Ch 100: Fundamentals for Chemistry Chapter 1: Introduction Lecture Notes

What is Chemistry? Chemistry is considered to be the central science Chemistry is the study of matter Matter is the “stuff” that makes up the universe The fundamental questions of Chemistry are: How can matter be described? How does one type of matter interact with other types of matter? How does matter transform into other forms of matter?

Scientific Method 1. Recognize a problem Make observation Ask a question 2. Make an educated guess - a hypothesis Predict the consequences of the hypothesis 3. Perform experiments to test the predictions Does experiment support or dispute hypothesis? 4. Formulate the simplest rule that organizes the 3 main ingredients - develop a theory

The Scientific Attitude All hypotheses must be testable (i.e. there must be a way to prove them wrong!!) Scientific: “Matter is made up of tiny particles called atoms” Non-Scientific: “There are tiny particles of matter in the universe that will never be detected”

Major Developments in Chemistry I ~400 BC: Democritus proposed the concept of the “atom” ~300 BC: Aristotle developed 1st comprehensive model of matter ~700 AD: Chinese alchemists invent gunpowder 1661: Robert Boyle proposed the concept of elements 1770-1790: Lavoisier proposed the concept of compounds & the Law of Mass Conservation 1774: Priestly isolates oxygen 1797: Proust proposed the Law of Definite Proportions 1803: Dalton re-introduces the concept of the atom and establishes Dalton’s Laws 1869: Mendeleev creates the 1st Periodic Table 1910: Rutherford proposes the “nuclear” model of the atom 1915: Bohr proposes a “planetary” model of the hydrogen atom 1920: Schroedinger publishes his wave equation for hydrogen 1969: Murray Gell-Mann proposes the theory of QCD (proposing the existence of quarks)

Major Developments in Chemistry II Discovery of subatomic particles: 1886: Proton (first observed by Eugene Goldstein) 1897: Electron (JJ Thompson) 1920: Proton (named by Ernest Rutherford) 1932: Neutron (James Chadwick) Other Important Discoveries: 1896: Antoine Henri Becquerel discovers radioactivity 1911: H. Kamerlingh Onnes discovers superconductivity in low temperature mercury 1947: William Shockley and colleagues invent the first transistor 1996: Cornell, Wieman, and Ketterle observe the 5th state of matter (the Bose-Einstein condensate) in the laboratory

Ch 100: Fundamentals for Chemistry Chapter 2: Measurements & Calculations Lecture Notes

Types of Observations Qualitative Quantitative Descriptive/subjective in nature Detail qualities such as color, taste, etc. Example: “It is really warm outside today” Quantitative Described by a number and a unit (an accepted reference scale) Also known as measurements Example: “The temperature is 85oF outside today”

Measurements Described with a value (number) & a unit (reference scale) Both the value and unit are of equal importance!! The value indicates a measurement’s size (based on its unit) The unit indicates a measurement’s relationship to other physical quantities

Scientific Notation Technique Used to Express Very Large or Very Small Numbers Based on Powers of 10 To Compare Numbers Written in Scientific Notation First Compare Exponents of 10 (order of magnitude) Then Compare Numbers 3

Writing Numbers in Scientific Notation Locate the Decimal Point Move the decimal point to the right of the non-zero digit in the largest place The new number is now between 1 and 10 Multiply the new number by 10n where n is the number of places you moved the decimal point Determine the sign on the exponent, n If the decimal point was moved left, n is + If the decimal point was moved right, n is – If the decimal point was not moved, n is 0 4

Writing Numbers in Standard Form Determine the sign of n of 10n If n is + the decimal point will move to the right If n is – the decimal point will move to the left Determine the value of the exponent of 10 Tells the number of places to move the decimal point Move the decimal point and rewrite the number 5

Measurement Systems There are 3 standard unit systems we will focus on: 1. United States Customary System (USCS) formerly the British system of measurement Used in US, Albania, and a couple others Base units are defined but seem arbitrary (e.g. there are 12 inches in 1 foot) 2. Metric Used by most countries Developed in France during Napoleon’s reign Units are related by powers of 10 (e.g. there are 1000 meters in 1 kilometer) 3. SI (L’Systeme Internationale) a special set of metric units Used by scientists and most science textbooks Not always the most practical unit system for lab work

Related Units in the Metric System All units in the metric system are related to the fundamental unit by a power of 10 The power of 10 is indicated by a prefix The prefixes are always the same, regardless of the fundamental unit 6

Units & Measurement When a measurement has a specific unit (i.e. 25 cm) it can can be expressed using different units without changing its meaning Example: 25 cm is the same as 0.25 m or even 250 mm The choice of unit is somewhat arbitrary, what is important is the observation it represents

Measurement & Uncertainty A measurement always has some amount of uncertainty Uncertainty comes from limitations of the techniques used for comparison To understand how reliable a measurement is, we need to understand the limitations of the measurement 10

Measurements & Significant Figures To indicate the uncertainty of a single measurement scientists use a system called significant figures The last digit written in a measurement is the number that is considered to be uncertain Unless stated otherwise, the uncertainty in the last digit is ±1 11

Rules for Counting Significant Figures Nonzero integers are always significant Zeros Leading zeros never count as significant figures Captive zeros are always significant Trailing zeros are significant if the number has a decimal point Exact numbers have an unlimited number of significant figures 12

