Podcast
Questions and Answers
What unit conversion factor would you use to convert 2011 lb to grams?
What unit conversion factor would you use to convert 2011 lb to grams?
Which of the following is an example of a chemical change?
Which of the following is an example of a chemical change?
Which statement accurately describes a characteristic of compounds?
Which statement accurately describes a characteristic of compounds?
What is the primary method for separating mixtures involving iron filings and sulfur powder?
What is the primary method for separating mixtures involving iron filings and sulfur powder?
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In the reaction where fuels in the space shuttle produce water, what are the reactants?
In the reaction where fuels in the space shuttle produce water, what are the reactants?
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What is the atomic weight of Boron calculated from its isotopes?
What is the atomic weight of Boron calculated from its isotopes?
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Which element forms a cation by losing one electron?
Which element forms a cation by losing one electron?
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What is the abundance of the 25Mg isotope in magnesium?
What is the abundance of the 25Mg isotope in magnesium?
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How many protons does the isotope 26Mg have?
How many protons does the isotope 26Mg have?
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What charge does sulfur typically exhibit when it gains two electrons?
What charge does sulfur typically exhibit when it gains two electrons?
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Which monatomic ion can have multiple oxidation states?
Which monatomic ion can have multiple oxidation states?
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What is the resulting charge of a phosphorus atom when it gains three electrons?
What is the resulting charge of a phosphorus atom when it gains three electrons?
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What type of ion is formed when metals typically lose electrons?
What type of ion is formed when metals typically lose electrons?
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Which of the following compounds uses Roman numerals in its name?
Which of the following compounds uses Roman numerals in its name?
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What is a characteristic property of ionic compounds?
What is a characteristic property of ionic compounds?
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Which formula represents a molecular compound?
Which formula represents a molecular compound?
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Which element exceptions does not require Roman numerals in its naming?
Which element exceptions does not require Roman numerals in its naming?
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What type of bond typically forms in ionic compounds?
What type of bond typically forms in ionic compounds?
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Which property of ionic compounds allows them to conduct electricity when molten?
Which property of ionic compounds allows them to conduct electricity when molten?
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What is the formula for iron(III) chloride?
What is the formula for iron(III) chloride?
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Which of the following is not a property of ionic compounds?
Which of the following is not a property of ionic compounds?
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What is the prefix for a compound containing three atoms of one element?
What is the prefix for a compound containing three atoms of one element?
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Which of the following compounds is correctly named using the prefixes for binary molecular compounds?
Which of the following compounds is correctly named using the prefixes for binary molecular compounds?
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What is the molecular formula for a branched alkane with five carbon atoms?
What is the molecular formula for a branched alkane with five carbon atoms?
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Which of the following is a common name for the compound with the formula N2O?
Which of the following is a common name for the compound with the formula N2O?
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In which type of hydrocarbon are all carbon atoms connected by single bonds?
In which type of hydrocarbon are all carbon atoms connected by single bonds?
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What is the prefix used for a hydrocarbon with eight carbon atoms?
What is the prefix used for a hydrocarbon with eight carbon atoms?
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Which of the following statements about isomers is true?
Which of the following statements about isomers is true?
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How can cyclic hydrocarbons be identified in their naming convention?
How can cyclic hydrocarbons be identified in their naming convention?
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What is the meaning of the atomic number (Z) for an element?
What is the meaning of the atomic number (Z) for an element?
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How is the mass number (A) calculated for an element?
How is the mass number (A) calculated for an element?
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Which statement about isotopes is TRUE?
Which statement about isotopes is TRUE?
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What is the purpose of a mass spectrometer?
What is the purpose of a mass spectrometer?
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What would be the number of neutrons in Magnesium given its atomic number is 12 and mass number is 25?
What would be the number of neutrons in Magnesium given its atomic number is 12 and mass number is 25?
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What is true about the value of 1 Unified atomic mass unit (u)?
What is true about the value of 1 Unified atomic mass unit (u)?
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How many copper atoms can fit across a diameter of 1.90 cm calculated in picometers?
How many copper atoms can fit across a diameter of 1.90 cm calculated in picometers?
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What denotes the element symbol and mass number format for carbon-12?
What denotes the element symbol and mass number format for carbon-12?
