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 Method1. 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 gunpowder1661: Robert Boyle proposed the concept of elements1770-1790: Lavoisier proposed the concept of compounds & the Law of Mass

Conservation1774: Priestly isolates oxygen1797: Proust proposed the Law of Definite Proportions1803: Dalton re-introduces the concept of the atom and establishes Dalton’s Laws1869: Mendeleev creates the 1st Periodic Table1910: Rutherford proposes the “nuclear” model of the atom1915: Bohr proposes a “planetary” model of the hydrogen atom1920: Schroedinger publishes his wave equation for hydrogen1969: Murray Gell-Mann proposes the theory of QCD (proposing the existence of quarks)

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 radioactivity1911: H. Kamerlingh Onnes discovers superconductivity in low temperature

mercury1947: William Shockley and colleagues invent the first transistor1996: Cornell, Wieman, and Ketterle observe the 5th state of matter (the Bose-

Einstein condensate) in the laboratory

Major Developments in Chemistry II

Ch 100: Fundamentals for Chemistry

Chapter 2: Measurements & CalculationsLecture Notes

Types of Observations• Qualitative

Descriptive/subjective in natureDetail qualities such as color, taste, etc.Example: “It is really warm outside today”

• QuantitativeDescribed by a number and a unit (an accepted

reference scale)Also known as measurementsExample: “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

NotationFirst Compare Exponents of 10 (order of

magnitude)Then Compare Numbers

Writing Numbers in Scientific Notation

1 Locate the Decimal Point2 Move the decimal point to the right of the

non-zero digit in the largest placeThe new number is now between 1 and 10

3 Multiply the new number by 10n

where n is the number of places you moved the decimal point

4 Determine the sign on the exponent, nIf 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

Writing Numbers in Standard Form1 Determine the sign of n of 10n

If n is + the decimal point will move to the rightIf n is – the decimal point will move to the left

2 Determine the value of the exponent of 10Tells the number of places to move the decimal

point3 Move the decimal point and rewrite the

number

Measurement SystemsThere are 3 standard unit systems we will focus on:

1. United States Customary System (USCS)formerly the British system of measurementUsed in US, Albania, and a couple othersBase units are defined but seem arbitrary (e.g. there are 12 inches in 1 foot)

2. MetricUsed by most countriesDeveloped in France during Napoleon’s reignUnits 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 unitsUsed by scientists and most science textbooksNot 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

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

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

Rules for Counting Significant Figures

• Nonzero integers are always significant• Zeros

Leading zeros never count as significant figuresCaptive zeros are always significantTrailing zeros are significant if the number has a

decimal point• Exact numbers have an unlimited number of

significant figures

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!!

Exact Numbers• Exact Numbers are numbers known with

certainty • Unlimited number of significant figures• They are either

counting numbersnumber of sides on a square

or defined100 cm = 1 m, 12 in = 1 ft, 1 in = 2.54 cm1 kg = 1000 g, 1 LB = 16 oz1000 mL = 1 L; 1 gal = 4 qts.1 minute = 60 seconds

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

VolumeMassDensity

Using Density in Calculations

VolumeMassDensity

DensityMass Volume

Volume Density Mass

Ch 100: Fundamentals for Chemistry

Chapter 3: Matter & Energy Lecture Notes

• 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)

Aristotle (384-322 BC)

4) Fire2) Water

3) Air1) Earth

4) Dry2) Moist

3) Hot1) Cold

Physical & Chemical Properties• Physical Properties are the characteristics of

matter that can be changed without changing its compositionCharacteristics 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 energyCharacteristics that describe the behavior of matter

Physical & Chemical Changes• Physical Changes are changes to matter

that do not result in a change the fundamental components that make that substanceState Changes : boiling, melting, condensing

• Chemical Changes involve a change in the fundamental components of the substanceProduce a new substanceChemical reactionReactants Products

Solid → Liquid → Gas

States of Matter

State Shape Volume Compress Flow

Solid Keeps Shape

Keeps Volume

No No

Liquid Takes Shape of Container

Keeps Volume

No Yes

Gas Takes Shape of Container

Takes Volume of Container

Yes Yes

+Energy +Energy

Solid ← Liquid ← Gas+Energy +Energy

Classification of MatterMatter 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

P u re S ub s ta n ceC o n s tan t C o m p o s it ion

H o m o ge n e o us

M ix tu reV a ria b le Co m p o s it ion

M a tte r

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 separationThere 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 propertiesUsages:

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:PotentialKineticHeat

• Energy cannot be created nor destroyed (but it does change from one type to another!)

