Syllabus Chemistry 101 Fall 2009 Sec. 501 (MWF 8:00-8:50) Sec 505 (MWF 12:40-1:30) RM 100 HELD Professor: Dr. Earle G. Stone Office: Room 123E Heldenfels (HELD) Telephone: 845-3010 (no voice mail) or leave a message at 845-2356 email: [email protected]

(put CHEM 101-Sec. # + subject in subject line of your email) Office Hours: HELD 123E: Tue. And Thur. 8:00-10:50 AM I.A. TBA S.I. Leader: TBA CHEM 101 and 102 are the first-year chemistry sequence in the core curriculum. These are 3-credit courses. All lecture sections strive to cover common content. The lecture component of Chemistry 101 covers stoichiometry, atomic and molecular structure, chemical bonding, fundamental acid/base chemistry, solution chemistry, properties of liquids and solids, the gas laws, and this class will additionally cover foundation work in inorganic and organic nomenclature and structure, including some mention of major biological organic and inorganic compounds. More importantly, it is the goal of my lecture section to help you develop the skill set to successfully complete your undergraduate degree and as most in this class are pre-something to prepare for your professional school entrance exam and enable you to succeed in your choice of professional school.

Kotz and Treichel 7 th ed. TEXTBOOKS Averill and Eldridge

Hardbound ~$200 Solution Manual ~$40 Online Tutor ~$45 Total ~$285 Ebook $45 per semester Includes Text Solution manual Online tutorial Helpful Yvette Freeman Publisher's Representative Pearson Education Online Dictionary of Chemistry Useful As A Second Language General Chemistry I and Organic Chemistry I (There are O-chem II and Physics books in this series if you find these useful and will have to take those classes.

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Grading:

Your grade will be based on

•Four one-hour examinations (each worth 200 points) •A final examination (400 points)

There are no bonuses, no extra credit, no soft points Major Examination Schedule Fall 2005: Wed. Sept. 16 Major Exam No.1

Wed. Oct. 7 Major Exam No.2

Mon. Nov. 9 Major Exam No.3

Fri. Dec. 4 Major Exam No. 4 Fri. Dec. 11 Section 501 Final Exam 10:00 to 12:00 Mon. Dec. 14 Section 506 Final Exam 10:30 to 12:30

How grades are determined

The way the real world works Individual Mastery compared to a large population What you are used to and I will report 1) Raw scores are determined. 2) Individual scores are normalized. 3) Normalized scores are transformed. 4) Letter grades are assigned >3 >2 >1 >0 <0 <-1 <-2 <-3 Normal Exam 1 Exam 2 Exam 3 Exam 4

1 10 61 0 4 78 0 3 73 0 5 75 0 4 64 154 154 61 10 1 146

154

56 12 2 163 134 65 12 2 146 156 58

10

2 175 133 59 15 2

Final

0 8 67 151 145 73 8 0 Sum of points assigned to correct responses A context-free evaluation of relative performance An absolute score is assigned to a defined scale >89.451 A >79.451 B >69.451 C >59.451 D <59.451 F

And we begin:

Problem A situation that presents difficulty, uncertainty, or perplexity: The mere formulation of a problem is far more often essential than its solution, which may be merely a matter of mathematical or experimental skill. To raise new questions, new possibilities, to regard old problems from a new angle requires creative imagination and marks real advances in science.

~ Albert Einstein

Question A request for data: inquiry, interrogation, query. Answer A spoken or written reply, as to a question.

Solution Something worked out to explain, resolve, or provide a method for dealing with and settling a problem.

Scientific Method -

answers to questions and solutions to problems, which consists of: A procedure that searches for

http://museum.nist.gov/exhibits/adx2/index.htm

http://museum.nist.gov/exhibits/adx2/index.htm

http://www.batteryequaliser.com/behome.html

http://www.patentstorm.us/patents/5945236-description.html

Observations

•

Can be qualitative or quantitative

•

Qualitative observations describe properties or occurrences in ways that do not rely on numbers.

•

Quantitative observations are measurements that consist of a number a unit and a label

Hypotheses

•

A tentative explanation for the observations that may not be correct, but puts the scientist’s understanding of the system being studied into a form that can be tested

Experiments

•

Tests the validity of the hypothesis

•

Are systematic observations or measurements made under controlled conditions, in which the variable of interest is clearly distinguished from any others

•

If experimental results are reproducible, they are summarized in a

Law

law.

•

A verbal or mathematical description of a phenomenon that allows for general predictions that describes what happens and not why and is unlikely to change greatly over time unless a major experimental error is discovered.