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What is the limiting reactant when mixing 25.0 mL of 0.234 M FeCl3 and 50.0 mL of 0.453 M NaOH?
What is the limiting reactant when mixing 25.0 mL of 0.234 M FeCl3 and 50.0 mL of 0.453 M NaOH?
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How many moles of Fe(OH)3 are produced when FeCl3 and NaOH are reacted?
How many moles of Fe(OH)3 are produced when FeCl3 and NaOH are reacted?
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What is the weight percent of oxalic acid in a 4.554 g mixture neutralized by 29.58 mL of 0.550 M NaOH?
What is the weight percent of oxalic acid in a 4.554 g mixture neutralized by 29.58 mL of 0.550 M NaOH?
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How many moles of NaOH are needed to neutralize one mole of oxalic acid?
How many moles of NaOH are needed to neutralize one mole of oxalic acid?
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What is the total volume used in the reaction of FeCl3 and NaOH?
What is the total volume used in the reaction of FeCl3 and NaOH?
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If 0.01627 mol of NaOH was used, how many moles of oxalic acid does that correspond to?
If 0.01627 mol of NaOH was used, how many moles of oxalic acid does that correspond to?
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What is the molar mass of oxalic acid (H2C2O4)?
What is the molar mass of oxalic acid (H2C2O4)?
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What does increasing the concentration of H+ ions in water indicate?
What does increasing the concentration of H+ ions in water indicate?
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Study Notes
Chapter 1: The Nature of Chemistry
- Chemistry is the science of matter and its transformations.
- Studying chemistry helps us understand our surroundings and how we function.
- Chemistry is central to medicine, engineering, and many other sciences.
Identifying Matter: Physical Properties
- Physical properties can be measured without changing the composition of a substance.
- Examples include temperature, pressure, mass, volume, melting point, boiling point, density, color, state (solid, liquid, or gas), and shape of crystals.
Physical Change
- A physical change does not alter the substance's composition.
- Examples include changes in state (e.g., ice melting to liquid water), shape (e.g., hammering lead into a sheet), and size (e.g., cutting a piece of wood in half).
Melting & Boiling Point
- Temperature (T) measures the relative energy content of an object.
- Energy transfers from higher-temperature objects to lower-temperature objects.
- Common temperature scales include Fahrenheit (°F) and Celsius (°C).
- Water freezes at 32°F (0°C) and boils at 212°F (100°C).
- Conversion formulas between Fahrenheit and Celsius are provided.
Density
- Density is a physical property defined as mass divided by volume.
- The density of a substance is a measure of how much mass is contained in a given volume.
- Different substances have different densities.
- The units of density are typically grams per milliliter (g/mL) or grams per cubic centimeter (g/cm³).
- Values for the density of various substances are provided.
Measurements, Units & Calculations
- Scientists use SI units (Système International d'unités).
- Key SI units and their symbols for length, mass, time, and temperature are listed.
SI Units
- Prefixes are used to multiply or divide units by multiples of ten.
- Examples and meanings of common SI prefixes (mega, kilo, deci, centi, milli, micro, nano, pico) are included.
Significant Digits & Rounding Numbers
- Measurements involve uncertainty.
- Significant figures represent the reliable digits in a measurement.
- Rules for determining significant figures and rounding are provided.
Significant Figures in Calculations
- Rules for rounding answers in addition, subtraction, multiplication, and division calculations involving significant figures.
Rules for Rounding
- Rules for rounding numbers to a specified number of significant figures are provided.
Classifying Matter: Elements & Compounds
- Elements cannot be broken down into simpler substances by chemical means.
- Compounds are formed when elements combine in fixed ratios.
- Sucrose is an example of a compound with a specific composition.
- Compounds have specific properties, such as melting and boiling points.
- Examples of elements and compounds frequently mention in chemistry are provided .
- Water is a compound that consistently melts at a specific temperature.
Types of Matter
- Matter is anything that occupies space and has mass.
- Matter can be categorized as heterogeneous or homogeneous substances.
- Substances are further classified as compounds or elements.
- Solutions are homogeneous mixtures of two or more substances.
Nanoscale Theories & Models
- Macroscale objects are large enough to be seen, measured, and handled without aids.
