Syllabus Chemistry 101 Fall 2010

Syllabus Chemistry 101 Fall 2010
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: (no voice mail) or leave a message at (put CHEM 101-Sec. # + subject in subject line of your ) 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. The lecture component of Chemistry 101 covers stoichiometry, atomic and molecular structure, chemical bonding, basic acid/base chemistry, solution chemistry, and the gas laws. In addition this class will cover foundation work in preparation for upper division course work in inorganic and organic nomenclature and structure, including discussion of major biological polymers; carbohydrates, sugars, fats, oils, and lipids, as well as peptides and proteins. More importantly as most in this class are pre-something, it is the goal of my lecture section to help you develop the skill set to not only successfully complete your undergraduate degree but to also prepare for professional school entrance exams and enable you to succeed in your choice of professional school. Check 12:40 section number

The instruction was applicable to my degree
All College 478 BIMS 128 27% 2014 379 Science 120 25% 2013 73 GEST and other 75 16% 2012 21 Agriculture other 42 9% 2011 5 Engineering 28 6% graduate Ag BICH, NUSC, GENE 22 5% Education 18 4% Liberal Arts Geosciences 11 2% 80% Pre-med, Pre-vet, Pre-Pharm, Pre-dent, Nursing, Physician Asst. Physical Therapist, Scientist, Eng. Architecture 9 Business and other 7 1% 20% degree requirement and thought this would be a fun class 501 505 228 250 33% 103 41% 39 17% 45 18% 25 11% 36 14% 20 23 19 8% 16 7% 8 3% 10 6 4 2 3 The instruction was applicable to my degree update

Kotz and Treichel 6th or 7th ed. TEXTBOOKS Chang
New Hardbound ~$200 paperback ~ $150 Solution Manual ~ $60 Online Tutor ~ $45 Total ~$ Ebook $45 per semester Includes Text Solution manual Online tutorial Optional Used Texts $5 - $50 + shipping no sell back Helpful 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. Change pearson ed rep name Chang’s Essentials

Supplemental Instruction Courses Texas Success Initiative
Tutoring Supplemental Instruction Courses Texas Success Initiative About Us Contact Us 118 Hotard North of Sbisa, between Neeley Hall and the Northside Post Office (979) The Student Learning Center has won the 2008 National College Learning Center Association Frank L. Christ Outstanding Learning Center Award! The award recognizes the center's commitment to supporting and strengthening the Academic experience of students at Texas A&M University by providing a variety of programs and services that promote retention and success. Read more... The Student Learning Center provides Supplemental Instruction and tutoring free of charge to all Texas A&M University students. The SLC oversees the STLC courses (formerly CAEN), which teach students how to improve their study skills and prepare for the job market. The SLC manages Developmental programs for students who have not yet passed the assessment tests required by the state. Study Tips General Time Management Reading Textbooks Setting Goals Preparing for Exams Success Tips from Fellow Aggies During the Fall 2008 semester, drop-in tutoring will be offered Sunday nights 5-8pm and Monday through Thursday nights from 5-10pm. Tutoring will begin on Monday, September 1st. Tutor Zones are currently planned for Studio 12 of The Commons. Look for our table and tutors and just ask for help! See the schedule Drop-in tutoring is available for most lower level math and science courses on the first floor of Hotard Hall. Tutors are also available to help out with many other courses. If you need help in a particular course and would like to check to see if a tutor is available for that course, you can contact our tutor coordinator, Linda Callen, at

Grading: Your grade will be based on
Four one-hour examinations (each worth 200 points) A final examination (400 points) There are no bonuses or soft points There will be no extra credit Major Examination Schedule Fall 2010: Fri. Sept. 17 Major Exam No.1 Wed. Oct. 6 Major Exam No.2 Mon. Nov. 1 Major Exam No.3 Fri. Dec. 3 Major Exam No. 4 Fri. Dec. 10 Section 501 Final Exam 10:00 to 12:00 Mon. Dec. 13 Section 505 Final Exam 10:30 to 12:30 Check section numbers

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 Sum of points assigned to correct responses Small dataset is bootstrapped to A context-free evaluation of relative performance 500,000+ exams 3) Individual scores are normalized. 4) Normalized scores are transformed An absolute score is assigned to a defined scale 5) Letter grades are assigned > A > D > B < F > C

