AP Chem- CONCEPT MAP

Postulate 1

• All elements are composed of indivisible particles called atoms

Postulate 1

Postulate 2

• Atoms of the same element are identical
• The atoms of any one element are different from those of another

Postulate 2

Postulate 3

• Atoms of different elements mix or combine in whole number ratios

Postulate 3

Postulate 4

• Chemical reactions occur when atoms separate, join, or rearrange
• In a chemical reaction, atoms of one element can never change into another substance

Postulate 4

Cathode Ray Experiment

• In the tube was an inert gas and two plates, a positive and negative
• Particles in the gas were attracted to the positive plate
• Therefore, the particles must have a negative charge

Oil Drop Experiment

Oil Drop Experiment

• Put a charge on a tiny drop of oil
• Measured how strong an electric field had to be to stop the oil drop from falling
• The mass of an electron is 9.10 × 10 -28 grams
• The charge was -1

Observation Of Gold Foil Experiment

• Most of the alpha particles went through he concluded...

Observation

Observation Of Gold Foil Experiment

• Very rarely particles were deflected at large angles he concluded...

Observation

Observation Of Gold Foil Experiment

Observation

• Few paticles were deflected at small angles he concluded...

Experiment Of Discovery The Neutron

Experiment

• Bombared a thin sheet of beryllium with alpha paticles
• A high radiation energy that is emitted by the metal
• Later experiments demonstrated that rays consistes of subatomic particles which are neutrons
• proved that the electrical neutral particles have aa mass slightly bigger than the mass of a proton

Average Atomic Mass

Average Atomic Mass

• Sometimes called atomic weight
• No difference between atomic mass and atomic weight

• The mass of the atom in atomic mass unit (amu)

Atomic Mass Unit

• Defined as a mass exactly equal to one-twelfth the mass of one carbon-12 atom

Atomic Mass Unit

Isotope

• Atoms having the same atomic number but different mass number

Isotope

Example

Video Explaining the Difference

Video

• What's the Difference between Mass Number and Atomic Mass?
• https://www.youtube.com/watch?v=m15DWkkGe_0

Example Calculation

Example

Video on how to calculate the average atomic mass

• How to Calculate Atomic Mass Practice Problems
• https://www.youtube.com/watch?v=ULRsJYhQmlo

Video

Mole

• Is the amount of a substance that contains as many elementary entities (moles, molecules, or other particles) as there are atoms in exactly 12 g (or 0.012 kg) of the carbon 12 isotope

Molar Mass

• the mass (in grams or kilograms) of 1 mole of units (such as atoms or molecules) of a substance

Molar Mass

Avogadro Number

• The number of particles in a mole

• 6.022 x 10^23 atoms, molecules, or formula units
• 602,000,000,000,000,000,000,000
• Six-Hundred and two sextillion

Molecular Mass

• Sometimes called molecular weight
• Is the sum of the atomic masses (amu) in the molecule
• Molar mass of a compound ( in grams) is numerically equal to the molecular mass (in amu)

Formula Mass

• also known as formula weight
• Is the sum of the atomic weights of the atoms in the empirical formula of the compound.
• Formula weight is given in atomic mass units (amu)

Example of Molecular Mass of H20

Example

• 2(atomic mass of H) + atomic mass of O

Or

• 2(1.008 amu) + 16.00 amu = 18.02 amu

Video On Calculating Molecular Mass

• How to Calculate Molar Mass (Molecular Weight)
• https://www.youtube.com/watch?v=o3MMBO8WxjY

Video

Mass Spectrometer

• Developed in the 1920s by the English physicist- F.W. Aston
• Analytical laboratory technique to separate the components of a sample by their mass and electrical charge
• Produces a mass spectrum that plots the mass-to-charge (m/z) ratio of compounds in a mixture.

Link To Image

• https://cdn1.byjus.com/chemistry/2018/02/12114729/Mass-spectrometry-diagram.jpg

Link To Image

Step 1- Ionization

• The atom is ionized by knocking one or more electrons off to give a positive ion.
• This is true even for things which you would normally expect to form negative ions (chlorine, for example) or never form ions at all (argon, for example).
• Mass spectrometers always work with positive ions.