Heat & Temperature• Temperature is _____.

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, °FWater’s freezing point = 32°F, boiling point = 212°F

• Celsius Scale, °CTemperature unit larger than the FahrenheitWater’s freezing point = 0°C, boiling point = 100°C

• Kelvin Scale, KTemperature unit same size as CelsiusWater’s freezing point = 273 K, boiling point = 373 K

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:

1. Exothermic = A process that releases heat energy. Example: when a match is struck, it is an exothermic process because

energy is produced as heat.2. Endothermic = A process that absorbs energy.

Example: melting ice to form liquid water is an endothermic process.

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

(T)Specific heat capacity (s) (a physical property unique to

the substance)

• To calculate heat (Q):Q = s . m . T

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:1. J/goC (SI)2. 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/(mT)

Table of Specific Heat for various substances @ 20oC

Substance c in J/gm K c in cal/gm K orBtu/lb F

Molar CJ/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: HOxygen: OAluminum: Al

• Each molecule has a chemical formula• The chemical formula indicates

the chemical symbol for each of the elements presentThe # 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 Theory1. Each element consists of individual particles

called atoms2. Atoms can neither be created nor destroyed3. All atoms of a given element are identical4. Atoms combined chemically in definite

whole-number ratios to form compounds5. Atoms of different elements have different

masses

The AtomThe 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 nucleusThe 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 nucleusThe positive charge of the protons repel each otherAll nucleons, protons and neutrons, possess a STRONG

attraction to each other that overcomes the protons’ mutual repulsion

Atomic Bookkeeping• Atomic number (Z)

The number of protons in an atomThe number of protons in an uncharged atomDetermines the identity of the atom

• Mass number (A)The number of protons & neutrons in an elementDetermines 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• 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 #

C126

C146

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 TableThe elements can be categorized as

MetalsThe leftmost elements of the periodic tableRoughly 70% of all of the elements

NonmetalsThe rightmost elements of the periodic table

Semimetals (metalloids)The elements between the metals and nonmetalsProperties 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:

1. Ionic2. Covalent (or molecular)3. Polar covalent

• Simple compounds are classified (and thus named) according to the type of chemical bond(s) that hold together its atomsNote: 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 ionsUsually 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 electronsCharge plays no direct role in their formationUsually 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 CompoundsA “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 endExample: 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 compoundExample: 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

cationRoman numerals must be used to indicate the charge of the cationExample:

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) Tri4) Tetra-5) Penta-

CCl4 is carbon tetrachloride

6) Hexa-7) Hepta-8) Octa-9) Nona-10) Deca-

C3H8 is tricarbon octahydrideNotes:

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 handoutExample: the nitrate ion (NO3

-)• The naming of this type of compound is

similar to that for ionic compounds

AcidsFrom 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 HExample: hydrochloric acid

HCl(aq) + H2O(l) H3O+ + Cl- (aq)

Taste bitterUsually metal containing hydroxidesSubstances that accept protons (H+) when dissolved in waterExample: potassium hydroxide

KOH(aq) + H3O+ K+(aq) + H2O (l)

Bases

Naming AcidsLets separate acids into 2 types:

Acids that contain oxygen Acids that do not contain oxygen

Naming acids containing oxygen: For acids containing “-ate” anions:

1. Use root of the anion (for sulfate, SO42-, use sulfur)

2. Add “-ic” suffix then end with “acid”

Example: H2SO4 is sulfuric acid

For acids with “-ite” anions:1. Use root of the anion (for sulfite, SO3

2-, use sulfur) 2. Add “-ous” suffix then end with “acid”