Theory

•

Attempts to explain why nature behaves as it does which is incomplete and imperfect and evolves with time to explain new facts as they are discovered

Natural Laws

• Law of Conservation of Mass –

matter, can be neither created nor destroyed.

The notion that mass, or

• Law of Conservation of Energy –

A law that states that in any system not involving nuclear reactions or velocities approaching the velocity of light, energy cannot be created or destroyed. The First Law of Thermodynamics.

• Law of Conservation of Mass – Energy –

Einstein’s General Theory of Relativity - E=mc 2 – with work becomes the special theory of relativity, which has been verified by experiment, has shown that the mass of a body changes as the energy possessed by the body changes. Such changes in mass are too small to be detected except in subatomic phenomena. Matter may be created by the materialization of a photon into an electron-positron pair; or it may be destroyed, by the annihilation of this pair of elementary particles to produce a pair of photons.

Natural Laws

Law of Definite Proportions –

When two or more elements combine to form a compound, their masses in that compound are in a fixed and definite ratio. This data helps justify an atomic view of matter.

Law of Multiple Proportions –

When two elements combine to form more than one compound, the mass of element A which combines in the first compound with a given amount of element B has a simple whole number ratio with the mass of element A which combines in the second compound with the same given mass of element B.

Dr. Stone’s patent pending chemistry problem solver 1. Write down everything you are given

• •Vocabulary •Numbers •Units

2. Write down what you want to know

•Vocabulary •Numbers •Units

3. Write down mathematical equation(s) that include(s) these values and units Principles 4. Write a balanced stoichiometric equation

•Mole concept

5. Convert everything to moles

•Dimensional analysis

6. Convert everything to the unknown’s units

•Rounding, significant figures, accuracy and

precision

1. Numbers –

Significant Figures, Rounding Rules, Accuracy, Precision, Statistical Treatment of the Data

2. Units –

1. Density?

2. Molecular Weight (Mass) 3. Mole Ratio, Molarity, molality

3. Vocabulary –

Approximately 100 new terms or words and applying new or more rigid definitions to words you may already own.

4. Principles (Theories and Laws) –

Mathematical formulas are used to model observed phenomena to predict possible outcomes Stoichiometry, Quantum Theory, Bonding, Chemical Periodicity, Solutions, Thermodynamics, Intermolecular Forces, Gas Laws c p = q/m D T D G = D H – T D S PV = nRT D T = Kmi rate = k[A] m [B] n E o cell = E cathode = E anode %yield = actual/theoretical * 100%  n = c (ms -1 ) l (m) ∆E = q + w K = [C] c [D] d [A] a [D] b

Use of Numbers

• Exact numbers – 1 dozen = 12 things for example • Accuracy – how closely measured values agree with the

correct value

• Precision – how closely individual measurements agree

with each other

• Significant Figures – start at the left and proceed

to the right

– If the number does not have a decimal point

count until there are no more non zero numbers

– If the number has a decimal point start

counting at the first non-zero number and continue counting until you run out of decimal places

• Scientific notation – use it.

• •

Use of Numbers

Multiplication & Division rule Easier of the two rules Product has the smallest number of significant figures of multipliers

4 .

24 2 2 .

783 2 x 1.2

3 x 1.

4 5 .

21766 3 .

89648 round off to 5.22

round off to 3.9

Addition & Subtraction rule More subtle than the multiplication rule Answer contains smallest decimal place of the addends.

3 .

692 3  1 .

23 4  2 .

0 2 8  .

793 2 .

7 12 3 6 .

9463 6 .

6707 round off to 6.95

round off to 6.671

When a 5 appears. Is there anything to the right of the 5 greater than zero?

Is the number to the left of the 5 odd?

Is the number to the left of the 5 even? (Treat 0 as even.)

How many sig figs?

a) 0.0713200

b) 7843000 c) 1.4800

d) 100 e) 100.0

f) 894.003

g) 89400 h) 0.03000

i) 74.000

Round off to two sig figs

a) 34.78

b) 17.51

c) 48.50

d) 45.50001

e) 24.33

f) 17.50

g) 20.5

h) 45.5000

Use of Numbers

How many sig figs in the answer?

a) 472x101 b) 4600x0.005

c) 36.0x4752

d) 45.08/36.2

e) 1.003/8500 f) 0.003/472x12 g) 3.003/475.0x0.30/524 h) 0.3005x4.1

i) 23.56+24.983

j) 4.78-2.892

k) 46.83-0.03

l) 34.892+5.0

m) 134.033-0.02

n) 48.2-46 (1.68) [ 23.56 – 2.3

1.248 x 10 3 ] =

Vocabulary

• Chemistry - Science that describes matter – its properties-composition-structure, the changes it undergoes, and the energy changes that accompany those processes • Matter - Anything that has mass and occupies space.