Rules for Rounding Off If the digit to be removed is less than 5, the preceding digit stays the same is equal to or greater than 5, the preceding digit is increased by 1 In a series of calculations, carry the extra digits to the final result and then round off Don’t forget to add place-holding zeros if necessary to keep value the same!! 13

Exact Numbers Exact Numbers are numbers known with certainty Unlimited number of significant figures They are either counting numbers number of sides on a square or defined 100 cm = 1 m, 12 in = 1 ft, 1 in = 2.54 cm 1 kg = 1000 g, 1 LB = 16 oz 1000 mL = 1 L; 1 gal = 4 qts. 1 minute = 60 seconds 14

Converting between Unit Systems To convert from one unit to another: Identify the relationship between the units (e.g. 100 cm = 1 m) Write out the starting measurement and multiply it by a quantity that will yield the desired value: 25 cm ( ) = _____ m The number in the “( )” is called the “conversion factor”

Metric Prefixes

Weight vs. Mass Mass is the amount of “stuff” in an object Mass is inertia Mass is the same everywhere in the universe SI Units of mass are kilograms (kg) Weight is the effect of gravity on an object’s mass Weight is a force Weight depends on location SI units of weight are newtons (N) USCS units are pounds (lb)

Volume The 3-D space an object occupies The SI unit is m3 (meters x meters x meters) The common metric unit is the Liter (L) Mass and volume are not the same thing Do not confuse mass & volume

Density Density is a property of matter representing the mass per unit volume For equal volumes, denser object has larger mass For equal masses, denser object has small volume Solids = g/cm3 1 cm3 = 1 mL Liquids = g/mL Gases = g/L Volume of a solid can be determined by water displacement Density : solids > liquids >>> gases In a heterogeneous mixture, denser object sinks 23

Using Density in Calculations 24

Ch 100: Fundamentals for Chemistry Chapter 3: Matter & Energy Lecture Notes

Aristotle (384-322 BC) Introduced observation as an important step in understanding the natural world All types of matter are mixtures of one of 4 basic “elements”: All matter has one or more of 4 basic “qualities”: According to Aristotle: Any substance could be transformed into another substance by altering the relative proportion of these qualities (i.e. lead to gold) 4) Fire 2) Water 3) Air 1) Earth 4) Dry 2) Moist 3) Hot 1) Cold

Physical & Chemical Properties Physical Properties are the characteristics of matter that can be changed without changing its composition Characteristics that are directly observable Chemical Properties are the characteristics that determine how the composition of matter changes as a result of contact with other matter or the influence of energy Characteristics that describe the behavior of matter 4

Physical & Chemical Changes Physical Changes are changes to matter that do not result in a change the fundamental components that make that substance State Changes : boiling, melting, condensing Chemical Changes involve a change in the fundamental components of the substance Produce a new substance Chemical reaction Reactants  Products 8

States of Matter Solid → Liquid → Gas Solid ← Liquid ← Gas +Energy 7

Classification of Matter Matter can be classified as either Pure or Impure: Pure Element: composed of only one type of atom Composed of either individual atoms or molecules (e.g. O2) Compound: composed of more than one type of atom Consists of molecules Impure (or mixture) Homogeneous: uniform throughout, appears to be one thing pure substances solutions (single phase homogeneous mixtures) Suspensions (multi-phase homogeneous mixtures) Heterogeneous: non-uniform, contains regions with different properties than other regions

Separation of Mixtures A pure substance cannot be broken down into its component substances by physical means only by a chemical process The breakdown of a pure substance results in formation of new substances (i.e. chemical change) For a pure substance there is nothing to separate (its only 1 substance to begin with) Mixtures can be separated by physical means (and also by chemical methods, as well) There are 2 general methods of separation Physical separation Chemical separation

Methods of Separation There are 2 ways of separating various substances: 1) Physical separation: separation of substances by their physical properties (such as size, solubility, etc.) Mixtures can be separated by physical separation There are several methods of separating mixtures Filtration (solids from liquids) Distillation (liquids from liquids) Centrifugation (liquids from liquids) 2) Chemical separation: separation of substances by their chemical properties Usages: Compounds can be separated into their individual elements Mixtures can be separated by chemical separation as well There are several methods of chemical separation Ion exchange (such as water purification systems) Chemical affinity (using antibodies to isolate specific proteins) Various Chemical reactions

Energy The capacity of something to do work chemical, mechanical, thermal, electrical, radiant, sound, nuclear The SI unit of energy is the Joule (J) Other common units are Calories (cal) Kilowatt-hour (kW.hr) Types of energy: Potential Kinetic Heat Energy cannot be created nor destroyed (but it does change from one type to another!)

Heat & Temperature Temperature is _____. Heat is energy that _____. how hot or cold something is (a physical property) related to the average (kinetic) energy of the substance (not the total energy) Measured in units of Degrees Fahrenheit (oF) Degrees Celsius (oC) Kelvin (K) Heat is energy that _____. flows from hot objects to cold objects is absorbed/released by an object resulting in its change in temperature Heat absorbed/released is measured by changes in temperature

Temperature Scales Fahrenheit Scale, °F Celsius Scale, °C Water’s freezing point = 32°F, boiling point = 212°F Celsius Scale, °C Temperature unit larger than the Fahrenheit Water’s freezing point = 0°C, boiling point = 100°C Kelvin Scale, K Temperature unit same size as Celsius Water’s freezing point = 273 K, boiling point = 373 K 22

Temperature of ice water and boiling water.