Step 1

Step 2- Acceleration

• The ions are accelerated so that they all have the same kinetic energy.

Step 2

Step 3- Deflection

Step 3

• The ions are then deflected by a magnetic field according to their masses.
• The lighter they are, the more they are deflected.
• The amount of deflection also depends on the number of positive charges on the ion
• On how many electrons were knocked off in the first stage.
• The more the ion is charged, the more it gets deflected.

Step 4- Detection

• The beam of ions passing through the machine is detected electrically

Step 4

Percent Composition

Percent Composition

By mass

• The percent by mass of each element in a compound

By Mole

• Percent by mole of each element in a compound

Step 1

• Find Atomic Mass Of An Element

Step 1

Step 2

• Find Molar Mass Of Certified Atom

Step 2

Step 3

• Multiply Molar Mass And Number Of Moles Of That Element

Step 3

Step 5

• Multiply By 100

Step 5

Step 4

• Divide Each Atom Mass By The Molecular Mass

Step 4

Link To Image

https://image.slideserve.com/671162/percent-composition-l.jpg

Link To Image

Molecular Formula

• An expression showing the exact number of atoms of each element in a molecule

• Example: C6 H10 O4

Empirical Formula

• An expression showing the types of elements present and the simplest ratios of the different kinds of atoms

• Example: CH

Step 1

• Find Percent Composition

Step 1

Step 2

• Assume 100 grams

Step 2

Step 3

• Convert To Moles

Step 3

Step 4

• Divide all moles by the lowest number

Step 4

Step 5

• Multiply By A Number To Make All Mole Ratio A Whole Number If Necessary

Step 5

Step 6

• Put The Answers As Subscripts Of The Element They Correspond Too

Step 6

Step 1

• Find the number of moles of each element in a sample of the molecule

Step 1

Step 2

• Find the ratios between the number of moles of each element

Step 2

Step 3

• Find the empirical formula

Step 3

Step 4

• Find the molecular weight of the empirical formula

Step 4

Step 5

• Find the number of empirical formula units in the molecular formula

Step 5

Step 6

• Find the molecular formula

Step 6

Example of Difference Between Formulas

Example

Video On The Difference Of Formulas

Video

• What are Chemical Formulas? What is the difference between Empirical & Molecular Formulas?
• https://www.youtube.com/watch?v=5ZtcByIUo8U

Chemical Reaction

• A process in which a substance (or substances) is changes into one or more new substances

Chemical Reaction

Chemical Equation

• Uses chemical symbols to show what happens during a chemical reaction

Chemical Equation

Example of Chemical Equation

Example

Step 1

• Writing the chemical formula

Step 1

Step 2

• List all the individual atoms on each side of the yield sign

Step 2

Step 3

• Change the coefficients on the the different molecules and make the atoms equal

Step 3

Step 4

• Make an educated guess and check until you get the answer

Step 4

Stoichiometry

• Is the quantitative study of reactants and products in a chemical reaction

Stoichiometry

Mole Method

• The stoichiometric coefficients in a chemicl reaction can be interpreted as the number of moles of each substance

Mole Method

Step 1

Writing the chemical equation

Step 1

Step 2

Convert everything to moles

Step 2

Step 3

Use the mole ratio to calculate the moles of a substane

Step 3

Step 4

Convert out of moles to the desired unit

Stoichiometric

Amounts

• In the proportions indicated by the balanced equation

Stoichiometric

Amounts

Limiting Reagent

• The reactant used up first in a reaction

Limiting Reagent

Excess Reagent

• The reagents present in quantities greater than necessary to react with the quantity of the limit

Excess Reagent

Example of Excess Reagent

Example

Videos On How To Find Excess Reagent

Video

• Limiting and Excess Reactant Stoichiometry Chemistry Practice Problems Tutorial
• https://www.youtube.com/watch?v=Q4ojTIOO-HY

• Calculating the Amount of Excess Reactant in a Limiting Reagent Problem
• https://www.youtube.com/watch?v=yO_8nViq55E