- Microscale objects require a microscope to view them.
- Nanoscale objects have dimensions comparable to atoms.
States of Matter: Solids, Liquids & Gases
- The kinetic-molecular theory describes matter as composed of tiny particles in constant motion.
- Solids have closely packed particles in regular arrays with fixed locations, vibrating in place.
- Liquids have slightly more open structures, with larger volumes and more randomly arranged particles, allowing them to flow.
- Gases have widely spaced particles in constant, rapid motion with no fixed volume or shape.
The Atomic Theory
- All matter is composed of atoms.
- Atoms of the same element have identical chemical properties.
- Compounds are formed when atoms of different elements combine in fixed proportions.
- During chemical reactions, atoms join, separate, or rearrange but are not created or destroyed.
The Chemical Elements
- Elements have unique names and symbols.
- Some symbols are obvious abbreviations, derived from their corresponding element names.
- Symbols for common elements, such as Helium, Hydrogen, Titanium, Gold, Silver, Tin, and Lead are shown.
- The origin of some element names is included.
The Chemical Discovery Elements
- Information about the discovery and origin of names of important elements are provided .
Types of Elements
- Over 110 elements are currently recognized.
- 90 occur naturally on Earth.
- Most elements are metals.
- Metals are typically solids, conductive of electricity and heat, and malleable (able to be bent or hammered into sheets).
- Nonmetals can exist as solids, liquids, or gases, and are typically poor conductors of electricity and heat.
- Metalloids have properties of both metals and nonmetals.
Periodic Table
- The periodic table organizes elements based on their properties.
- Rows are called periods and columns are called groups.
- Elements within the same group exhibit similar properties.
Elements that Consist of Molecules
- Many nonmetal elements exist as molecules.
- Diatomic molecules include H₂, O₂, N₂, F₂, Cl₂, Br₂, and I₂.
- Polyatomic molecules include molecules with more than two atoms like O3, P4, and S8.
Communicating Chemistry: Symbolism
- Chemical formulas show the number and type of atoms in a molecule.
- Relative proportions of atoms in a compound, such as sucrose (C₁₂H₂₂O₁₁) and methanol (CH₃OH) are represented by chemical formulas.
- Chemical equations show how reactants convert into products(e.g., sucrose decomposing to carbon and water).
Biological Periodic Table
- Some elements are more prevalent in living organisms than others.
- Major elements like hydrogen, oxygen, carbon, nitrogen, calcium, phosphorus, sulfur, sodium, potassium, chlorine, and magnesium are important building blocks of life.
- Trace elements are also vital to biological processes.
Chapter 2: Chemical Compounds
- Atoms are made of electrons, protons, and neutrons.
- Radioactivity was discovered using uranium ore.
- Experimentation demonstrated three types of rays: alpha (α), beta (β), and gamma (γ).
- Atoms contain subatomic units
Atomic Structure & Subatomic Particles
- Atoms are made up of subatomic particles.
- Electrons (e⁻) are negatively charged.
- Protons (p⁺) are positively charged.
- Neutrons (n⁰) are neutral.
Radioactivity
- Uranium ore emits rays that can fog photographic plates.
- Marie and Pierre Curie isolated new elements, polonium (Po), and radium (Ra).
- Radioactivity is the term for the phenomenon of an element undergoing transformations that result in rays such as the ones produced by uranium and polonium.
Radioactivity (Three Distinct Types of Radiation)
- Alpha particles are positively charged.
- Beta particles are negatively charged.
- Gamma rays have no charge.
- The different types of rays are deflected by an electric field in different directions.
Electrons
- Thomson discovered electrons using cathode ray tubes.
- Electrons have a negative charge and very small mass.
- Millikan measured the charge of individual electrons.
Electrons (Millikan 1911)
- Millikan experimented using electrically charged oil droplets.
- Individual electron charges were measured using a precisely controlled electric field.
- Data obtained from these experiments were consistent with the assumption that an electron's charge is a multiple of some smallest, elementary unit of charge.
Protons
- Atoms gain a positive charge when electrons are lost; this suggests the existence of a positive subatomic particle.
- Hydrogen ions, called protons (P⁺), have the lowest mass, as they consist of a single proton.