Results of Bootstrap Overall T-score A B C D F 501 15.5% 37.0% 38.8%
6.8% 1.8% 503 13.3% 37.6% 37.3% 10.6% 1.2% All 14.3% 38.0% 8.9% 1.5% 16.0% 35.0% 14.0% 2.0% Exam 1 18.5% 33.8% 30.2% 15.8% 17.6% 38.3% 30.1% 10.2% 3.9% 18.0% 36.2% 12.8% 2.9% Exam 2 20.1% 29.7% 34.7% 2.7% 16.5% 39.2% 27.5% 3.5% 18.1% 34.8% 30.8% 13.1% 3.2% Exam 3 17.4% 33.5% 0.0% 17.0% 31.6% 36.4% 14.6% 0.4% 17.2% 32.5% 36.3% 13.8% 0.2% Exam 4 19.5% 34.9% 32.6% 9.8% 3.3% 19.3% 26.5% 13.7% 19.4% 29.3% 11.9% Final 19.8% 29.5% 35.9% 0.5% 18.7% 31.3% 35.7% 19.2% 30.5% 35.8% 13.6% 0.9% Results of Bootstrap Simulation Stats 5000 Repetitions 549 Secs Sheet1!$B$112 Sheet1!$B$113 Average 35.319 SD 3.3410 1.8824 Max 41.707 Min 29.000 Form Diff Raw T Grade Exam 1 Student X A 1 146.00 89.19 B Student Y C 4 138.00 85.08 Exam 2 2 130.00 84.50 131.00 80.41 Exam 3 75.00 77.61 87.00 79.74 Exam 4 25.00 57.83 F D 108.00 79.57 Final Exam 234.00 84.54 3 230.00 76.62 Avg Exam Diff 1.4 3.8 Avg Raw 50.8% 57.8% Avg T Score 79.7% Course Grade

Maximum Cumulative Points after Last Exam Taken
current avg grade desired avg grade extra raw score points needed for desired grade avg raw score needed on every remaining exam to earn desired grade 50.000% 79.501% 200 100 400 318 118 170.80 85.401% 600 300 477 177 188.50 94.252% 800 636 236 218.00 % 1200 900 954 54 277.01 % 90.000% 180 360 -42 154.80 77.401% 540 -63 148.50 74.252% 720 -84 138.00 69.002% 1080 -126 117.01 58.503%

And we begin: 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 Problem - A situation that presents difficulty, uncertainty, or perplexity: 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.

Numbers – Significant Figures, Rounding Rules, Accuracy, Precision, Statistical Treatment of the Data Units – 5 of the 7 Time – seconds Conversion factors – realtionship Length – Meters between two of the basic units Mass – grams Molecular Weight (Mass) Amount – Moles Molar Mass, Mole Ratio, Molarity, Temperature – Kelvins molality, Density? Vocabulary – Approximately 300 new terms or words and applying new or more rigid definitions to words you may already own. Principles (Theories and Laws) – Stoichiometry, Quantum Theory, Bonding, Chemical Periodicity, Solutions, Thermodynamics, Intermolecular Forces, Gas Laws, Colligative Properties, Kinetics, Equilibrium, Electrochemistry cp = q/mDT rate = k[A]m[B]n ∆E = q + w DG = DH – TDS Eocell = Ecathode = Eanode PV = nRT %yield = actual/theoretical * 100% K = DT = Kmi [C]c[D]d [A]a[D]b l (m) c (ms-1) E = n =

Dr. Stone’s patent pending chemistry problem solver
Write down everything you are given Vocabulary Numbers Units Write down what you want to know Write down mathematical equation(s) that include(s) these values and units Principles Write a balanced stoichiometric equation Mole concept Convert everything to moles Dimensional analysis Convert everything to the unknown’s units Rounding, significant figures, accuracy and precision

Observations Hypotheses Experiments Law Theory
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. 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 Theory

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

Natural Laws Law of Conservation of Mass – The notion that mass, or matter, can be neither created nor destroyed. 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=mc2 – 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.

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

[ ] Use of Numbers How many sig figs? How many sig figs in the answer?
1.4800 100 100.0 89400 74.000 How many sig figs in the answer? 472x101 4600x0.005 36.0x4752 45.08/36.2 1.003/8500 0.003/472x12 3.003/475.0x0.30/524 0.3005x4.1 Round off to two sig figs 34.78 17.51 48.50 24.33 17.50 20.5 Add clicker questions (1.68) [ ] 23.56 – 2.3 1.248 x 103 = Day 2

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.

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 the same number of protons, and hence the same atomic number, but 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.