Actual Yield

• The amount of product actually obtained from a reaction

Actual Yield

Theoretical Yield

• The amount of product that wild result if all the limiting reagent reacted

Theoretical

Yield

Percent Yield

• The proportion of the actual yield to the theoretical yield

Percent Yield

Exist at 25° and 1 atm

Exist at 25° and

1 atm

Do Not Exist at 25° and 1 atm

Physical Characteristics #1

• Gases assume the volume and shape of their container

Number 1

Physical Characteristics #2

• Gases are the most compressible

Number 2

Physical Characteristics #3

• Gases will mix evenly and completely when confined to the same container

Number 3

Physical Characteristics #4

• Gases have much lower densities than liquids and solids

Number 4

Pascal (Pa)

• A pressure of one newton per square meter

Pascal

Pressure

• Forced applied per unit are

Pressure

Atmospheric Pressure

• The pressure exerted by Earth's atmosphere

Atmospheric Pressure

Standard Atmospheric Pressure (1 atm)

Standard Atmospheric Pressure

• Equal to the pressure that supports a column of mercury exactly 760 mm or 76 cm high at 0°C at sea level

• 1 atm = 760 mmHg

Barometer

• An instrument for measuring atmospheric pressure

Barometer

Torr

• a unit of pressure used in measuring partial vacuums

• 1 Torr= 1 mmHg

Torr

Manometer

• A device used to measure the pressure of gases other than the atmosphere

Manometer

Velocity

• Velocity = Distance Moved / Elapsed Time

Velocity

Acceleration

• Acceleration = Change in Velocity / Elapsed Time

Acceleration

SI Unit of Force (Newton)

• 1 N = 1 kg Meters Per Second Squared

SI Unit of Force

Kelvin

• SI base unit of temperature

• One Kelvin is equal in magnitude in one degree in Celsius

Kelvin

Absolute Zero

• Theoretically the lowest attainable temperature
• -273.15 °C

Absolute Zero

Proportionally Constant

• PV = K1
• K1 = Proportionally Constant

Proportionally Constant

Boyle's Law

• The pressure of the a fixed amount of gas at a constant temperature is inversely proportional to the volume of the gas

• P1V1 = P2V2 or PV = K1
• P = Pressure
• V= Volume
• K1 = Constant

Charles's Law (Charles's and Gay-Lussac's Law)

Charles's Law

• The volume of a fixed amount of gas maintained at constant pressure is directly proportional to the absolute temperature of the gas

• V Form
• V1/T1 = V2/T2
• V = Volume
• T= Temperature
• P Form
• P1/T1 = P2/T2
• P = Pressure
• T = Temperature

Avogradro's Law

• At constant pressure and temperature, the volume of a gas is directly proportional to the number of moles of the gas present

• V1/N1 = V2/N2 or V = K4n
• V = Volume
• N= Quantity of Gas in Moles
• K4= Constant

Gas Constant (R)

Gas Constant

• The constant that appears in the ideal gas equation

• Expressed as...
• 0.08206 L x atm / K x mol
• 8.314 J/ K x mol

Ideal Gas

• A hypothetical gas whose pressure-volume-temperature behavior can be completely accounted for by the ideal gas equation

Ideal Gas

Ideal Gas Equation

• An equation expressing the relationships among pressure, volume, temperature, and amount of gas

• Equation:
• PV = nRT
• P = Pressure
• V = Volume
• T = Temperature
• n = Number of Molecules
• R = Gas Constant

Standard Temperature and Pressure (STP)

• Conditions 0°C and 1 atm

STP

Calculate R

Calculate R

• R Form (Gas Constant)
• R= PV / nT
• = (1 atm) (22.414 L)/ (1 mol) (273.15 K)

• Temperature scale
• 0°C or 273.15 Kelvin

• Molar Volume of Gas at STP
• 1 mole = 22.41 Liter or 0.082057 Liter

• Calculation of Molar Volume of Gas
• 0.082057 L x atm/ K x mol

Molar Mass

Molar Mass

• Equation
• M= dRT/ P
• d = Density
• R = Gas Constant
• T = Temperature
• P = Pressure

Density

Density

• Equation
• d= m/V = PM/RT
• m = Mass
• V= Volume
• P = Pressure
• M = Molar Mass
• R = Gas Constant
• T = Temperature