- Protons have a positive charge equal in magnitude to the electron's negative charge.
The Nuclear Atom
- J. J. Thomson proposed the "plum pudding" model, picturing the atom as a sphere of positive charge with negatively charged electrons embedded within.
- Rutherford's gold foil experiment demonstrated a concentrated positive charge at the center, called the nucleus.
Neutrons
- Rutherford's model explained most of the mass and positive charge concentrated in a small core called the nucleus.
- Chadwick discovered neutrons as neutral particles in the nucleus.
- Neutrons have a similar mass to that of protons.
The Nuclear Atom
- Atoms consist of a nucleus (containing protons (p+ ) and neutrons (n0)) and electrons (e⁻) orbiting the nucleus.
- The nucleus accounts for the mass of an atom.
- Electrons occupy the space surrounding the nucleus.
- A neutral atom has equal numbers of protons and electrons, therefore equal positive and negative charges.
Atomic Numbers & Mass Numbers
- Definition of atomic number (Z), mass number (A).
- Examples of calculating the number of protons (p⁺), neutrons (n⁰), and electrons for specific atomic symbols.
Isotopes & Average Atomic Mass
- Isotopes are atoms of the same element that have the same number of protons but different numbers of neutrons.
- The atomic mass of an element is the weighted average of the masses of its isotopes.
- The mass of the atom is a precise amount. A precise value represents the weight of the single atom.
Mass Spectrometer
- Mass spectrometers measure the mass-to-charge ratio (m/z) of charged atoms or molecules.
- A mass spectrometer helps determine the mass to charge ratio. This is helpful to distinguish between isotopes.
- The area under the peaks on a mass spectrum represents the relative abundance of different isotopes.
Isotopes & Atomic Weight
- Most elements exist as a mixture of isotopes.
- Atoms of the same element differ in their masses due to the different numbers of neutrons.
- The atomic weight of an element reflects the weighted average of all naturally occurring isotopes of that element.
Average Atomic Mass
- The weighted average of multiple masses is called the average atomic mass.
Ions & Ionic Compounds
- Ions are atoms or groups of atoms with a positive or negative charge.
- Metals tend to form cations (positive ions).
- Non-metals tend to form anions (negative ions).
- A neutral atom formed ions in an ionic compound because of the transfer of electrons between elements.
Monatomic Ions
- Add or lose electrons in the process of forming cations or anions.
- The charge of an ion is equal to the group number minus eight.
- Ions are commonly formed so that they have the same number of electrons as the nearest noble gas.
Monatomic Ions (Transition Metals)
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Transition metals can have multiple possible charges.
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The old and new group numbers for transition metals do not correspond to the charge values.
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The charge states from transition metals are shown with Roman numerals in parenthesis.
Polyatomic Ions
- Polyatomic ions are groups of atoms with a net electrical charge.
- Examples like ammonium (NH⁺₄), hydroxide (OH⁻), sulfate (SO₄²⁻), hydrogen sulfate (HSO₄-), cyanide (CN⁻) and nitrate (NO₃⁻) are shown.
Oxoanions
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Oxoanions are polyatomic ions containing oxygen and another element, like phosphate (PO₄³⁻), sulfate (SO₄²⁻), nitrate (NO₃⁻), phosphite (PO₃³-), sulfite (SO₃²⁻) and so on.
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A series of related names (e.g., perchlorate, chlorate, chlorite, hypochlorite) for a family of oxoanions is shown.
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If names begin with "H", like "hydrogen sulfate", the name is modified to include "hydrogen" for naming purposes.
Ionic Compounds
- Ionic compounds are held together by electrostatic forces between positively charged cations and negatively charged anions, are always electrically neutral.
Recognizing Ionic Compounds
- Ionic compounds contain a metal and a nonmetal, or a metal and a polyatomic ion.
Naming Ions & Ionic Compounds
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Positive metal ions (cations), except for ammonium (NH⁺₄), are named using the metal's name.
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Positive ions with more than one charge have their charge state shown with a Roman numeral.
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Negative ions (anions) are named by adding “-ide” to the element's root name.
Naming Ionic Compounds
- Name the cation followed by the anion, dropping “ion” from the names.