Heat Heat is the flow of energy due to a temperature difference Heat flows from higher temperature to lower temperature Heat is transferred due to “collisions” between atoms/molecules of different kinetic energy When produced by friction, heat is mechanical energy that is irretrievably removed from a system Processes involving Heat: Exothermic = A process that releases heat energy. Example: when a match is struck, it is an exothermic process because energy is produced as heat. Endothermic = A process that absorbs energy. Example: melting ice to form liquid water is an endothermic process. 17

Heat (cont.) The heat energy absorbed by an object is proportional to: The mass of the object (m) The change in temperature the object undergoes (DT) Specific heat capacity (s) (a physical property unique to the substance) To calculate heat (Q): Q = s . m . DT

Specific Heat Capacity (s) The amount of heat energy (in J or Cal) required to increase the temperature of 1 gram of a substance by 1oC (or 1K) The Units of Specific Heat Capacity: J/goC (SI) cal/goC (metric & more useful in the lab) Specific Heat Capacity is a unique physical property of different substances Metals have low specific heat capacity Non-metals have higher specific heat capacity Water has an unusually large specific heat capacity s = Q/(mDT)

Table of Specific Heat for various substances @ 20oC c in J/gm K c in cal/gm K or Btu/lb F Molar C J/mol K Aluminum 0.900 0.215 24.3 Bismuth 0.123 0.0294 25.7 Copper 0.386 0.0923 24.5 Brass 0.380 0.092 ... Gold 0.126 0.0301 25.6 Lead 0.128 0.0305 26.4 Silver 0.233 0.0558 24.9 Tungsten 0.134 0.0321 24.8 Zinc 0.387 0.0925 25.2 Mercury 0.140 0.033 28.3 Alcohol(ethyl) 2.4 0.58 111 Water 4.186 1.00 75.2 Ice (-10 C) 2.05 0.49 36.9 Granite .790 0.19 Glass .84 0.20

Ch 100: Fundamentals for Chemistry Chapter 4: Elements, Ions & Atoms Lecture Notes

Dmitri Mendeleev (1834-1907) Russian born chemist Considered one of the greatest teachers of his time Organized the known elements into the first “periodic table” Elements organized by chemical properties (& by weight) -> called periodic properties Predicted the existence of 3 new elements

Chemical Symbols & Formulas Each element has a unique chemical symbol Examples of chemical symbols: Hydrogen: H Oxygen: O Aluminum: Al Each molecule has a chemical formula The chemical formula indicates the chemical symbol for each of the elements present The # of atoms of each element present in the molecule Examples of chemical formulas: Elemental oxygen: O2 (2 O atoms per molecule) Water: H2O (2 H atoms & 1 O atom) Aluminum sulfate: Al2(SO4)3 (2 Al, 3 S & 12 O atoms)

Dalton’s Atomic Theory Each element consists of individual particles called atoms Atoms can neither be created nor destroyed All atoms of a given element are identical Atoms combined chemically in definite whole-number ratios to form compounds Atoms of different elements have different masses

The Atom The atom has 2 primary regions of interest: 1) Nucleus Contains protons & neutrons (called nucleons, collectively) Establishes most of the atom’s mass Mass of 1 neutron = 1.675 x10-27 kg Mass of 1 proton = 1.673 x10-27 kg Small, dense region at the center of the atom The radius of the nucleus ~ 10-15 m (1 femtometer) 2) The Electron Cloud Contains electrons Mass of 1 electron = 9.109 x10-31 kg Establishes the effective volume of the atom The radius of the electron cloud ~ 10-10 m (1 Angstrom) Determines the chemical properties of the atom During chemical processes, interactions occur between the outermost electrons of each atom The electron properties of the atom will define the type(s) of interaction that will take place

Structure of the Atom

Electric Charge Electric charge is a fundamental property of matter We don’t really know what electric charge is but we do know that there are 2 kinds: Positive charge (+) Negative charge (-) Opposite charge polarity is attractive: + attracts - Same charge polarity is repulsive: + repels + and – repels – The magnitude of electric charge (q) is the same for protons and electrons: The charge of a proton or electron is the smallest amount that occurs in nature, it is called the quantum of charge: qproton = +1.602 x 10-19 Coulombs qelectron = -1.602 x 10-19 Coulombs

What holds the atom together? Electromagnetic interaction (a.k.a. electric force) holds the electrons to the nucleus The negative charge (-) of the electrons are attracted to the positive charge (+) of the nucleus Strong interaction (a.k.a. strong force) holds the nucleons together within the nucleus The positive charge of the protons repel each other All nucleons, protons and neutrons, possess a STRONG attraction to each other that overcomes the protons’ mutual repulsion

# of neutrons = (Mass #) – (Atomic #) Atomic Bookkeeping Atomic number (Z) The number of protons in an atom The number of protons in an uncharged atom Determines the identity of the atom Mass number (A) The number of protons & neutrons in an element Determines the weight of the atom To determine number of neutrons in an atom: # of neutrons = (Mass #) – (Atomic #) Or # of neutrons = A - Z