Video

• Ideal gas equation: PV = nRT

Additional Video

Partial Pressure

• The pressure of individual gas components in the mixture

Partial Pressure

Dalton Law of Partial Pressure

• The total pressure of a mixture of gases is just the sum of the pressures that each gas would exert if it were present alone

• Equation
• PT = P1 + P2 + P3...
• P1 + P2 + P3... = the Partial Pressures of Component
• T = Total

Mole Faction

• A dimensionless quantity that expresses the ration of the number of moles of one component to the number of moles of all components present

Kinetic Energy (KE)

• The type of energy expended by a moving object, or energy of motion

Kinetic Energy

Kinetic Molecular Theory of Gases

• A number of generalizations about gas behavior

• Formula
• KE = 0.5 x m x u squared
• m = Mass of Molecule
• U = Speed

Root-Mean-Square (RMS)

Root-Mean - Square

• An average molecular speed

• Formula
• (U^2)½ = Urms = (3RT/M)½
• Urms = root mean square velocity in m/sec
• R = Gas Constant
• T= Temperature
• M = Mass of Mole

Diffusion

• The gradual mixing of molecules of one gas with molecule of another by virtue of their kinetic properties

Diffusion

Effusion

• The process by which one gas gradually escapes from one compartment of a container to another by passing through a small opening

Assumption #1

• A gas is composed of molecules that are separated from each other by distances far greater than their own dimensions
• The molecules can be considered to be "points"
• They posses mass but have negligible volume

Assumption #2

• Gas molecules are in constant motion in random directions, and frequently collide with one another
• Collisions among molecules are elastic

Assumption #3

• Gas molecules exert neither attractive nor repulsive forces on one another

Assumption #3

Assumption #4

• The average kinetic energy of the molecules is proportional to the temperature of the gas in kelvins
• Average kinetic energy of molecules is given by KE = 1/2 m u squared

Video

The Kinetic Molecular Theory of Gas (part 1 and 2)

Video

Graph

Example

• Graham's Law of Effusion Practice Problems, Examples, and Formula

Example

Video

• Graham's Law of Diffusion

Video

Video / Example

• Van der Waals equation

Video / Example

Hydration

• As ionic solids dissolve in water (break up into individual anions and cations), the positive ends of the water attract the anions and the negative ends of the water attract the cations

Solubility

• The maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature

Solubility

Reversible Reaction

• The reaction can occur in both directions

Reversible Reaction

Water

• One of the most significant substances on Earth

• Can dissolve many different substances

• A polar molecule because of its unequal charge distribution

• Is “bent” or V-shaped
• 104.5°C between hydrogens

• Is a polar molecule because the oxygen atom has a greater attraction for electrons

Solubility

Solubility

Water As A Solvent

Video #1

Solution

• A homogeneous mixture made up of a solute and solvent

Solution

Solute

• Substance being dissolved

Solute

Solvent

• The substance doing the dissolving

Solvent

Aqueous Solution

• Solutions in which water is the solvent

Aqueous Solution

Electrolyte

• Electrical conductivity is the ability of a solution to conduct electricity

Electrolyte

Strong Electrolyte

• These are completely ionized (completely dissociate into separate ions) when dissolved in water

Strong Electrolyte

Weak Electrolyte

• These dissociate only a little, while the majority of the substance does not

Weak Electrolyte

Nonelectrolyte

• These dissolve in water but do not produce ions

Nonelectrolyte

Classification Of Solutes in Aqueous Solutions

Classification Of Solutes

Aqueous Solutions, Acids, Bases and Salts

Video #1

Identifying Strong Electrolytes, Weak Electrolytes, and Nonelectrolytes

Video #2

Solubility Chemistry

Video #3

Water & Solutions - for Dirty Laundry: Crash Course Chemistry #7

Video #4

Molarity

• Moles of solute per volume of solution in liters

Molarity

Concentration of Solutions

• The amount of solute present in a given amount of solvent, or a given amount of solution

Concentration of Solutions

Dilution

• The procedure for preparing a less concentrated solution from a more concentrated one

Dilution

Standard Solution

• Solution whose concentration is accurately known

Standard Solution

Stock Solution

• Solutions in concentrated forms used for dilutions of different concentration

Stock Solution

Example

• For a 0.25 M CaCl2 solution:

CaCl2 → Ca2+ + 2Cl–

• Ca2+: 1 × 0.25 M = 0.25 M Ca2+
• Cl–: 2 × 0.25 M = 0.50 M Cl–

Example

• Calculate the concentration of Cl- ions if 0.70mol of ZnCl2 are dissolved in 1.75L of solution.