Ionic Compound Properties
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Ionic compounds are typically solids at room temperature, characterized by high melting and boiling points, hardness, and brittleness.
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Poor conductors of heat and electricity when solid, but good when molten.
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Properties of ionic compounds indicate that the ions are held tightly in fixed positions; therefore, they also possess strong electrostatic attractions.
The Crystal Lattice
- Ionic compounds exist as crystal lattices, with each ion surrounded by multiple ions of opposite charge.
- A cubic structure in a repeating pattern is typical of many ionic compounds.
Properties of Ionic Compounds
- The structure of ionic crystals explains their properties, such as hardness and brittleness.
Molecular Compounds
- Molecular compounds are composed of nonmetal atoms.
- Molecules are neutral units that are formed when nonmetals bond in covalent interactions.
Molecular Formulas
- Molecular formulas show the type and number of atoms in a molecule—but not the arrangement of atoms.
- Condensed formulas provide similar information but in a more compact form.
- Functional groups are specific groups of atoms within a molecule that impart characteristic chemical behaviors.
Naming Binary Molecular Compounds
- If one element is hydrogen, it is named first, then the second element is named with the "-ide" suffix.
- If the elements are not hydrogen: the first element is named according to the number of atoms present, and the second element name has the "-ide" suffix following the prefix that reflects the number of atoms.
Hydrocarbons
- Hydrocarbons are binary molecules containing carbon and hydrogen.
- Alkanes are hydrocarbons with only single bonds between carbon atoms.
- Alkanes are traditionally named by adding the suffix "ane" to the prefix reflecting the number of carbons present.
Isomers
- Isomers are compounds with the same molecular formula but different structural arrangements of atoms.
- Branched alkanes are isomers of their linear counterparts.
Alkanes & Their Isomers
- Alkyl functional groups result from removing a hydrogen atom from an alkane.
- The removed hydrogen is replaced with a functional group that reflects the characteristic chemical behavior.
- Rules for naming alkyl functional groups are provided in the text.
- Examples of alkyl groups are shown.
- The possible isomers using alkane structures and their corresponding names are presented.
Amount of Substance: The Mole
- The mole (mol) is a unit in chemistry that measures the amount of a substance, similar to a "dozen".
- The mass of a mole of an element is numerically equal to that element’s atomic weight in grams.
- Avogadro's number is the number of entities (atoms, molecules, ions, electrons, etc.) in a mole.
Gram-Mole Calculations
- Procedures for determining the number of moles of an element or compound are provided, given the mass of the element or compound and the molar mass of the element or compound.
Moles of Compounds
- A mole of a compound contains a certain number of moles of atoms from the elements that are in the compound.
- The molar mass of a compound is the sum of the atomic masses of the atoms in the compound.
- Examples of specific calculations using moles and molar mass are provided using the formulas for the substances presented in the text.
Molar Mass of Ionic Compounds
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Formula weight is the sum of the atomic weights of all atoms in a formula unit.
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Molar Mass of Ionic Compound is the mass of one mole of formula units, and is numerically equivalent to its formula mass. Molar Mass (g/mol) is numerically equal to the formula weight.
Ionic Hydrates
- Ionic hydrates are ionic compounds with water molecules trapped within their crystal structure.
- Names are formed by adding a Greek prefix indicating the number of water molecules to the anhydrous salt name.
Gram-Mole Calculations.
- Calculations for determining the number of moles present in a compound given a mass and the molar mass or vice versa.
Composition & Chemical Formula
- Percentage composition by mass: calculating the percentage of each element in a compound from its chemical formula, or vice versa.
Empirical & Molecular Formulas
- Empirical formula: represents the simplest whole-number ratio of atoms in a compound.
Empirical & Molecular Formulas
- Determining the empirical formula of a compound from its elemental percentage composition, given a mass of the compound.
Empirical & Molecular Formulas
- Calculate the empirical formula and molecular formula of a compound given percentage composition and molar mass
Solutions Concentration
- Solutions are mixtures of solute and solvent.
- Molarity is a key concentration unit; moles of solute/litres of solution.
Molarity
- Calculation of the molarity of a solution is shown; given the mass of solute and the volume of solution.
Molarity
- Calculation of the molarity of a solution; given the moles of solute and the volume of the solution; including calculations of the concentrations for the constituents of the solution.