Mass # vs. Atomic Mass (Atomic Symbol) C C Isotopes are the equivalent of sibling members of an element Unique atoms of the same element with different mass numbers (i.e. they have different numbers of neutrons) Unique isotopes are identified by their mass number Isotope notation: Example: carbon-12 ( ) & carbon-14 ( ) Atomic mass The average total mass of an element’s various naturally occuring isotopes The unit of Atomic Mass is the Dalton (formerly called the amu) 1 Dalton = one twelfth mass of one 12C atom = 1.661x10-27 kg Note: There 6 protons & 6 neutrons in a 12C atom but the mass of a 12C atom is actually less than the combined mass of all of the nucleons individually. Where is this lost mass? It’s released as energy when the nucleons combine (bind) to form the nucleus of the atom. (Atomic Symbol) Mass # Atomic # C 12 6 C 14 6

Examples of Isotopes

The Periodic Table All of the known elements are arranged in a chart called the Periodic Table The elements are arranged by similarity of chemical properties Each element is identified by its Atomic Number The elements are organized left-to-right and top-to-bottom by their Atomic Number The columns are called Groups Elements of each group have similar properties The rows are called Periods

Elements and the Periodic Table The elements can be categorized as Metals The leftmost elements of the periodic table Roughly 70% of all of the elements Nonmetals The rightmost elements of the periodic table Semimetals (metalloids) The elements between the metals and nonmetals Properties are not quite metal or non-metal

Ions Atoms (or molecules) that have gained or lost one or more electrons Ions that have lost electrons are called cations Ions that have gained extra electrons are called anions Ionic compounds have both cations and anions (so that their net charge is zero)

- - Ions (cont.) + + Ions have electric charge: “+” when 1 or more electrons are lost “-” when 1 or more electrons are gained When an atom/molecule is an ion, its charge must be specified: Sodium ion: Na+ Chloride ion: Cl- Hydroxide ion: OH- Notes on Electric Charge: Opposite charges attract Like charges repel + - + -

Ch 100: Fundamentals for Chemistry CH 100: Chemical Nomenclature (a.k.a. naming compounds)

Antoine Lavoisier (1743-1794) Considered by many to be the “Father of Modern Chemistry” Major contributions included Demonstrated that water cannot be transmuted to earth Established the Law of Conservation of Mass Developed a method of producing better gunpowder Observed that oxygen and hydrogen combined to produce water (dew) Invented a system of chemical nomenclature (still used in part today!) Wrote the 1st modern chemical textbook

Types of Compounds When compounds are formed they are held together by the association of electrons This association is called a chemical bond There are 3 general types of chemical bonds: Ionic Covalent (or molecular) Polar covalent Simple compounds are classified (and thus named) according to the type of chemical bond(s) that hold together its atoms Note: many compounds have more than one type of chemical bond present, but we will only work with “simple compounds”

Types of Compounds (cont.) For “practical” purposes will separate compounds into 2 general categories: Ionic Made up of ions (both positive and negative charge) Must have no net charge (i.e. combined charge of zero) Depend on the attraction between positive and negative charges of the ions Usually a metal is present as a cation and a nonmetal is present as an anion Molecular (or covalent) Made up of atoms that share their outer electrons Charge plays no direct role in their formation Usually no metals are present

Naming Compounds Easiest way to identify an ionic compound is to ask whether or not it has a metal present: Yes -> ionic (e.g. CaCl2) No -> covalent (e.g. CCl4) Covalent compounds require the use of Greek prefixes to indicate the number of each element present in one molecule Ionic compounds do not use the Greek prefixes

Naming Simple Compounds A “simple” or binary compound is a compound made of only 2 types of elements When the first element is a metal: The first element (metal) keeps its full name The non-metal goes by its root with the suffix “-ide” added to the end Example: NaCl is sodium chloride When there are no metals present Same as above except Greek prefixes must be used to identify the number of each element present in the compound Example: CO2 is carbon dioxide

Ionic Charges & the Periodic Table Group 1 metals form 1+ cations (Na+ sodium ion) Group 2 metals form 2+ cations (Ca2+ calcium ion) Group 13 metals form 3+ cations (Al3+ aluminum ion) All other metals (i.e. the transitional metals, Pb, etc.) form more than one type of cation Roman numerals must be used to indicate the charge of the cation Example: Fe3+ is called iron(III) FeCl3 is called iron(III) chloride Exceptions: Ag+, Cd2+ & Zn2+ Group 15 nonmetals form 3- anions (N3- nitride ion) Group 16 nonmetals form 2- anions (O2- oxide ion) Group 17 nonmetals form 1- anions (Cl- chloride ion) Group 18 nonmetals do not form ions

Greek Prefixes for Compound Names 1) Mono- 2) Di- 3) Tri 4) Tetra- 5) Penta- CCl4 is carbon tetrachloride 6) Hexa- 7) Hepta- 8) Octa- 9) Nona- 10) Deca- C3H8 is tricarbon octahydride Notes: 1) Prefixes are used when the compound does not have a metal present (or when H is the first element in the formula) 2) Prefixes must be used for every element present in the compound 3) mono- is not used for the first element in a compound name (e.g. carbon dioxide)

Ionic Compounds containing Polyatomic ions Some ionic compounds are made up of polyatomic ions When you encounter this, do not freak out!! Become familiar with the polyatomic ions on the handout Example: the nitrate ion (NO3-) The naming of this type of compound is similar to that for ionic compounds

Acids From the Latin term for “sour”{Acids are sour to the taste} Acids are substances that donate protons (H+) (usually when dissolved in water) Chemical formula usually begins with H Example: hydrochloric acid HCl(aq) + H2O(l)  H3O+ + Cl- (aq) Taste bitter Usually metal containing hydroxides Substances that accept protons (H+) when dissolved in water Example: potassium hydroxide KOH(aq) + H3O+  K+(aq) + H2O (l) Bases