• When ZnCl2 dissolves: ZnCl2 ​​ Zn+2 + 2Cl
• 0.70 mole x 2Cl =0.80 M Cl
• 1.75L sol-

Molality Practice Problems

Video #1

Solutions: Crash Course Chemistry #27

Video #2

Molarity Practice Problems

Video #3

Synthesis

• Elements combine with other elements to form a compound

Synthesis

Decomposition

• Compounds break down into elements and/or smaller compounds

Decomposition

Replacement Reactions

Combustion

• A reaction in which a substance reacts with oxygen, usually with the release of heat and light to produce a flame

Combustion

Combination #1

• NH4+ + OH- produces NH3 + H2O

• ammonium chloride + sodium hydroxide ​ ​ammonia + water
• sodium chloride

Combination #1

Combination #2

• H+ + CO3-2 produces CO2 + H2O

• hydrochloric acid + calcium carbonate ​​carbon dioxide + water
• calcium chloride

Combination #2

Combination #3

• H+ + SO3-2 produces SO2 + H2O

• acetic acid + lithium sulfite ​​sulfur dioxide + water + lithiu
• acetate

Combination #3

Combination #4

• H+ + S-2 produces H2S

• phosphoric acid + silver sulfide ​​ hydrogen sulfide + silve
• phosphate

Combination #4

Chemical Reactions

Video #1

t

Video #2

t

Video #3

Precipitation Reactions

• A reaction that results in the formation of a precipite

Precipitation Reactions

Precipitate

• An insoluble solid that separates from the solutio

Precipitate

Soluble

• Solid dissolves in solution;
• (aq) is used in reaction

Soluble

Insoluble

• Solid does not dissolve in solution
• (s) is used in reaction

Insoluble

Precipitate

• Insoluble and slightly soluble are often used interchangeably

Precipitate

Precipitation Reactions - Explained

Video #1

Precipitation Reactions: Crash Course Chemistry #9

Video #2

Precipitation Reactions and Net Ionic Equations

Video #3

Molecular Equation

• The formulas of the compounds are written as though all species existed as molecules or whole units

Molecular Equation

Ionic Equation

• Shows dissolved species as free ions

• Represents as ions all reactants and products that are strong electrolytes

Net Ionic Equation

• Includes only those solution components undergoing a change

• Show only components that actually react

Spectator ions

• Are not included
• Ions that do not participate directly in the reaction
• Na+ and NO3-

Spectator Ions

Example

• AgNO3(aq) + NaCl(aq) → AgCl(s) + NaNO3(aq)

Example

Step #1

• Write a balanced molecular equation for the reaction

• Using the correct formulas for the reactant and product ionic compounds

• Refer to the insoluble products and therefore will appear as a precipitate

Step #2

• Write the ionic equations for the reaction

• The compound that does not appear as the precipitate should be shown as free ions

Step #2

Step #3

• Identify and cancel the spectator ions on both sides of the equation

• Write the net ionic equation for the reaction

Step #3

Step #4

• Check that the charges and number of atoms balance in the net ionic equation

Step #4

Molecular, Ionic, and Net Ionic Equations

Video #1

How to Write Complete Ionic Equations and Net Ionic Equations

Video #2

Complete Ionic and Net Ionic Equations

Video #3

Writing Molecular, Total & Net Ionic Equations

Video #4

Step #1

• Identify the species present in the combined solution, and determine what reaction if any occurs

Step #1

Step #2

• Write the balanced net ionic equation for the reaction

Step #2

Step #3

• Calculate the moles of reactants

Step #3

Step #4

• Determine which reactant is limiting

Step #4

Step #5

• Calculate the moles of product(s), as required

Step #5

Step #7

• Convert to grams or other units, as required

Step #6

Solution Stoichiometry Practice Problems & Examples -

Video #1

Stoichiometry of a Precipitation Reaction

Video #2

Precipitation Stoichiometry

Video #3

Acid–Base Reactions

• A type of chemical reaction that involves the exchange of one or more hydrogen ions, H+, between species that may be neutral (molecules, such as water, H2O) or electrically charged (ions, such as ammonium, NH4+; hydroxide, OH−; or carbonate, CO32−).