Solution Preparation
- Preparing a solution by dissolving a measured amount of solute and diluting to a fixed volume is addressed.
- Preparing a solution by diluting a more concentrated solution is addressed
Solution Preparation from Pure Solute
- Calculations for preparing a solution from a pure substance; given the desired molarity, and volume of solution.
- Steps to prepare the solution are addressed
Solution Preparation by Dilution
- Calculation of the final concentration of a diluted solution, given the initial concentration and volume of the concentrated solution, and the final volume of the diluted solution.
Stoichiometry in Aqueous Solution
- Calculations for determining how many moles of a specific product are produced, given moles of the reactant(s), using stoichiometric ratios from a chemical reaction, using the moles of product to calculate mass.
Molarity & Reactions in Aqueous Solution
- Calculating the moles of a product given the concentration and volume of a solution participating in a chemical reaction..
Stoichiometry in Aqueous Solution
- Calculations of weight percent of component in a mixture of chemicals given the volume of reactant(s), concentration of reactant(s), and moles of the desired product.
Acids
- Acids increase the concentration of H+ ions in water.
- Strong acids completely ionize in water.
- Weak acids partially ionize in water.
Bases
- Bases increase the concentration of OH⁻ ions in water.
- Strong bases completely ionize in water.
- Weak bases partially ionize in water.
Common Acids & Bases
- List of strong and weak acids and bases commonly found in chemical reactions is given.
Neutralization Reactions
- Neutralization reactions occur when an acid reacts with a base to form a salt and water.
- Chemical reactions are typically exchange reactions. The positive ion exchanges from the acid to combine with the negative ion in the base.
Net Ionic Equations for Acid-Base Reactions
- Net ionic equation for reactions involving strong acids and strong bases.
- Net ionic equation for reactions involving weak acids and strong bases.
- Equations are simplified so only the ions responsible for the reaction are shown.
Titrations in Aqueous Solution
- Titration is a process for determining the concentration of an unknown solution by reacting it with a solution of known concentration.
Titrations in Aqueous Solution
- Calculation for determining the concentration of a solution from neutralization of acid with base from known volume and concentration, or vice-versa.
Chapter 4: Energy and Chemical Reactions (Part 1)
- Energy is the capacity to do work.
- Work occurs when an object moves against a resisting force.
- All energy is either potential or kinetic energy.
The Nature of Energy
- Kinetic energy is the energy of motion.
- Potential energy is the energy of position.
- Energy units include the joule (J), calorie (cal), and dietary calorie (Cal).
- Conversion between units often involves a specific factor, such as 1 cal = 4.184 J.
Conservation of Energy
- Energy can neither be created nor destroyed, only transformed between forms.
- The total energy of the universe is constant.
Energy & Working
- Work is done when an object moves against a force.
- Energy is transferred and stored by different forms like potential and kinetic energy including chemical energy.
Energy, Temperature & Heating
- Temperature is a measure of the average kinetic energy of particles in a substance.
- Heat is the transfer of energy from a hotter object to a cooler object.
- Heat capacity and specific heat capacity describe the thermal properties of a substance.
Keeping Track of Energy Transfers
- System is the part of the universe under study.
- Surroundings include the rest of the universe.
- Internal energy is the sum of all energy (kinetic and potential) of the particles within the system.
Calculating Thermodynamic Changes
- Calculating changes in the internal energy (ΔE) of a system from initial and final energies.
Conservation of Energy & Reactions
- Energy changes in chemical reactions are tracked using the system concept.
- The conservation of energy is shown in a balanced chemical equation where the total energy input must equal the sum of the energy output.
Energy Changes (Heat Transfer)
- Calculations for determining change of energy using energy transfers that can involve heat transfer and work performed on, by, or within the system.
Heat Capacity
- Heat capacity is the energy required to raise the temperature of a substance by 1°C.
- Different substances have different specific heats or molar heat capacities, describing energy transfer rates.
Heat Capacity
- Calculations for determining the heat needed to raise the temperature of a sample or vice versa.
- Calculating the energy required to change the temperature of solids or liquids, given mass, specific heat, initial temperature and final temperature.