Example: H2SO3 is sulfurous acid

Naming Acids (cont.)Naming acids not containing oxygen:

Add “hydro-” prefix to beginningUse 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:HFHNO2

HCNH3PO4

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

reactionThey 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 rateHow quickly the concentration of products increasesHow quickly the concentration of reactants decreases

• Factors that influence reaction rates:Reactants must be in contact

Reactions occur due to collisionsWithout contact between reactants there can be no reaction

Concentration of reactantsThe more reactant molecules packed into a given space the more likely a

collision (& reaction) will occurTemperature

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 in Chemical ReactionsInternal Energy

Reactants

Products

Activation Energy (EA)

Energy Released (Q)

Exothermic Reactions

Internal Energy

Reactants

ProductsActivation

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

reactionsAllow reactions to occur that might not otherwise take place

(due to low temperature for example)Lower activation energy for a chemical reactionDo 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 EnergyActivation 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 reactantsEnergy 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 destroyedAll 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 reactionNow 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.

1. Write out the unbalanced chemical reaction:

2. 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”:1. Formation of a solid (precipitation reaction)2. Formation of water (acid-base reaction)3. Transfer of electrons (oxidation-reduction reaction)4. 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

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) + AgIs

K+

I-

Ag+

NO3-

K+

NO3-

Ag I

Dissociation• ionic compounds

metal + nonmetal (Type I & II)metal + polyatomic anionpolyatomic 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

Dissociation (examples)• potassium chloride dissociates in water into

potassium cations and chloride anionsKCl(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

Dissociation (example)• potassium sulfate dissociates in water into

potassium cations and sulfate anionsK2SO4(aq) = 2 K+ (aq) + SO4

2-(aq)

K+

SO42-

K+

KK SO4

Process for Predicting the Products of

a Precipitation Reaction1) Determine what ions each aqueous reactant has2) Exchange Ions

(+) ion from one reactant with (-) ion from other3) Balance Charges of combined ions to get formula of

each product4) Balance the Equation

count atoms5) Determine Solubility of Each Product in Water

solubility rules if product is insoluble or slightly soluble, it will precipitate

Solubility Rules1. Most compounds that contain NO3

- ions are soluble2. Most compounds that contain Na+, K+, or NH4

+ ions are soluble

3. Most compounds that contain Cl- ions are soluble, except AgCl, PbCl2, and Hg2Cl2

4. Most compounds that contain SO42- ions are soluble,

except BaSO4, PbSO4, CaSO4

5. 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

6. Most compounds that contain S2-, CO32-, or PO4

3- ions are slightly soluble (will precipitate)

Ionic Equations• equations which describe the chemicals put into the water and the

product molecules are called molecular equationsKCl(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 ionsK+

(aq) + Cl- (aq) + Ag+ (aq) + NO3

- (aq) K+

(aq) + NO3-

(aq) + AgCl(s)

• an ionic equation in which the spectator ions are dropped is called a net ionic equation

Cl- (aq) + Ag+ (aq) AgCl(s)

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

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

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 saltacid + 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)

Reactions of Metals with Nonmetals

(Oxidation-Reduction)• The metal loses electrons and becomes a

cationWe call this process oxidation

• The nonmetal gains electrons and becomes an anionWe call this process reduction

• In the reaction, electrons are transferred from the metal to the nonmetal

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.

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

1) determine the charges on the cation and anion(from their position on the Periodic Table)

2) determine numbers of cations and anions needed to have charges cancel

3) balance the equation

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)

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

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

Ways to Classify Reactions

• Reactions that involve electron transfer are called oxidation-reduction reactionsMetals + NonmetalO2 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)

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 + AY

• Examples:Zn(s) + 2 HCl(aq) ZnCl2(aq) + H2(g)

Fe2O3(s) + 2 Al(s) 2 Fe(s) + Al2O3(s)

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)

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 ReactionsN2(g) + O2(g) 2 NO(g)

• These two types of Synthesis Reactions are also subclasses of Oxidation-Reduction Reactions

Other Ways to Classify Reactions• Reactions in which one reactant breaks

down into smaller molecules are called Decomposition Reactions

• Generally initiated by addition of energyAddition of electric current or heat

• Opposite of a Synthesis Reaction2 NaCl(l) 2 Na(l) + Cl2(g)

electriccurrent

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-LussacEstablishing 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 sameMuch 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 TableThe atomic mass is the quantity (in grams) of 1 mole of that elementThe units of atomic mass are grams/moleMass is used by chemists as a way of “counting” number of

atoms/molecules of a substance• Mole calculations

Got mole problems? Call Avogadro at 602-1023.