Neutralization Reaction

• A reaction between an acid and a base

Neutralization Reaction

Hydronium Ion

• The hydrated proton, H30+

Hydronium Ion

Salt

• An ionic compound made up of a cation other than H + and an anion ther than OH- or O2-

Salt

Acid–Base Titrations

Calculations

Calculating Acid–Base Reactions

Acids

Acid

Base

Step #1

• The exact reaction with titrant and analyte has to be known

Step #1

Step #2

• The equivalence point must be marked accurately

Step #2

Step #3

• The volume of titrant needed to react equivalence point must be known accurately

Step #3

Acid Base Neutralization Reactions

Video #1

Acid Base Titration Problems

Video #2

Combustion

• A reaction in which a substance reacts with oxygen, usually with the release of heat and light to produce a flame

Combustion

Combination Reaction

• A reaction in which two or more substances combine to form a single product

Combination Reaction

Decomposition

• Compounds break down into elements and/or smaller compounds

Decomposition

Displacement Reactions

• An atom or an ion in a compound is replaced by an atom of another element

Displacement Reactions

Disproportionation Reaction

• t

Disproportionation Reaction

Redox Reaction

• Reactions in which one or more electrons are transferred

Oxidation

Oxidation

Rules

Rules For Assigning Oxidation Numbers

Introduction to Oxidation Reduction (Redox) Reactions

Video #1

Oxidation and Reduction Reactions - Basic Introduction

Video #2

Redox Reactions: Crash Course Chemistry #10

Video #3

Oxidation and Reduction (Redox) Reactions Step-by-Step Example

Video #4

How to Calculate Oxidation Numbers Introduction

Video #5

Half-Reaction

• Explicitly shows the electrons involved in a redox reaction

Half- Reaction

Oxidation–Reduction

Reactions by Half Reactions

Half Reactions

• Oxidation states are assigned to each element
• Elements being oxidized or reduced are identified
• Then separate the reaction into two half-reactions:
• one involving oxidation and the other dealing with reduction
• The method for balancing half-reactions varies depending on whether the reactions are taking place in acidic or basic (alkaline) solution

Acidic Solution

The Half-Reaction Method for Redox Reactions in Acidic Solution

Basic Solution

The Half-Reaction Method for Redox Reactions in Basic Solution

How to Balance Redox Equations in Acidic Solution

Video #1

How to Balance Redox Equations in Basic Solution

Video #2

Energy

• The capacity to do work or to produce heat

Energy

Work

• Directed energy change resulting from a process

Work

Process

Heat

• Transfer of energy between two bodies that are at different temperatures
• Heat flows from hotter to cooler

Heat

Surroundings

• The rest of the universe outside a system

Surroundings

System

• Any specific part of the universe that is of interest to us

System

Thermochemistry

• The study of heat changes in chemical reactions

Thermochemistry

State Function

• Property that does not depend in any way on the system’s past or future (only depends on present state)

State Function

Position

Energy

Energy

• On any given pathway, the total energy change is always constant, but the work and heat expended will differ
• Energy is a state function

Formula

• Work = -PΔV
• P is the external pressure
• ΔV is the change in volume

Formula

Work

• For an expanding gas, ΔV is a positive quantity because the volume is increasing
• To convert between L·atm and Joules, use 1 L·atm = 101.3 J.