Conservation of Energy & Changes of State
- During changes of state (e.g., melting, freezing, vaporization, condensation), energy is transferred but temperature remains constant.
Conservation of Energy & Changes of State (Heat of fusion)
- The enthalpy of fusion is the energy required (or released) to melt (or freeze) a substance at a fixed temperature and pressure.
Conservation of Energy & Changes of State (Heat of vaporization)
- Similarly, enthalpy of vaporization is the energy required (or released) to vaporize (or condense) a substance at a fixed temperature and pressure.
State Functions & Path Independence
- State functions have values that depend only on the current state of a system, regardless of how the system arrived at that state.
- Examples like Enthalpy (H), Internal Energy (E) and Volume (V), which depend only on the current state of the system, are shown.
- Changes in enthalpy (∆H) for a system remain constant independent of the pathway used.
Reaction Enthalpies for Chemical Reactions
- ∆H is the heat absorbed or released during a chemical reaction at constant temperature and pressure.
- Given that the total internal energy of the system remains the same during a chemical reaction, we can calculate the amount of heat released or absorbed in kJ/mol.
- Calculating the amount of heat release or absorbed, (qp ), using the moles, and the molar enthalpy.
Where Does the Energy Come From? (Bond Enthalpies)
- Bond enthalpy is the energy required to break a bond, with the energy released when a bond is formed.
Bond Enthalpies (Thermodynamic & Kinetic Stability)
- Calculating how much energy is released or absorbed in a chemical reaction given the bond enthalpies of the bonds formed or broken. The sum of all the energies required to break all bonds during the chemical reaction is mathematically compared to the sum of all the energies released when forming new bonds during the chemical reaction.
Measuring Reaction Enthalpies
- Calorimetry is used to measure heat transferred during chemical reactions.
- Types of calorimeters include bomb calorimeters (constant volume, for reactions in a combustion environment) and coffee-cup calorimeters (constant pressure).
- Using calorimetry, heat capacity, enthalpy or internal energy can be calculated, given the change of temperature of the substance, and conditions present during the chemical reaction, like temperature and pressure
Constant Volume Calorimetry
- Calculations involving bomb calorimetry are presented and used to measure energy released during a chemical reaction.
Constant Volume Calorimetry
- Calculations using bond enthalpies to calculate the energy released during a chemical reaction, involving the use of a bomb calorimeter and the energy change from the temperature readings are shown.
Constant Volume Calorimetry
- Calculating heat released by the reaction, given the conditions and equations of the reactions are shown.
Constant Volume Calorimetry
- Calculations for determining energy change during a chemical reaction given the conditions (temperature changes and mass) using a coffee-cup calorimeter. The heat capacity is also considered during the calculation.
Hess's Law
- If a reaction is the sum of multiple intermediate reactions, then enthalpy change of the overall reaction equals the sum of the enthalpy changes of the intermediate reactions; also, the enthalpy change for a reaction remains the same regardless of pathway.
Hess's Law
- Calculating the enthalpy change of a chemical reaction using Hess's law.
Standard Formation Enthalpies
- The standard enthalpy of formation (∆,H°) is the enthalpy change when 1 mole of a substance is formed from its elements in their standard states at a standard temperature (25°C ) and pressure ( 1 bar ).
Standard Formation Enthalpies
- Calculated using Hess's law, where the standard enthalpy of formation of an element (given it is in its most stable state at 25°C and 1 bar) is zero.
Standard Formation Enthalpies
- Finding the standard enthalpy change of a specific chemical reaction by considering the changes in the standard enthalpies of formation for each component involved in the reaction.
Fuels for Society & Our Bodies
- Chemical fuels react exothermiccally with O, in air (oxygen in the atmosphere).
- Calculating the relative energy production per mole based on bond enthalpy provided for chemical fuels.
Fuels for Society & Our Bodies
- Carbohydrates, and other fuels, are broken-down in the body, to produce energy (chemical energy).
- Energy is measured using kJ/ g or Cal/ g.
- Energy consumption rates and fuel values are presented for food.
Fuels for Society & Our Bodies (Fossil fuels)
- The major sources for energy worldwide are fossil fuels.
- Data on the amounts of energy produced by each energy source historically and presently include various charts depicting the changes in fossil fuel consumption over the years.