Answer: molasses (a mole of asses)

What do you get if you have Avogadro's number of

donkeys?

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 g1 mole of Cl has a mass of 35.451 mole of Cl2 has a mass of 70.90 g

• Compounds (molecules)Examples: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 gMass 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/MoleculesMultiply # of atoms (or molecules) by Avogadro’s #Example: How many molecules are in 2.5 moles?

2423

1(1.0 10 ) 1.76.022 10

moleatoms moles

23246.022 10(2.5 ) 1.5 10

1moles molecules

mole

Mole Calculations (2)• Moles to Grams

Multiply the # of moles by atomic massExample: How many grams in 2.5 moles of carbon?

• Grams to MolesDivide the mass in grams by atomic massExample: How many moles are in 2.5 grams of lithium?

112.01(2.5 ) 30. ( 3 10 )1gramsmoles grams ormole

11(2.5 ) 0.36 ( 3.6 10 )6.941

molegrams moles orgrams

Percent Composition• Percentage of each element in a compound (by

mass)• Can be determined from:

1. the formula of the compound or2. the experimental mass analysis of the compound

Note: The percentages may not always total to 100% due to rounding

% 100%partCompositionwhole

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 salt1. The molar mass of NaCl is 58.44 g/mol2. There is 1 atom of Na in each NaCl molecule3. The atomic mass of Na is 22.99

(# )( )% 100%atoms of A atomic mass of AComposition of Amolar mass of compound

1 22.99% 100% 39.33%58.44

Composition of Na

Percent Composition Calculations

Perform the following % Composition calculations:1.The % composition of carbon in carbon

monoxide2.The % composition of oxygen in water3.The % composition of chlorine in sodium

hypochlorite

Ch 100: Fundamentals for Chemistry

Ch 9: More on Chemical ReactionsLecture Notes (Sections 9.1 to 9.2)

Chemical Equations: What do they tell us?

• A properly written chemical equation will provide the following information:

1. All reactants & products involved in the reaction

2. The physical state of all reactants & products3. The presence of any catalysts involved in the

chemical reaction4. The relative quantity of all reactants &

products

Information Given by theChemical 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

Information Given by theChemical Equation

• Since the information given is relative:2 CO + O2 2 CO2

200 CO molecules react with 100 O2 molecules to produce 200 CO2 molecules2 billion CO molecules react with 1 billion O2 molecules to produce 20 billion CO2 molecules2 moles CO molecules react with 1 mole O2 molecules to produce 2 moles CO2 molecules12 moles CO molecules react with 6 moles O2 molecules to produce 12 moles CO2 molecules

Information Given by theChemical 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

Information Given by theChemical 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

1. Write the balanced equation 2 CO + O2 2 CO2

2. Use the coefficients to find the mole relationship

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

3. Use dimensional analysis

CO moles 6.4O mole 1CO moles 2xO moles 3.2

2 2

22

22 CO moles 6.4

O mole 1CO moles 2 xO moles 3.2

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

Ch 100: Fundamentals for Chemistry

Ch 14: Solutions & ConcentrationLecture Notes

Solutions• Solutions are single phase homogenous mixtures• Solutions consist of:

Solvent: the component in largest quantitySolute(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)

solventofamount

soluteofamountsoluteofconc .

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?

%100%

solutionofmasssoluteofmassMass

%0.20%1000.1500.30%

solutiongramsNaClgramsMass

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?

Mgramsmol

LgramsNaCl 513.0

44.581

10.30

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 = 9.0pH = 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!