energy is released when chemical bonds are formed

When is energy consumed?

energy is consumed when chemical bonds are broken.

First Law of Thermodynamics

•Energy can be converted from one form to another, but it is neither created nor destroyed.

system and surroundings

•The portion of the universe that we single out to study is called the system (here, the hydrogen and oxygen molecules). •The surroundings are everything else (here, the cylinder, piston and everything beyond).

1)Open System:

1)Open System: a region of the universe being studied that can exchange heat and mass with its surroundings.

Closed System:

1)Closed System: a region of the universe being studied that can only exchange heat with its surroundings (not mass)

1)Isolated System

1)Isolated System: a region of the universe that can not exchange heat or mass with its surroundings

internal energy

•The internal energy of a system is the sum of all kinetic and potential energies of all components of the system; we use E to represent it.

q

system gains heat from surroundings

system loses head to surroundings

w

work done on system by surroundings

work done by system on surroundings

delta E

net gain of E by system

net loss of E by system

endothermic

When heat is absorbed by the system from the surroundings, the process is endothermic.

Delta H is positive

exothermic

When heat is released by the system into the surroundings, the process is exothermic.

Delta H is negative

State functions

E is a state function

It depends only on the present state of the system, not on the path by which the system arrived at that state

q and w are not state functions

Enthalpy

•Enthalpy is the internal energy plus the product of pressure and volume

•If a process takes place at constant pressure (and we usually work at atmospheric pressure) and the only work done is this pressure–volume work, we can account for heat flow during the process by measuring the enthalpy (H) of the system.

•So, at constant pressure, the change in enthalpy is the heat gained or lost.

Delta H

Change in enthalpy

Enthalpy of products minus enthalpy of reactants

Truth about Enthalpy

1.Enthalpy is an extensive property.

2.The enthalpy change for a reaction is equal in magnitude, but opposite in sign, to delta H for the reverse reaction

3.The enthalpy change for a reaction depends on the states of the reactants and the products.

calorimetry

•Since we cannot know the exact enthalpy of the reactants and products, we measure delta H through calorimetry, the measurement of heat flow.

•The instrument used to measure heat flow is called a calorimeter.

heat capacity

The amount of energy required to raise the temperature of a substance by 1 K (1 degree Celsius) is its heat capacity.

Molar heat capacity

The amount of energy required to raise the temperature of a substance by 1 K (1 degree Celsius) is its heat capacity. If the amount of the substance heated is one gram, it is the specific heat. If the amount is one mole, it is the molar heat capacity.

Constant Pressure Calorimetry

•By carrying out a reaction in aqueous solution in a simple calorimeter, the heat change for the system can be found by measuring the heat change for the water in the calorimeter.

Bomb Calorimetry

•Reactions can be carried out in a sealed “bomb” such as this one.

•Because the volume in the bomb calorimeter is constant, what is measured is really the change in internal energy,

Delta E

Hess’s Law

•Hess’s law: If a reaction is carried out in a series of steps, Delta H for the overall reaction equals the sum of the enthalpy changes for the individual steps.

An enthalpy of formation,

An enthalpy of formation, Delta H, is defined as the enthalpy change for the reaction in which a compound is made from its constituent elements in their elemental forms.

Standard Enthalpies of Formation

By definition, the standard enthalpy of formation of the most stable form of any element is zero. Thus, the values of  for C(graphite), and the standard states of other elements are zero by definition.

Bond Enthalpy

The enthalpy associated with breaking one mole of a particular bond in a gaseous substance.

•The bond enthalpy is always positive because energy is required to break chemical bonds.

•Energy is always released when a bond forms between gaseous fragments.

•The greater the bond enthalpy, the stronger the bond.

fuel value.

•The energy released when one gram of food is combusted is its fuel value.

The vast majority of the energy consumed in this country comes from _____

fossil fuels.

Renewable energy sources,

•Renewable energy sources, like solar, wind, geothermal, hydroelectric, and biomass sources produce 9.9% of the U.S. energy needs.

Solutions

•Solutions are defined as homogeneous mixtures of two or more pure substances.

solvent

•The solvent is present in greatest abundance.

Solutes

Smallest amount

aqueous solution.

•When water is the solvent, the solution is called an aqueous solution.

dissociation

–Ionic compounds dissolve by dissociation, where water surrounds the separated ions.

–Molecular compounds interact with water, but most do NOT dissociate.

Some molecular substances react with water when they dissolve.

solvation

▪All substances dissolve by solvation, surrounding of the solute by solvent.

electrolyte

•An electrolyte is a substance that dissociates into ions when dissolved in water.

nonelectrolyte

•A nonelectrolyte may dissolve in water, but it does not dissociate into ions when it does so.

strong electrolyte

•A strong electrolyte dissociates completely when dissolved in water.

•A strong electrolyte dissociates completely when dissolved in water. Its equations for the reaction of its dissociation look familiar:

weak electrolyte

•A weak electrolyte only dissociates partially when dissolved in water.

•A weak electrolyte only dissociates partially when dissolved in water. Its equations indicate a chemical equilibrium, where a reaction goes both forward and backward:

nonelectrolyte

•A nonelectrolyte does NOT dissociate in water.

•Precipitation reactions

•Precipitation reactions occur when two solutions containing soluble salts are mixed and an insoluble salt is produced. The solid is called a precipitate.

Metathesis (Exchange) Reactions

•It appears as though the ions in the reactant compounds exchange, or transpose, ions, as seen in the given equation.

Molecular Equation

•The molecular equation lists the reactants and products without indicating the ionic nature of the compounds. The only indication of the overall reaction is the states of matter given.

Complete Ionic Equation

•In the complete ionic equation all strong electrolytes (strong acids, strong bases, and soluble ionic salts) are dissociated into their ions.

Net Ionic Equation

•To form the net ionic equation, cross out anything that does not change from the left side of the equation to the right.

spectator ions

•The ions crossed out are called spectator ions in net ionic equations

Acids

•Acids as substances that ionize in aqueous solution to form H+

Bases

•Bases are substances that react with, or accept, H+

neutralization reactions

•Reactions between an acid and a base are called neutralization reactions.

Salt

ionic compound

oxidation

•Loss of electrons is oxidation.

reduction

•Gain of electrons is reduction.

Displacement Reactions

•In displacement reactions, ions oxidize an element.

Metal/Acid Displacement Reactions

•Elements higher on the activity series are more reactive.

•They will exist as ions.

•The element below will exist as the element.

Molarity

•The quantity of solute in a solution can matter to a chemist.

•Molarity is one way to measure the concentration of a solution:

concentration

•We call the amount dissolved its concentration.

diluted

A solution can be diluted by adding ONLY solvent. The concentration is LOWER, but the MOLES don’t change.

titration

A titration is an analytical technique in which one can calculate the concentration of a solute in a solution.

standard solution

•A solution of known concentration, called a standard solution, is used to determine the unknown concentration of another solution.