Petroleum
- Petroleum is a complex mixture of various hydrocarbons.
- Refining is important in separating the mixtures into specific fractions based on their boiling points.
Fractional Distillation
- Separate components of a mixture based on differing boiling points.
Petroleum Refining
- Methods for refining crude oil into various fractions with different boiling points.
- These are fractionated based on their boiling points; consequently, they have useful properties depending on their components.
Octane Number
- Octane ratings are used in fuel, based on how well compounds with similar structures will burn smoothy in a combustion environment.
Catalytic Cracking
- Breaking down large hydrocarbon molecules into smaller hydrocarbon molecules with higher octane numbers using heat and catalysts.
Catalytic Reforming
- Transforming less desirable straight-chain hydrocarbons into branched or aromatic hydrocarbons.
- This improves fuel performance by raising the octane number of the fuel.
Octane Enhancers
- Additives to increase octane ratings, often containing lead in the past, but now often using compounds containing oxygen.
Oxygenated & Reformulated Gasoline
- Oxygen is often added because it promotes complete combustion and reduces emissions.
U.S. Energy Sources & Consumption
- Total energy production (in joules), by source, historically and in a recent year, along with percentage of energy source relative to production from other sources.
Natural Gas
- Natural gas is a mixture mainly of methane (CH₄).
- The increased production of natural gas by fracking in U.S. supplies is described and compared to other sources of energy.
- Potential environmental effects and concerns are addressed
Coal
- Coal is a fuel derived from fossilized plant matter, often containing sulfur (S).
Chemical Reactions & Dispersal of Energy
- The concept of dispersal of energy as relevant to probability-driven chemical reactions is described and why.
- Most exothermic reactions are product-favored, as energy increases in the surroundings.
Probability & Dispersal of Energy
- Predicting probability of energy dispersal when multiple atoms are in contact or interacting.
Dispersal of Matter
- Explaining how matter can disperse due to expansion in a gas or liquid.
Measuring Dispersal of E: Entropy
- Calculating the change in entropy (∆S°) for a reaction at a given temperature.
- Using standard molar entropies of reactants and products to estimate the change in entropy caused by a reaction.
- Qualitative guidelines to predict the direction of change in entropy given the specific reaction.
Qualitative Guidelines for Entropy
- Entropy and molecular complexity are related; increasingly complex molecules can have increased entropy.
Qualitative Guidelines for Entropy (Ionic Solids)
- Predicting the enthalpy change of ionic solids using the physical forces that maintain the solids' structural formation.
- The structural forces can also indicate how extensive is the motion of the solid's constituents.
Qualitative Guidelines for Entropy (Solid or Liquid Dissolving)
- Calculating the change in entropy caused by dissolving a specific compound, whether the solute is a solid, liquid or gas, using values for standard entropies of pure and dissolved compounds.
Predicting Entropy Changes (Gas Dissovling)
- Calculating the change in entropy of dissolving specific gases in various solvents.
- Predicting whether dissolving a gas will increase or decrease the entropy. Using standard entropy data of gases to calculate the change in entropy from dissolving them in a solvent or solution.
Calculating Entropy Changes
- Calculating the entropy change of a chemical reaction.
Entropy & the 2nd Law of Thermodynamics
- Understanding entropy and its relationship to the second law of thermodynamics.
Gibbs Free Energy
- Combining enthalpy (H) and entropy (S) to determine Gibbs free energy (G).
- Predicting whether a reaction is product-favored from a positive or negative Gibbs energy.
Effect of T on Reaction Direction
- Examining how temperature affects the spontaneity and direction of a reaction (reactant-favored or product-favored), based on whether the reaction is exothermic or endothermic.
Gibbs Free Energy (Constant T)
- Calculating the Gibbs free energy given the reaction enthalpy and entropy.
Gibbs Free Energy & Maximum Work
- Gibbs free energy is the maximum possible useful work from a reaction.
Thermodynamic & Kinetic Stability
- Categorizing chemical processes as either thermodynamically or kinetically stable and unstable.
End of Study Notes
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Test your knowledge on fundamental chemistry concepts with this quiz. Questions cover unit conversions, types of changes in matter, and methods of separation in mixtures. Perfect for high school chemistry students!