• Energy - The capacity to do work or transfer heat.

• Chemical Properties - chemical changes - describes the characteristic ability of a substance to react to form new substances (flammability and corrosion). – rusting or oxidation, chemical reactions • Physical Properties - physical changes - Characteristics that scientists can measure without changing the composition of the sample under study (mass, color, volume, amount of space occupied by the sample).

– changes of state, density, color, solubility • Extensive Properties - depend on quantity a. Vary with the amount of the substance b. Include mass, weight, and volume.

• Intensive Properties

-

do not depend on quantity a. Include color, melting and boiling point, electrical conductivity, and physical state at a given temperature b. Determine a substance’s identity, c. Have an important intensive property called density (d), a ratio of two extensive properties, mass and volume density = mass

d = m volume V

Vocabulary

• Three distinct

states of matter:

1. Solids — relatively rigid and have fixed shapes and volumes — volumes of solids independent of temperature and pressure 2. Liquids — — have fixed volumes but flow to assume the shape of their containers Volumes of liquids independent of temperature and pressure 3. Gases — have neither fixed shapes nor fixed volumes and expand to fill their containers completely — Depends strongly on temperature and pressure

Vocabulary

–

Isotopes and Atomic Masses

    

Atoms of different elements exhibit different chemical behavior.

Identity of an element is defined by its atomic number. (Z) is the number of protons in the nucleus of an atom of the element.

The atomic number is therefore different for each element.

Known elements are arranged in order of increasing Z in the periodic table.

      

Vocabulary

–

Isotopes and Atomic Masses

The chemistry of each element is determined by its number of protons and electrons.

In a neutral atom, the number of electrons equals the number of protons.

Symbols for elements are derived directly from the element’s name.

Nuclei of atoms contain neutrons as well as protons.

The number of neutrons is not fixed for most elements, unlike protons.

Atoms that have number, but the same number of protons, and hence the same atomic different numbers of neutrons are called isotopes . Isotopes All isotopes of an element have the same number of protons and electrons, which means they exhibit the same chemistry. Isotopes of an element differ only in their atomic mass.

Vocabulary

–

•

Atomic mass Isotopes and Atomic Masses

1. The mass of any given atom is not simply the sum of the masses of its electrons, protons, and neutrons.

2. Atoms are too small to measure individually and do not have a charge.

3. The arbitrary standard that has been established for describing atomic mass is the atomic mass unit (amu), defined as one-twelfth of the mass of one atom of 12 C.

4. Most elements exist as mixtures of several stable isotopes. The weighted average is of the masses of the isotopes is called the atomic mass. 5. Electrons added or removed from an atom produce a charged particle called an ion, whose charge is indicated by a superscript after the symbol for the element.

Vocabulary

–

Essential Elements

• Elements that are absolutely required in the diets of humans are called

essential elements (highlighted in purple).

• Essential elements are restricted to the first four rows of the periodic table

with only two exceptions (Mo and



).

• An essential element is one that is required for life and whose absence

results in death.

• An element is considered to be essential if a deficiency consistently causes

abnormal development or functioning and if dietary supplementation of that element and only that element prevents this adverse effect.

Classification of the Essential Elements

• Most living matter consists primarily of

diet.

bulk elements— oxygen, carbon, hydrogen, nitrogen, and sulfur. They are the building blocks of the compounds that make up our organs and muscles; they also constitute the bulk of our

• Six elements—sodium, magnesium, potassium, calcium,

chlorine, and phosphorus—are called structural components of the body.

macrominerals and provide essential ions in body fluids and form the major

• Remaining essential elements called

present in small amounts.

trace elements and are

The Trace Elements

• It is difficult to detect low levels of some of the essential

elements, so the trace elements were relatively slow to be recognized.

• Many compounds of trace elements are toxic. • Dietary intakes of elements range from deficient to optimum

to toxic with increasing quantities; the optimum levels differ greatly for the essential elements.

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Chemical Compounds

Atoms in all substances that contain more than one atom are held together by electrostatic interactions—interactions between electrically charged particles such as protons and electrons.

Compounds & Molecules

• COMPOUNDS

are a combination of 2 or more elements in definite ratios by mass. (Law of Definite Proportions)

• The character of each element is lost when forming a

compound.

• MOLECULES

are the smallest unit of a compound that retains the characteristics of the compound.