1. The mass of any given atom is not simply the sum of the masses of its electrons, protons, and neutrons. 2. Most elements exist as mixtures of several stable isotopes. The weighted average is of the masses of the isotopes is called the atomic mass. 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 12C. 4. 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 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 diet. Six elements—sodium, magnesium, potassium, calcium, chlorine, and phosphorus—are called macrominerals and provide essential ions in body fluids and form the major structural components of the body. Remaining essential elements called trace elements and are present in small amounts (V, Cr, Ni, F, Sn, Si, As).

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.

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 – chemical symbols and number of each representing composition representing actual number but not position Structural Formula – chemical symbols and number of each representing composition representing actual number and position The structural chemical formula for glycine can be written as H2NCH2COOH to show atom ordering or in the bond-line structural formula 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 Day 3

In one molecule of glycine there are 2 C atoms 5 H atoms 1 N atom
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 _______________ Clicker question?

Representations of Molecular Structures
Different ways of representing the structure of a molecule Molecular formula gives only the number of each kind of atom present. Structural formula shows which atoms are present and how they are connected. Ball and stick model shows the atoms as spheres and the bonds as sticks. A perspective drawing, called a wedge-and-dash representation, attempts to show the three-dimensional structure of the molecule. The space-filling model shows the atoms in the molecule but not the bonds. 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. 1. 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 between oppositely charged particle species (positive and negative) results in a force that causes them to move toward each other. Electrostatic repulsion between two species that have the same charge (either both positive or both negative) results in a force that causes them to repel each other When the attractive electrostatic interactions between atoms are stronger than the repulsive interactions, atoms form chemical compounds and the attractive interactions between atoms are called chemical bonds.

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
– 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 SiO PbS 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.

Common Positive Ions (Cations)
Table of Common Ions Common Positive Ions (Cations) Monovalent Divalent Trivalent Hydronium H3O+ magnesium Mg2+ aluminium Al3+ (or hydrogen) H+ calcium Ca2+ Antimony III Sb3+ lithium Li+ strontium Sr2+ Bismuth III Bi3+ Sodium Na+ Beryllium Be2+ Potassium K+ manganese II Mn2+ Rubidium Rb+ barium Ba2+ Cesium Cs+ zinc Zn2+ Francium Fr+ Cadmium Cd2+ Silver Ag+ Nickel II Ni2+ Ammonium NH4+ Palladium II Pd2+ Thalium Tl+ Platinum II Pt2+ copper I Cu+ copper II Cu2+ Mercury II Hg2+ mercury I Hg22+ iron II Fe2+ iron III Fe3+ Cobalt II Co2+ Cobalt III Co3+ chromium II Cr2+ chromium III Cr3+ lead II Pb2+ tin II Sn2+

Common Negative Ions (Anions)
Table of Common Ions Common Negative Ions (Anions) Monovalent Divalent Trivalent Hydride H- Oxide O2- Nitride N3- Fluoride Fl- Peroxide O22- Chloride Cl- Sulfide S2- Bromide Br- Selenide Se2- Iodide I- Oxalate C2O42- Hydroxide OH- Chromate CrO42- Permangante MnO4- Dichromate Cr2O72- Cyanide CN- Tungstate WO42- Thiocynate SCN- molybdate MoO42- Acetate C2H3O2- tetrathionate S4O62- Nitrate NO3- Thiosulfate S2O32- Bisulfite HSO3- Sulfite SO32- Bisulfate HSO4- Sulfate SO42- Bicarbonate HCO3- Carbonate CO32- Dihydrogen phosphate H2PO4- Hydrogen phosphate HPO42- Phosphate PO43- Nitrite NO2- phosphite Amide NH2- Hypochlorite ClO- Chlorite ClO2- Chlorate ClO3- Perchlorate ClO4-

Prefixes for indicating numbers of species in chemical names
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. CuSO4.5H2O Prefixes for indicating numbers of species in chemical names Prefix Number mono - 1 hepta- 7 di- 2 octa- 8 tri- 3 nona- 9 tetra- 4 deca- 10 penta- 5 undeca- 11 hexa- 6 dodeca- 12 Day 4

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: Place the ions in their proper order: cation and then anion Name the cation 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 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. Name the anion Monatomic anions — named by adding the suffix –ide to the root of the name of the parent element Polyatomic anions Have common names that must be learned 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 C8H10N4O2 C4H10OBCl3 C3H6N6O6
Covalent compound are represented by a molecular formula, which gives the 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 C4H10OBCl3 C8H10N4O2 C3H6N6O6

Covalent Compounds Inorganic compounds Organic 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. 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. 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 a or o of the prefix is dropped to avoid awkward pronunciation. Write the name of the compound. 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