Naming Acids Lets separate acids into 2 types: Acids that contain oxygen Acids that do not contain oxygen Naming acids containing oxygen: For acids containing “-ate” anions: Use root of the anion (for sulfate, SO42-, use sulfur) Add “-ic” suffix then end with “acid” Example: H2SO4 is sulfuric acid For acids with “-ite” anions: Use root of the anion (for sulfite, SO32-, use sulfur) Add “-ous” suffix then end with “acid” Example: H2SO3 is sulfurous acid

Naming Acids (cont.) Naming acids not containing oxygen: Add “hydro-” prefix to beginning Use root of the anion (i.e. Cl- use chlor) Add “-ic” suffix then end with “acid” Example: HCl is hydrochloric acid Name the following acids: HF HNO2 HCN H3PO4

Ch 100: Fundamentals for Chemistry Chapter 6: Chemical Reactions

Chemical Reactions (Intro) When matter undergoes chemical changes these processes are called chemical reactions Substances that undergo the change(s) are called the reactants The resulting substances are called the products Standard form of a chemical reaction: Reactant(s)  Product(s) Example: 2H2 (g) + 1O2 (g)  2H2O (g) The underlined numbers are called coefficients. The number of each molecule for each reactant & product in the chemical reaction They are always whole numbers

Chemical Reactions (cont.) Balanced chemical reactions indicate the ____ identity of each reactant & product involved in the reaction phase of each reactant and product involved in the reaction (i.e. solid (s), liquid (l) or gas (g)) relative quantity of each reactant and product involved in the reaction (the coefficients!) relative molar quantity of each reactant and product involved in the reaction (the coefficients!)

Rates of Chemical Reactions How quickly a chemical reaction occurs is indicated by its reaction rate How quickly the concentration of products increases How quickly the concentration of reactants decreases Factors that influence reaction rates: Reactants must be in contact Reactions occur due to collisions Without contact between reactants there can be no reaction Concentration of reactants The more reactant molecules packed into a given space the more likely a collision (& reaction) will occur Temperature the average KE of each reactant affects how much energy will be transferred between reactants during a molecular collision Molecules must transfer enough KE to break the existing bonds Energy due to collisions must be great enough to overcome the energy barrier for existing chemical bonds. The magnitude of this energy is referred to as activation energy.

Energy in Chemical Reactions Internal Energy Reactants Products Activation Energy (EA) Energy Released (Q) Exothermic Reactions Internal Energy Reactants Products Activation Energy (EA) Energy Absorbed (Q) Endothermic Reactions

Energy in Reactions (cont.) Internal Energy 2Na(s) + 2H2O(l) 2NaOH(aq) + H2(g) Low Activation Energy (EA) Large amount of Energy Released (Q) Example: Sodium Water Reaction

Catalysts Catalysts are substances that speed up chemical reactions Allow reactions to occur that might not otherwise take place (due to low temperature for example) Lower activation energy for a chemical reaction Do not participate in the reaction They may undergo a chemical change as a reactant but they are always recycled as a product (so there is no net change in the catalyst molecule) Catalysts are indicated in a chemical reaction by placing the chemical formula over/under the reaction arrow. Example: Reactants  Products catalyst

Catalysts & Energy in Reactions Internal Energy Reactants Products Activation Energy with catalyst Catalysts lower Activation Energy Activation Energy without catalyst

Endothermic or Exothermic? (that is the question…) In chemical reactions: Energy is required to break bonds (energy absorbed) Energy is released when bonds are formed The amount of energy required to break a chemical bond is the same as the energy released when the bond is formed, this is called Bond Energy During a chemical reaction: Energy is absorbed equal to the bond energies for all bonds broken in the reactants Energy is released equal to the bond energies for all bonds formed in the products Endothermic reactions absorb more energy than they release Exothermic reactions release more energy than they absorb

Balancing Chemical Reactions According to the Law of Mass Conservation (& John Dalton!) matter is never created nor destroyed All atoms in the reactants of a chemical reaction must be accounted for in the products The Basic Process: Identify all reactants & products in the reaction & write out their formulas (this is the unbalanced chemical equation) Count the number of each atom for each compound for each reactant & product (these values must be the same for both reactants & products when the reaction is balanced!) Starting with the most “complicated” molecule, systematically adjust the coefficients to balance # of the atoms on each side of the reaction (balance one atom at a time) Repeat until all atoms are balanced for the reaction Now you have a balanced chemical equation!

Balancing Chemical Reactions (example) When sodium metal is added to water a violent reaction takes place producing aqueous sodium hydroxide and releasing hydrogen gas. Write out the unbalanced chemical reaction: Now, balance the chemical reaction:

Balancing Chemical Reactions (Hint) When a polyatomic ion(s) appears on both the reactant & product side of the reaction unchanged, treat the whole ion as a “unit” when balancing the reaction Example: Note the nitrate ion (NO3-) gets swapped between the Ag + and the Ca2+ ions in this reaction So NO3- can be treated as a whole unit when balancing this reaction Balance it! AgNO3(aq) + CaCl2 (aq) AgCl(s) + Ca(NO3)2(aq)

Ch 100: Fundamentals for Chemistry Chapter 7: Chemical Reactions in Aqueous Solutions