•

The composition of molecular compounds is given by a CHEMICAL FORMULA

WRITING FORMULAS

Chemical Formula –

chemical symbols and number of each representing composition

Empirical Formula –

simplest ratio of elements that does not represent the actual number and is non-positional

Molecular Formula – Structural Formula –

chemical symbols and number of each representing composition representing actual number but not position chemical symbols and number of each representing composition representing actual number and position

The structural chemical formula for glycine can be written as H 2 NCH 2 COOH to show atom ordering or in the bond-line structural formula H H O H N C H C O H Enables chemists to create a three-dimensional model that provides information about the physical and chemical properties of the compound

WRITING FORMULAS

Empirical and molecular formulas are precise and informative but have disadvantages: – inconvenient for routine verbal communications – many compounds have the same empirical and molecular formulas but different arrangements of atoms, which results in different chemical and physical properties

Compounds & Molecules

STRUCTURAL FORMULA BOND-LINE FORMULA

MOLECULAR FORMULAS

• In one molecule of glycine there are – 2 C atoms – 5 H atoms – 1 N atom – 2 O atoms

Empirical Formula _______________ Molecular Formula _______________ Structural Formula _______________



Representations of Molecular Structures

Different ways of representing the structure of a molecule

1.

1. Molecular formula gives only the number of each kind of atom present.

2. Structural formula shows which atoms are present and how they are connected.

3. Ball and stick model shows the atoms as spheres and the bonds as sticks.

4. A perspective drawing, called a attempts to show the three-dimensional structure of the molecule. wedge-and-dash representation, 5. The space-filling model shows the atoms in the molecule but not the bonds. 6. The condensed structural formula is the easiest and most common way to represent a molecule—it omits the lines representing bonds between atoms and simply lists the atoms bonded to a given atom next to it. Multiple groups attached to the same atom are shown in parentheses, followed by a subscript that indicates the number of such groups.

2.

3.

4.

5.

6.

•

Chemical Compounds Chemical bonds

– two different kinds 1. Ionic — ionic compounds consist of positively and negatively charged ions held together by strong electrostatic forces. (Formula Mass) 2. Covalent — covalent compounds consist of molecules, which are groups of atoms in which one or more pairs of electrons are shared between bonded atoms. Atoms are held together by the electrostatic attraction between the positively charged nuclei of the bonded atoms and the negatively charged electrons they share. (Molecular Mass)

• •

Ionic Chemical Compounds

Electrostatic attraction each other.

between oppositely charged particle species (positive and negative) results in a force that causes them to move toward Electrostatic repulsion (either both positive or both negative) results in a force that causes them to repel each other between two species that have the same charge

• When the attractive electrostatic interactions between atoms are stronger

than the repulsive interactions, atoms form chemical compounds attractive interactions between atoms are called chemical bonds .

and the

• • • • •

Ionic Chemical Compounds

IONS are atoms or groups of atoms with a net positive or negative charge. Taking away an electron from an atom gives a CATION with a positive charge Adding an electron to an atom gives an ANION with a negative charge .

Ionic compounds contain both cations and anions in a ratio that results in no net electrical charge.

CATION + ANION → COMPOUND A neutral compound requires equal number of + and - charges.

In general

metals (Mg) lose electrons to become

cations

nonmetals (F) gain electrons to become anions

Ionic compounds are held together by the attractive electrostatic interactions between cations and anions . Cations and anions are arranged in space to form an extended three-dimensional array that maximizes the number of attractive electrostatic interactions and minimizes the number of repulsive electrostatic interactions.



Physical Properties of Ionic Compounds

Ionic compounds – Usually form hard crystalline solids that melt at high temperatures and are very resistant to evaporation – Properties stem from the characteristic internal structure of an ionic solid, which is a three-dimensional array of alternating positive and negative ions held together by strong electrostatic attractions PbS SiO FeS



Binary Ionic Compounds

An ionic compound that contains only two elements, one present as a cation and one as an anion, is called a binary ionic compound.



For such compounds, the subscripts in the empirical formula can also be obtained using the absolute value of the charge on one ion as the subscript for the other ion and then reduce the subscripts to their simplest ratio to write the empirical formula.

Polyatomic Ions

   

Groups of atoms that bear a net electrical charge Atoms that make up a polyatomic atom are held together by the same covalent bonds that hold atoms together in molecules Many more kinds of polyatomic ions than monatomic ions and polyatomic anions are more numerous than polyatomic cations Method used to predict empirical formula for ionic compounds that contain monatomic ions can be used for compounds containing polyatomic ions. Overall charge on the cations must balance the overall charge on the anions in the formula unit.