To Be Covered in Weeks 6 and 7
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 Aromatics — contain rings of six carbon atoms that can be drawn with alternating single and double bonds To Be Covered in Weeks 6 and 7 Biologically Important Compounds Carbohydrates –Sugars, starches Lipids, Fats, Oils Biopolymers Proteins Nucleic Acids To Be Covered in Week 8

Coordinate Inorganic Compounds
To Be Covered in Week 9

Acids and Bases To Be Covered in Weeks 10 and 11 Acids Bases Neutral
– 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 10 and 11

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 x 1023 There is Avogadro’s number of particles in a mole of any substance. 518 g of Pb, 2.50 mol Amedeo Avogadro

Molar Mass One-mole Amounts
1 mol of 12C = g of C = x 1023 atoms of C 12.00 g of 12C is its MOLAR MASS Taking into account all of the isotopes of C, the molar mass of C is g/mol One-mole Amounts Day 5

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

PROBLEM: What amount of Mg is represented by 0.200 g? How many atoms?
Molar mass PROBLEM: What amount of Mg is represented by g? How many atoms? Mg has a molar mass of g/mol. PROBLEM: How many hydrogen atoms in moles of propane C3H8? In 2.7 moles? Clicker question part two

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 – C2H6O? 1 mol contains 2 mol C (12.01 g C/1 mol C/1 mol C2H6O) = 6 mol H (1.01 g H/1 mol H/1 mol C2H6O) = 1 mol O (16.00 g O/1 mol O/1 mol C2H6O) = 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 = Molar mass = C8H9NO2 Clicker question Tylenol

Molar Mass PROBLEM: How many moles of alcohol are there in a “standard” can of beer if there are 21.3 g of C2H6O? (a) Molar mass of C2H6O = g/mol (b) Calc. moles of alcohol PROBLEM: How many molecules of alcohol are there in a “standard” can of beer if there are 21.3 g of C2H6O? (a) g of C2H6O is mol of C2H6O. (b) Calc. molecules of alcohol PROBLEM: How many atoms of C from ethanol are there in a “standard” can of beer if there are 21.3 g of C2H6O? (a) g of C2H6O is mol of C2H6O which is 2.78 E23 . (b) Calc. Atoms of carbon.

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, C2H6O 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 NO2, Molar mass = ?
What is the weight percent of N and of O? What are the weight percentages of N and O in NO?

A compound of B and H is 81.10% B. What is its empirical formula?
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 g of sample. Take the ratio of moles of B and H. Always divide by the smaller number. Now, recognize that atoms combine in the ratio of small whole numbers. 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 B2H5 = 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% Day 6

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 product(s), which are written on the right side. An arrow, read as yields or reacts to form, points from the reactants to the products. Abbreviations are added in parentheses as subscripts to indicate the physical state of each species:—(s) for solid, (l) for liquid, (g) for gas, and (aq) for an aqueous solution. A balanced chemical equation 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 number of atoms.

Chemical Equations Look at the information an equation provides:
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 Identify the most complex substance. 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. 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. 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,C3H8, burns in oxygen to give carbon dioxide and water.

Law of Conservation of Matter
NH3 burns in oxygen to form NO & water C7H16 burns in oxygen to form carbon dioxide and water. KClO C12H22O KCl + CO H2O CaSO4 + CH4 + CO CaCO3 + S + H2O

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 Convert the mass of one substance (Substance A) to the corresponding number of moles using its molar mass. 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). 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 Fe2O3? How many iron atoms can be produced by the reaction of 2.50E5 MOLES of iron (III) oxide with excess carbon monoxide? 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? 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.

Percent Yields from Reactions
A 10.0 g sample of ethanol, C2H5OH, was boiled with excess acetic acid, CH3COOH, to produce 14.8 g of ethyl acetate, CH3COOC2H5. What is the percent yield?

Classifying Chemical Reactions To Be Covered in Weeks 10 and 11
Most chemical reactions can be classified into one or more of only four basic types: 1. Acid-base reactions acid + base  salt 2. Exchange reactions (Single Displacement, Double Displacement, Metathesis) AB + C  AC + B or AB + CD  AD + CB 3. Condensation reactions (and the reverse, cleavage reactions) (Combination, Decomposition) Condensation: A + B  AB Cleavage: AB  A + B 4. Oxidation-reduction reactions (covered in Chem 102) oxidant + reductant  reduced oxidant + oxidized reductant To Be Covered in Weeks 10 and 11 Day 7

Syllabus Chemistry 101 Fall 2010

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Syllabus Chemistry 101 Fall 2010

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