Driving Forces & Chemical Reactions The tendency for reactants to undergo chemical changes (reactions) to form products are called “driving forces” There are 4 common “driving forces”: Formation of a solid (precipitation reaction) Formation of water (acid-base reaction) Transfer of electrons (oxidation-reduction reaction) Formation of a gas (bad taco reaction ) When 2 or more chemicals are brought together, if any of these things can happen, a chemical change is likely to occur When one of these processes occurs, we describe the resulting chemical reaction based on the driving force Write on the dry marker board the type of reaction that may occur

Solubility A measure of how much of a solute will dissolve in a solvent is called its solubility Solubility is temperature dependent Solid solubility increases with increased temperature (i.e. you can dissolve more sugar in hot water than in cold water) Gas solubility increases with decreased temperature (i.e. you can dissolve more CO2 in cold water than hot water) A solute is soluble if any of it will dissolve in a solvent NaCl is soluble in water A solute is insoluble if no appreciable amount of it will dissolve in solvent AgCl is insoluble in water When 2 solutions are combined and result in the formation of an insoluble product: The product will not dissolve in the solvent The product will form a precipitate Precipitation (formation of a solid) is one indication that a chemical change has occurred!

Precipitation Reactions in all precipitation reactions, the ions of one substance are exchanged with the ions of another substance when their aqueous solutions are mixed At least one of the products formed is insoluble in water KI(aq) + AgNO3(aq)  KNO3(aq) + AgIs K+ I- Ag+ NO3- Ag I 3

Dissociation ionic compounds metal + nonmetal (Type I & II) metal + polyatomic anion polyatomic cation + anion when ionic compounds dissolve in water the anions and cations are separated from each other; this is called dissociation we know that ionic compounds dissociate when they dissolve in water because the solution conducts electricity 4

Dissociation (examples) potassium chloride dissociates in water into potassium cations and chloride anions KCl(aq) = K+ (aq) + Cl- (aq) copper(II) sulfate dissociates in water into copper(II) cations and sulfate anions CuSO4(aq) = Cu+2(aq) + SO42-(aq) K+ Cl- K Cl Cu+2 SO42- Cu SO4 5

Dissociation (example) potassium sulfate dissociates in water into potassium cations and sulfate anions K2SO4(aq) = 2 K+ (aq) + SO42-(aq) K+ SO42- K SO4 6

Process for Predicting the Products of a Precipitation Reaction Determine what ions each aqueous reactant has Exchange Ions (+) ion from one reactant with (-) ion from other Balance Charges of combined ions to get formula of each product Balance the Equation count atoms Determine Solubility of Each Product in Water solubility rules if product is insoluble or slightly soluble, it will precipitate 7

Solubility Rules Most compounds that contain NO3- ions are soluble Most compounds that contain Na+, K+, or NH4+ ions are soluble Most compounds that contain Cl- ions are soluble, except AgCl, PbCl2, and Hg2Cl2 Most compounds that contain SO42- ions are soluble, except BaSO4, PbSO4, CaSO4 Most compounds that contain OH- ions are slightly soluble (will precipitate), except NaOH, KOH, are soluble and Ba(OH)2, Ca(OH)2 are moderately soluble Most compounds that contain S2-, CO32-, or PO43- ions are slightly soluble (will precipitate) 8

Ionic Equations equations which describe the chemicals put into the water and the product molecules are called molecular equations KCl(aq) + AgNO3(aq)  KNO3(aq) + AgCl(s) equations which describe the actual ions and molecules in the solutions as well as the molecules of solid, liquid and gas not dissolved are called ionic equations K+ (aq) + Cl- (aq) + Ag+ (aq) + NO3- (aq) K+ (aq) + NO3- (aq) + AgCl(s) ions that are both reactants and products are called spectator ions an ionic equation in which the spectator ions are dropped is called a net ionic equation Cl- (aq) + Ag+ (aq) AgCl(s) 9

Electrolytes electrolytes are substances whose aqueous solution is a conductor of electricity all electrolytes have ions dissolved in water in strong electrolytes, virtually all the molecules are dissociated into ions in nonelectrolytes, none of the molecules are dissociated into ions in weak electrolytes, a small percentage of the molecules are dissociated into ions 11

Reactions that Form Water: Acids + Bases Acids all contain H+ cations and an anion Bases all contain OH- anions and a cation when acids dissociate in water they release H+ ions and their anions when bases dissociate in water they release OH- ions and their cations 12

acid + base salt + water Acid-Base Reactions in the reaction of an acid with a base, the H+ from the acid combines with the OH- from the base to make water the cation from the base combines with the anion from the acid to make the salt acid + base salt + water H2SO4(aq) + Ca(OH)2(aq)  CaSO4(aq) + 2 H2O(l) the net ionic equation for an Acid-Base reaction is always H+ (aq) + OH- (aq)  H2O(l) 13

Reactions of Metals with Nonmetals (Oxidation-Reduction) The metal loses electrons and becomes a cation We call this process oxidation The nonmetal gains electrons and becomes an anion We call this process reduction In the reaction, electrons are transferred from the metal to the nonmetal 14

Oxidation-Reduction Reactions All reactions that involve a transfer of one or more electrons are called oxidation-reduction reactions We say that the substance that loses electrons in the reaction is oxidized and the substance that gains electrons in the reaction is reduced. 15