Energy & Chemistry: Crash Course Chemistry #17

Video #1

5.1 Nature of Energy (Chemistry)

Video #2

First Law of Thermodynamics, Basic Introduction - Internal Energy, Heat and Work - Chemistry

Video #3

Thermochemistry Equations & Formulas - Lecture Review & Practice Problems

Video #4

Internal Energy, Heat, and Work Thermodynamics, Pressure & Volume, Chemistry Problems

Video #5

Enthalpy

• A thermodynamic quantity used to describe heat changes taking place at constant pressure

• Enthalpy is a state function

Calorimetry

• The measurement of heat changes

Calorimetry

Heat Capacity

• The energy needed to raise a specific mass of an object by one degree

Heat Capacity

Endothermic

• If two reactants at the same temperature are mixed and the resulting solution gets warmer, this means the reaction taking place

Endothermic

Exothermic

• A reaction that cools the solution

Exothermic

Change In Enthalpy

Part 1

• State function

• ΔH = q at constant pressure
• ΔH = Hproducts – Hreactants

• If ΔH is positive: endothermic reaction

• If ΔH is negative: exothermic reaction

Change In Enthalpy

Part 2

• For a particular reaction, ΔH values are linked to the coefficients in the balanced equation
• 1 CH4 + 2 O2 ​​ 1 CO2 + 2 H2O ΔH = -890

• When 1 mole of CH4 reacts with 2 moles of O2t
• produce 1 mole of CO2 and 2 moles of H2O, then 890 k
• of energy is produce

• If any other quantity of reactants is used, a differet
• amount of energy is evolve

• Use dimensional analysis to calculate these amount.

Example

• 1 CH4 + 2 O2 ​​ 1 CO2 + 2 H2O ΔH = -890 k

• How much energy is produced from the complet rreaction of 2.56 moles of methane (CH4)?
• 2.56 moles CH x -890 k = -2280 k

1 mole CHJ

Example

Calorimetry

Formula

Formula

“Coffee Cup” Calorimeter: Constant-Pressure Calorimeter

“Coffee Cup” Calorimeter

• The thermometer records temperature change as the chemicals react in the water

• The temperature change is then converted into units of energy

Bomb Calorimeter:

Constant-Volume Calorimeter

Bomb Calorimeter

• A bomb calorimeter is often used to determine the heat of combustion for a combustible material

• Weighed reactants are placed inside a steel container and ignited

• The energy change is measured by the temperature change of the surrounding water

• Use q = Ccal × ΔT

• Ccal = the heat capacity

Calorimetry Problems

Video #1

Bomb Calorimeter vs Coffee Cup Calorimeter Problem

Video #2

Coffee Cup Calorimeter

Video #3

Enthalpy: Crash Course Chemistry #18

Video #4

Calorimetry: Crash Course Chemistry #19

Video #5

Intensive Property

• A property that does not depend on how much matter is being considered

Intensive Property

Extensive Property

• A property that depends on how much matter is being considered

Extensive Property

Thermochemical Equations

• The enthalpy changes as well as the mass relationships

Thermochemical Equations

Problem-Solving Strategy

Characteristics #1

• If a reaction is reversed, the sign of ΔH is also reversed

Characteristics #1

Characteristics #2

• The magnitude of ΔH is directly proportional to the quantities of reactants and products in a reaction

• If the coefficients in a balanced reaction are multiplied by an integer, the value of ΔH

Hess Law Chemistry Problems

Video #1

Hess's Law Problems & Enthalpy Change

Video #2

Thermochemical Equations Practice Problems

Video #3

Hess's law and reaction enthalpy change

Video #4

Hess's law example

Video #5

Standard Enthalpy of Formation (ΔHf°)

• Change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states.

Standard Enthalpy of Formation (ΔHf°)

Standard State

• The condition of 1 atm of pressure

Standard State

Lattice Energy

• The energy required to completely separate one mole of a sold ionic compound into gaseous ions

Lattice Energy

Heat

Step #1

• When a reaction is reversed, the magnitude of ΔH remains the same, but its sign changes

Step #1

Step #2

• When the balanced equation for a reaction is multiplied by an integer, the value of ΔH for that reaction must be multiplied by the same integer

Step #3

• The change in enthalpy for a given reaction can be calculated from the enthalpies of formation of the reactants and products:
• ΔH°rxn = ΣnpΔHf°(products) - ΣnrΔHf°(reactants)

Step #4

• Elements in their standard states are not included in the ΔHreaction calculations because ΔHf° for an element in its standard state is zero

Step #4

Enthalpy of Formation Reaction & Heat of Combustion

Video #1

Thermochemical Equations Practice Problems

Video #2