Monovalent

Hydronium (or hydrogen) Lithium Sodium Potassium Rubidium Cesium Francium Silver Ammonium Thalium Copper I

Table of Common Ions Common Positive Ions (Cations)

H 3 O + H + Li + Na + K + Rb + Cs + Fr + Ag + NH 4 + Tl + Cu +

Divalent

Magnesium Calcium Strontium Beryllium Manganese II Barium Zinc Cadmium Nickel II Palladium II Platinum II Copper II Mercury II

Mercury I

Iron II Cobalt II Chromium II Lead II Tin II Mg 2+ Ca 2+ Sr 2+ Be 2+ Mn 2+ Ba 2+ Zn 2+ Cd 2+ Ni 2+ Pd 2+ Pt 2+ Cu 2+ Hg 2+

Hg 2 2+

Fe 2+ Co 2+ Cr 2+ Pb 2+ Sn 2+

Trivalent

Aluminium Antimony III Bismuth III Iron III Cobalt III Chromium III Al 3+ Sb 3+ Bi 3+ Fe 3+ Co 3+ Cr 3+

Table of Common Ions Common Negative Ions (Anions)

Monovalent

Hydride Fluoride Chloride Bromide Iodide Hydroxide Permangante Cyanide Thiocynate Acetate Nitrate Bisulfite Bisulfate Bicarbonate Dihydrogen phosphate Nitrite Amide Hypochlorite Chlorite Chlorate Perchlorate H Fl Cl Br I OH MnO 4 CN SCN C 2 H 3 O 2 NO 3 HSO 3 HSO 4 HCO 3 H 2 PO 4 NO 2 NH 2 ClO ClO 2 ClO 3 ClO 4 O 2-

O 2 2-

S 2 Se 2 C 2 O 4 2 CrO 4 2 Cr 2 O 7 2 WO 4 2 MoO 4 2 S 4 O 6 2 S 2 O 3 2 SO 3 2 SO 4 2 CO 3 2 HPO 4 2-

Trivalent

Nitride Phosphate N 3 PO 4 3-

  

Hydrates

Ionic compounds that contain specific ratios of loosely bound water molecules, called waters of hydration.

Waters of hydration can be removed by heating.

Compounds that differ only in the numbers of waters of hydration can have very different properties.

Prefixes for indicating numbers of species in chemical names

Prefix mono Number 1 Prefix hepta Number 7 di tri tetra penta hexa 2 3 4 5 6 octa nona deca undeca dodeca 8 9 10 11 12

Naming Polyatomic Ionic Compounds

Many compounds have more than one name: 1. Common name — have historical origins 2. Systematic name — write structure of the compound from its name and vice versa Procedure for naming binary ionic compounds, which contain only two elements, uses the following steps: 1) Place the ions in their proper order: cation and then anion 2) Name the cation 1) Metals that form only one kind of positive ion. These metals are in Groups 1–3, 12, and 13. The name of the cation of a metal that forms only one kind of positive ion is the same as the name of the metal 2) Metals that form more than one cation. These metals are transition metals, actinides, and the heaviest elements of Groups 13–15. Positive charge on the metal is indicated by a Roman numeral in parentheses following the name of the metal. 3) Name the anion 1) Monatomic anions — named by adding the suffix –ide to the root of the name of the parent element

2) Polyatomic anions

1) Have common names that must be learned 2) Polyatomic anions that contain a single metal or nonmetal atom plus one or more oxygen atoms are called oxoanions. Relationship between names of oxoanions and number of oxygen atoms present is:

Naming Polyatomic Ionic Compounds

Write the names of the compound as the name of the cation followed by the name of the anion.

– It is not necessary to indicate the number of cations or anions present per formula unit in the name of an ionic compound because information is implied by the charges on the ions – When writing the formula for an ionic compound from its name the charge of the ions must considered.