Predicting Products of Metal + Nonmetal Reactions metal + nonmetal  ionic compound ionic compounds are always solids unless dissolved in water in the ionic compound the metal is now a cation in the ionic compound the nonmetal is now an anion to predict direct synthesis of metal + nonmetal determine the charges on the cation and anion (from their position on the Periodic Table) determine numbers of cations and anions needed to have charges cancel balance the equation 16

Another Kind of Oxidation-Reduction Reaction Some reactions between two non-metals are also oxidation-reduction reaction Any reaction in which O2 is a reactant or a product will be an oxidation-reduction reaction Examples: CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(g) 2 SO3(g)  2 SO2(g) + O2(g) 17

Ways to Classify Reactions Reactions that involve solid formation are called precipitation reactions Reactions that involve water formation are called acid-base reactions Both precipitation reactions and acid-base reactions involve compounds exchanging ions, ion exchange reactions are called double displacement reactions 18

Double Displacement Reactions two ionic compounds exchange ions X Y (aq) + AB (aq)  XB + AY reaction will not occur unless one of the products either (1) precipitates or (2) is water 19

Ways to Classify Reactions Reactions that involve electron transfer are called oxidation-reduction reactions Metals + Nonmetal O2 as a reactant or product Reactions that occur in aqueous solution where one of the products is a gas are called gas forming reactions NaHCO3(aq) + HCl(aq)  NaCl(aq) + CO2(g) + H2O(l) 20

Ways to Classify Reactions Reactions that involve one ion being transferred from one cation to another are called single replacement reaction X Y + A X + AY Examples: Zn(s) + 2 HCl(aq)  ZnCl2(aq) + H2(g) Fe2O3(s) + 2 Al(s)  2 Fe(s) + Al2O3(s) 21

Other Ways to Classify Reactions Reactions in which O2(g) is reacted with a carbon compound are called Combustion Reactions Combustion reactions release a lot of energy Combustion reactions are a subclass of Oxidation-Reduction reactions Combustion of carbon compounds produces CO2(g) Combustion of compounds that contain hydrogen produces H2O(g) C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g) 22

Other Ways to Classify Reactions Reactions in which chemicals combine to make one product are called Synthesis Reactions Metal + Nonmetal reactions can be classified as Synthesis Reactions 2 Na(s) + Cl2(g)  2 NaCl(s) Reactions of Metals or Nonmetals with O2 can be classified as Synthesis Reactions N2(g) + O2(g)  2 NO(g) These two types of Synthesis Reactions are also subclasses of Oxidation-Reduction Reactions 23

Other Ways to Classify Reactions Reactions in which one reactant breaks down into smaller molecules are called Decomposition Reactions Generally initiated by addition of energy Addition of electric current or heat Opposite of a Synthesis Reaction 2 NaCl(l)  2 Na(l) + Cl2(g) electric current 24

Ch 100: Fundamentals for Chemistry Chapter 8 Lecture Notes (Sections 8.1 to 8.5)

Amadeo Avogadro (1743-1794) Italian lawyer turned chemist Major contributions included: Established difference between atoms & molecules: Oxygen & nitrogen exist as molecules O2 & N2 Reconciled the work of Dalton & Guy-Lussac Establishing Avogadro’s Principle: equal volumes of all gases at the same temperature and pressure contain the same number of molecules. Did not determine Avogadro’s number nor the mole (these concepts came later) Avogadro is honored because the molar volume of all gases should be the same Much of Avogadro’s work was acknowledged after he died, by Stanislao Cannizarro

The Mole A counting unit (similar to the dozen) A large unit used to describe large quantities such as number of atoms 1 mole = 6.022 x 1023 units 6.022 x 1023 is known as Avogadro’s number (NA) Relationship between the mole & the Periodic Table The atomic mass is the quantity (in grams) of 1 mole of that element The units of atomic mass are grams/mole Mass is used by chemists as a way of “counting” number of atoms/molecules of a substance Mole calculations

What do you get if you have Avogadro's number of donkeys? Got mole problems? Call Avogadro at 602-1023. What do you get if you have Avogadro's number of donkeys? Answer: molasses (a mole of asses)

Molar Mass Mass in grams of 1 mole of a substance Refers to both atoms & molecules Elements (atoms) Examples: 1 mole of Na has a mass of 22.99 g 1 mole of Cl has a mass of 35.45 1 mole of Cl2 has a mass of 70.90 g Compounds (molecules) 1 mole of NaCl has a mass of 58.44 g Mass of Na (22.99 g) + Mass of Cl (35.45 g) 1 mole of CO2 has a mass of 44.01 g Mass of C (12.01 g) + 2 x Mass of O (16.00 g)

Mole Calculations (1) Atoms/Molecules to Moles Divide # of atoms (or molecules) by Avogadro’s # Example: How many moles are 1.0x1024 atoms? Moles to Atoms/Molecules Multiply # of atoms (or molecules) by Avogadro’s # Example: How many molecules are in 2.5 moles?

Mole Calculations (2) Moles to Grams Grams to Moles Multiply the # of moles by atomic mass Example: How many grams in 2.5 moles of carbon? Grams to Moles Divide the mass in grams by atomic mass Example: How many moles are in 2.5 grams of lithium?

Percent Composition Percentage of each element in a compound (by mass) Can be determined from: the formula of the compound or the experimental mass analysis of the compound Note: The percentages may not always total to 100% due to rounding 12

Percent Composition Calculations To determine % Composition from the chemical formula: Determine the molar mass of compound Multiply the molar mass of the element of interest by the number of atoms per molecule then Divide this value by the molar mass of the compound Example: The % Composition of sodium in table salt The molar mass of NaCl is 58.44 g/mol There is 1 atom of Na in each NaCl molecule The atomic mass of Na is 22.99

Percent Composition Calculations Perform the following % Composition calculations: The % composition of carbon in carbon monoxide The % composition of oxygen in water The % composition of chlorine in sodium hypochlorite

Ch 100: Fundamentals for Chemistry Ch 9: More on Chemical Reactions Lecture Notes (Sections 9.1 to 9.2)

Chemical Equations: What do they tell us? A properly written chemical equation will provide the following information: All reactants & products involved in the reaction The physical state of all reactants & products The presence of any catalysts involved in the chemical reaction The relative quantity of all reactants & products

Information Given by the Chemical Equation Balanced equation provides the relationship between the relative numbers of reacting molecules and product molecules 2 CO + O2  2 CO2 2 CO molecules react with 1 O2 molecules to produce 2 CO2 molecules 2

Information Given by the Chemical Equation Since the information given is relative: 2 CO + O2  2 CO2 200 CO molecules react with 100 O2 molecules to produce 200 CO2 molecules 2 billion CO molecules react with 1 billion O2 molecules to produce 20 billion CO2 molecules 2 moles CO molecules react with 1 mole O2 molecules to produce 2 moles CO2 molecules 12 moles CO molecules react with 6 moles O2 molecules to produce 12 moles CO2 molecules 3

Information Given by the Chemical Equation The coefficients in the balanced chemical equation shows the molecules and mole ratio of the reactants and products Since moles can be converted to masses, we can determine the mass ratio of the reactants and products as well 4

Information Given by the Chemical Equation 2 CO + O2  2 CO2 2 moles CO = 1mole O2 = 2 moles CO2 Since 1 mole of CO = 28.01 g, 1 mole O2 = 32.00 g, and 1 mole CO2 = 44.01 g 2(28.01) g CO = 1(32.00) g O2 = 2(44.01) g CO2 5

2 moles CO = 1 mol O2 = 2 moles CO2 Example: Determine the Number of Moles of Carbon Monoxide required to react with 3.2 moles Oxygen, and determine the moles of Carbon Dioxide produced Write the balanced equation 2 CO + O2  2 CO2 Use the coefficients to find the mole relationship 2 moles CO = 1 mol O2 = 2 moles CO2 6

Use dimensional analysis Example (cont.) Determine the Number of Moles of Carbon Monoxide required to react with 3.2 moles Oxygen, and determine the moles of Carbon Dioxide produced Use dimensional analysis 7

Ch 100: Fundamentals for Chemistry Ch 14: Solutions & Concentration Lecture Notes

Solutions Solutions are single phase homogenous mixtures Solutions consist of: Solvent: the component in largest quantity Solute(s): the other components The solute is considered to be dissolved in the solvent When a solution has not reached its limit of dissolved solute it is an unsaturated solution When a solution has reached its limit of dissolved solute and any added solute will not dissolve, it is a saturated solution

Concentration A measure of how much of a substance (solute) is dissolved in another substance (solvent) To calculate [concentration]: Common usages of concentration: Mass (m/v) conc. (units are grams/L, grams/mL, etc.) Volume (v/v) conc. (unit-less, often % is used) Molarity (units are moles/L or M)

Mass Percent (%) Concentration of a solute dissolved in a solvent (in grams per unit gram of solution) To determine mass percent Divide mass of solute (in grams) by the total mass of solution (in grams) and multiply this ration by 100% Example: What is the mass percent of 30.0 grams of NaCl in a 150.0 gram solution? Questions: (a) How much CaCl2 is in 250.0 grams of solution where the mass percent of CaCl2 is 30.0%? (b) How much Cl- is in this solution?

Molarity Concentration of a solute (in moles per unit volume) dissolved in a solvent SI units are moles/liter, or M (molarity or molar concentration) To determine molarity from mass concentration Simply a unit conversion from grams to moles (using atomic or molar mass as the unit conversion) Example: What is the molarity of a NaCl solution with concentration of 30.0 grams/L? The molar mass of NaCl is 22.99 + 35.45 = 58.44 grams/mol Question: What is the molarity of Na+ and Cl- in this solution?

The pH Scale The acidity (or concentration of H+) of a solution is often measured using the pH scale The pH scale is based on the molarity of H+ ions in solution The pH scale ranges from 0 (acidic) to 14 (basic) When pH=7.0 the solution is neutral acidity, there is equal concentration of H+ and OH- in the bulk liquid To calculate pH from [H+] (in mol/L): pH=-log10[H+] Example: a solution with [H+]=1.0x10-5 M [H+]=1.0x10-5 M then pH = -log10(1.0x10-5) = 5.0 To calculate [H+] from pH: [H+]=10-pH Example: a solution with pH = 9.0 pH = 9.0 then [H+]=10-9.0 M = 1.0x10-9 M

pH Concept Questions What is the [H+] for a 0.5 M HCl solution? 0.5 M What is the [H+] for a 0.5 M H3PO4 solution? Less than 0.5 M Do the 2 solutions above have the same pH? No, pH depends on [H+] not [acid] Why or why not? HCl is a strong acid but H3PO4 is a weak acid How does a strong acid differ from a weak acid? Strong acids dissociate all of their H+ ions when in water whereas weak acids do not!