Covalent Compounds



Physical Properties of Covalent compounds

– Can be gases, liquids, or solids at room temperature and pressure, depending on the strength of the intermolecular interactions – Covalent molecular solids tend to form soft crystals that melt at low temperatures and evaporate easily – Consist of discrete molecules held together by comparatively weak intermolecular forces (the forces between molecules) even though the atoms within each molecule are held together by strong intramolecular covalent bonds (the forces within the molecule)

Covalent Compounds

• Covalent compound are represented by a

atomic symbol for each component element, in a prescribed order, accompanied by a subscript indicating the number of atoms of that element in the molecule molecular formula , which gives the C 4 H 10 OBCl 3 C 8 H 10 N 4 O 2 C 3 H 6 N 6 O 6

• •

Covalent Compounds Inorganic compounds

– Compounds that consist primarily of elements other than carbon and hydrogen – Include both covalent and ionic compounds – Formulas are written when the component elements are listed beginning with the one farthest to the left in the periodic table with those in the same group listed alphabetically

Organic compounds

– Covalent compounds that contain predominantly carbon and hydrogen – Formulas of organic compounds written with carbon first, followed by hydrogen and then by other elements in alphabetical order

Inorganic Covalent Compounds

• Some pure elements exist as covalent molecules • Hydrogen, nitrogen, oxygen, and the halogens

occur as diatomic molecules and contain two atoms

• A few pure elements, such as elemental phosphorus

and sulfur, are polyatomic molecules and contain more than two atoms

 

Binary Inorganic Compounds

Binary covalent compounds — covalent compounds that contain only two elements The procedure for naming them uses the following steps: 1. Place the elements in their proper order.

a. Element farthest to the left in the periodic table is named first. If both elements are in the same group, the element closer to the bottom of the column is named first.

b. Second element is named as if it were a monatomic anion in an ionic compound with the suffix – ide attached to the root of the element name 2. Identify the number of each type of atom present.

a. Prefixes derived from Greek stems are used to indicate the number of each type of atom in the formula unit.

b. If the molecule contains more than one atom of both elements, then prefixes are used for both. c. With some names, the final pronunciation.

a or o of the prefix is dropped to avoid awkward 3. Write the name of the compound.

a. Binary compounds of the elements with oxygen are named as “element oxide” with prefixes that indicate the number of atoms of each element per formula unit.

b. Certain compounds are always called by their common names assigned long ago when names rather than formulas were used.

Binary Inorganic Compounds



Organic Covalent compounds

Hydrocarbons and Biologically Important Compounds Hydrocarbons



Consist entirely of carbon and hydrogen



Four major classes of hydrocarbons: 1. Alkanes — contain only carbon-hydrogen and carbon carbon single bonds 2. Alkenes — contain at least one carbon-carbon double bond 3. Alkynes — contain a least one carbon-carbon triple bond 4. Aromatics — contain rings of six carbon atoms that can be drawn with alternating single and double bonds Biologically Important Compounds 1. Carbohydrates 2. Lipids 3. Biopolymers 1. Proteins 2. Nucleic Acids

To Be Covered in Weeks 6 and 7

Coordinate Inorganic Compounds

To Be Covered in Week 8

  

Acids and Bases

Acids – A substance with at least one hydrogen atom that can dissociate to form an anion and an H + ion (a proton) in aqueous solution, thereby forming an acidic solution Bases – Compounds that produce hydroxide ions (OH – ) and a cation when dissolved in water, thus forming a basic solution Neutral – Solutions that are neither basic nor acidic

To Be Covered in Weeks 9 and 10

Counting Atoms

Chemistry is a quantitative science—we need a “counting unit.”

MOLE -

1 mole is the amount of substance that contains as many particles (atoms, molecules) as there are in 12.0 g of 12C.

Avogadro’s Number

6.02214199 x 10

23

There is Avogadro’s number of particles in a mole of any substance.

Amedeo Avogadro 1776-1856

Molar Mass

1 mol of 12 C = 12.00 g of C = 6.022 x 10 23 atoms of C 12.00 g of 12 C is its MOLAR MASS Taking into account all of the isotopes of C, the molar mass of C is 12.011 g/mol

One-mole Amounts

Molar mass

•

Flowchart for converting between mass, number of moles, and number of atoms, molecules, or formula units

Molar mass

PROBLEM: What amount of Mg is represented by 0.200 g? How many atoms?

Mg has a molar mass of 24.3050 g/mol.

PROBLEM: How many hydrogen atoms in 0.036 moles of propane C 3 H 8 ? In 2.7 moles?

MOLECULAR MASS & MOLAR MASS

Molecular weight = sum of the atomic weights of all atoms in the molecule.

Molar mass = molecular weight in grams

Problem: What is the molar mass of ethanol – C

2

H

6

O?

1 mol contains

2 mol C (12.01 g C/1 mol C/1 mol C 2 H 6 O) = 6 mol H (1.01 g H/1 mol H/1 mol C 2 H 6 O) = 1 mol O (16.00 g O/1 mol O/1 mol C 2 H 6 O) = TOTAL = Molar Mass =

MOLECULAR MASS & MOLAR MASS

• Molecular mass of a substance

– Sum of the average masses of the atoms in one molecule of the substance – Calculated by summing the atomic masses of the elements in the substance, each multiplied by its subscript in the molecular formula – Units of molecular mass are atomic mass units (amu)

• Formula =

C 8 H 9 NO 2

• Molar mass =

Tylenol

Molar Mass

PROBLEM: How many moles of alcohol are there in a “standard” can of beer if there are 21.3 g of C 2 H 6 O?

(a) Molar mass of C 2 H 6 O = 46.08 g/mol (b) Calc. moles of alcohol

Molar Mass

PROBLEM: How many molecules of alcohol are there in a “standard” can of beer if there are 21.3 g of C 2 H 6 O?

(a) 21.3 g of C 2 H 6 O is 0.462 mol of C 2 H 6 O.

(b) Calc. molecules of alcohol

Molar Mass

PROBLEM: How many C 2 H 6 O?

atoms of C from ethanol are there in a “standard” can of beer if there are 21.3 g of (a) 21.3 g of C 2 H 6 O is 0.462 mol of C 2 H 6 O which is 2.78 E 23 .

(b) Calc. Atoms of carbon.

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Calculating Mass Percentages

A pure compound always consists of the same elements combined in the same proportions by weight.

Therefore, we can express molecular composition as PERCENT BY WEIGHT

Ethanol, C 2 H 6 O 52.13% C 13.15% H 34.72% O

Law of definite proportions states that a chemical compound always contains the same proportion of elements by mass Percent composition — the percentage of each element present in a pure substance—is constant Calculation of mass percentage 1. Use atomic masses to calculate the molar mass of the compound 2. Divide the mass of each element by the molar mass of the compound and then multiply by 100% to obtain percentages 3. To find the mass of an element contained in a given mass of the compound, multiply the mass of the compound by the mass percentage of that element expressed as a decimal

Percent Composition

Consider NO 2 , Molar mass = ?

What is the weight percent of N and of O?

What are the weight percentages of N and O in NO?

1.

Percent Composition

A compound of B and H is 81.10% B. What is its empirical formula?

Calculate the number of moles of each element in 100.0 g of sample.

2. Take the ratio of moles of B and H. Always divide by the smaller number. 3. Now, recognize that atoms combine in the ratio of small whole numbers.

4. Find the ratio of moles of elements in the compound.

We need to do an EXPERIMENT to find the MOLAR MASS. Experiment gives 53.3 g/mol Compare with the mass of B 2 H 5 = 26.66 g/unit , i.e. Find the ratio of these masses.

 

Determining the Empirical Formula of Penicillin

Can use the empirical formula of a substance to determine its percent composition Can use the percent composition of a sample to determine its empirical formula, which then can be used to determine the molecular formula—a procedure used to determine the empirical and molecular formulas of penicillin

Determining the Empirical and Molecular Formulas of Penicillin

The combustion analysis of Penicillin G C = 53.9% H = 4.8% N = 7.9% S = 6.5% Na = 6.5% Total = 82.1%

Chemical Equations

• A chemical equation is an expression that gives the identities and

quantities of the substances in a chemical reaction

• Chemical formulas and other symbols are used to indicate the starting

material(s) or reactant(s), which are written on the left side of the equation, and the final compound(s) or the reactants to the products.

product(s), which are written on the right side. An arrow, read as yields or reacts to form, points from

• • Abbreviations are added in parentheses as subscripts to indicate the

physical state of each species:—( and ( aq ) for an aqueous solution.

s ) for solid, ( l ) for liquid, ( g ) for gas, A balanced chemical equation number of atoms.

is when both the numbers of each type of atom and the total charge are the same on both sides. A chemical reaction represents a change in the distribution of atoms but not in the

Chemical Equations

• Look at the information an equation provides:

Fe 2 O 3(s) + 3 CO (g) 2 Fe (s) + 3 CO 2 (g)

• Balanced chemical equation

– Provides qualitative information about the identities and physical states of the reactants and products – Provides quantitative information because it tells the relative amounts of reactants and products consumed or produced in the reaction – The number of atoms, molecules, or formula units of a reactant or product in a balanced chemical equation is the coefficient of that species – Mole ratio of two substances in a chemical reaction is the ratio of their coefficients in the balanced chemical equation



Balancing Simple Chemical Equations

Method for balancing chemical equations 1. Identify the most complex substance.

2. Beginning with that substance, choose an element that appears in only one reactant and one product. Adjust the coefficients to obtain the same number of atoms of this element on both sides.

Optionally - Balance polyatomic ions (if present) as a unit.

3. Balance the remaining atoms, usually ending with the least-complex substance and using fractional coefficients if necessary. If a fractional coefficient is used, multiply both sides of the equation by the denominator to obtain whole numbers for the coefficients.

4. Count the numbers of atoms of each kind on both sides of the equation to be sure that the chemical equation is balanced.

Chemical Equations

• Law of Conservation of Matter – There is no detectable change in quantity of matter in

an ordinary chemical reaction.

– Balanced chemical equations must always include the

same number of each kind of atom on both sides of the equation.

– This law was determined by Antoine Lavoisier. • Propane,C

3 H 8 , burns in oxygen to give carbon dioxide and water.

D C 3 H 8  5 O 2 3 CO 2  4 H 2 O

Law of Conservation of Matter

• NH

3 burns in oxygen to form NO & water

Law of Conservation of Matter

• C

7 H 16 burns in oxygen to form carbon dioxide and water.

Balance this reaction

KClO

3

+ C

12

H

22

O

11

KCl + CO

2

+ H

2

O

Balance this reaction

CaSO

4

+ CH

4

+ CO

2

CaCO

3

+ S + H

2

O

Mass Relationships in Chemical Equations

• A balanced chemical equation gives the identity of the

reactants and products and the accurate number of molecules or moles of each that are consumed or produced.

•

Stoichiometry

is a collective term for the quantitative relationships between the masses, numbers of moles, and numbers of particles (atoms, molecules, and ions) of the reactants and products in a balanced reaction.

• A stoichiometric quantity is the amount of product or

reactant specified by the coefficients in a balanced chemical equation.

Stoichiometry Problems

• Steps in converting between masses of reactants and

product 1.

Convert the mass of one substance (Substance A) to the corresponding number of moles using its molar mass.

2. From the balanced chemical equation, obtain the number of moles of another substance (B) from the number of moles of substance A using the appropriate mole ratio (the ratio of their coefficients).

3. Convert the number of moles of substance B to mass using its molar mass.

• Converting amounts of substances to moles, and vice versa, is the key to all stoichiometry problems.

Calculations Based on Chemical Equations

• How many CO molecules are required to react with 25

MOLES of Fe 2 O 3 ?

Calculations Based on Chemical Equations

• How many iron atoms can be produced by the reaction of

2.50

E 5 MOLES of iron (III) oxide with excess carbon monoxide?

Calculations Based on Chemical Equations

• What mass of CO is required to react with 146 g of iron (III)

oxide?

Limiting Reactants

• If one or more of the reactants is not used up completely

but is left over when the reaction is completed, then the amount of product that can be obtained is limited by the amount of only one of the reactants

• A limiting reactant is the reactant that restricts the amount

of product obtained. The reactant that remains after a reaction has gone to completion is present in excess .

Limiting Reactant Concept

• What is the maximum mass of sulfur dioxide that can be

produced by the reaction of 95.6 g of carbon disulfide with 110. g of oxygen?

CS 2  3 O 2  CO 2  2 SO 2

Limiting Reactant Concept

• What is the maximum mass of sulfur dioxide that can be

produced by the reaction of 95.6 g of carbon disulfide with 110. g of oxygen?

Percent Yields from Reactions

• The

Limiting Reactant determines the maximum amount of product that can be formed from the reactants when reactants are not present in stoichiometric quantities.

• The

Theoretical Yield is calculated by assuming that the reaction goes to completion.

– Determined from the limiting reactant calculation. • Actual Yield

is the amount of a specified pure product made in a given reaction.

– In the laboratory, this is the amount of product that is

formed in your beaker, after it is purified and dried.

• Percent Yield

indicates how much of the product is obtained from a reaction.

actual yield % yield = theoretical yield  100%

Percent Yields from Reactions

• A 10.0 g sample of ethanol, C

2 H 5 OH, was boiled with excess acetic acid, CH 3 COOH, to produce 14.8 g of ethyl acetate, CH 3 COOC 2 H 5 . What is the percent yield?

Classifying Chemical Reactions



Most chemical reactions can be classified into one or more of only four basic types: 1. Acid-base reactions acid + base



2. Exchange reactions Metathesis) AB + C



(Single Displacement, Double Displacement, AC + B or AB + CD



AD + CB 3. Condensation reaction salt s (and the reverse, (Combination, Decomposition) Condensation: A + B



Cleavage: AB



A + B AB cleavage reactions ) 4. Oxidation-reduction reactions oxidant + reductant



reduced oxidant + oxidized reductant

To Be Covered in Week 9