CHEMISTRY 101



CHEMISTRY 101

Sections 578 - 588

T, R 11:10 to 12:25 p.m.

Room 100 Heldenfels Hall

Texas A&M University

M. L. Peck

413 Heldenfels Hall

Telephone: 845-2356

e-mail: Peck@tamu.edu

Textbook: General Chemistry (sixth edition)

by: Whitten, Davis, and Peck

Saunders College Publishing Ó 2000

OFFICE HOURS:

Tue. and Thurs. 12:30-2:00

Several additional hours are

available each week by appointment

HOMEWORK

SET 1 Ð due 9/30

Chapter 1 -- 3, 8, 9, 13, 23, 35, 54, 58, 73 and 77

Chapter 2 -- 4, 6, 16, 18, 20, 26, 32, 34, 44, 46, 48, 56, 62, 70, 75, 94, 96, 101 and 108

Chapters 3, 4 and 5 (will be announced at a later date)

CHAPTER 1 The Foundations of CHEMISTRY

GOALS: Structure of matter (states of matter, chemical and physical properties and changes, metric system, significant; figures, solving problems, density, heat transfer, and synthesis problems).

Define Chemistry, Matter, Energy, Elements, Compounds, Mixtures, Atoms and Molecules. You will need to know the names and symbols of the elements in Table 1-2 on page 16 of the textbook. ThatÕs not an easy task if you donÕt already know most of them. If you donÕt know them, I suggest that you make flashcards with the names on one side and the symbols on the other. For the first test, you will also need to know the names and symbols of the following additional elements: argon/Ar arsenic/As cesium/Cs germanium/Ge radium/Ra selenium/Se xenon/Xe

You will always have the Periodic Table during exams (itÕs one that we provide) so you may find it easier to learn the names if you look at a Periodic Table that has the symbols, etc., but not the names. The Periodic Tables that we provide for use during exams are very similar to the one on the front inside cover of your textbook or your lab manual.

Monatomic MoleculesÑ elements (He, Ne, Ar, Kr, Xe)

Diatomic MoleculesÑ elements [H2, N2, O2, F2, Cl2, Br2 (l or g), and I2 (g, l or s)]

compounds (NaCl, CO, KBr)

More Complex MoleculesÑ elements (carbon, P4, S8)

compounds [CH3CH2OH (or C2H5OH), CH4, CO2 and Na2SO4]

Law of Conservation of MatterÑthere is no detectable change in the quantity of matter during a chemical reaction or during a physical change.

Law of Conservation of EnergyÑenergy cannot be created or destroyed in a chemical reaction or in a physical change. It can only be changed from one form to another.

Law of Conservation of Matter and EnergyÑthe combined amount of matter and energy in the universe is fixed. (E = mc2)

A Scientific (Natural) LawÑa statement, based on experiments (observations), that is believed to be true and to which exceptions are not known.

Chemical PropertiesÑproperties that matter exhibits as it undergoes change in composition (chemical change).

States of MatterÑsolid, liquid and gas

Physical PropertiesÑcan be observed in the absence of any change in composition.

Examples: Chemical Properties Physical Properties

reactivity (reacts with water, etc.) change in state (melt, boil, etc.)

reactivity(unstable in air) no change in composition (soft)

Extensive and Intensive Properties (Extensive: depends upon amount.

Intensive: Independent of sample size)

1.6 Measurements in Chemistry

(You will need to know the prefixes kilo-, centi-, and milli-.)

7. Units of Measurement (On a test, I will not ask you to convert from English to Metric or vice versa. We will work many problems that involve conversions within the metric system.)

1.12 Heat and Temperature (On a test, I may ask you to convert from Fahrenheit to Celsius or vice versa. You will often need to convert between Celcius and Kelvin.)

Significant Figures (More for laboratory than for lecture)

Dimensional Analysis Factor-label, Unit factor, and Ratio methods will be used to work various examples in lecture. You should use the method that you understand best.

REVIEW If 80.0 grams of a corrosive, volatile liquid are needed for a reaction, what volume of liquid would you use? The density of the liquid is 1.48 g×mL-1.

A 37.74 gram piece of copper is dropped into a graduated cylinder that contains 25.00 mL of water. After the addition of the metal, the water level rises to 29.23 mL. What is the specific gravity of copper?

Calculate the amount of heat required to raise the temperature of 150.0 grams of water from 20.0¡C to 60.0¡C. Specific heat of water equals 4.18 J/g¡C.

Answers: 54.1mL, 8.92, & 25.1 kJ

We will come back to temperature and heat calculations much later in this course.

Key Terms are listed at the end of chapters. It is important that you understand each term. Again, I suggest that you prepare flashcards.

SAMPLE TEST QUESTIONS

Which one of the following statements is FALSE?

(a) The term ÒmoleculeÓ can refer to elements.

(b) The term ÒmoleculeÓ can refer to compounds.

(c) The term ÒatomÓ can refer to elements.

(d) The term ÒatomÓ can refer to compounds.

(e) Molecules can be monatomic, diatomic, or polyatomic.

Answer (d)

HOMEWORK

SET 1 Ð due 9/30

Chapter 1 -- 3, 8, 9, 13, 23, 35, 54, 58, 73 and 77

Chapter 2 -- 4, 6, 16, 18, 20, 26, 32, 34, 44, 46, 48, 56, 62, 70, 75, 94, 96, 101 and 108

Chapter 3 -- 4, 6, 8, 10, 12, 14, 22, 34, 36, 38, 40, 42, 56, 60, 62, 70, 72, 80, 98, 100, 101, and 104

Chapters 4 and 5 (will be announced at a later date

OFFICE HOURS

Tuesdays and Thursdays 12:30 Ð 1:30

CHAPTER 2

CHEMICAL FORMULAS AND COMPOSITION STOICHIOMETRY

Recall examples of Monatomic Molecules (elements) He, Ne, Ar, Kr

Recall examples of Diatomic Molecules (elements) O2, H2, Cl2, F2, Br2, I2, N2

Recall examples of Complex Molecules (elements) P4, O3, S8

Formula of a substance must show its composition.

Allotropic modifications (allotropes) are different forms of the same element.

Examples are O2 and O3, diamond and graphite, and white and red phosphorus.

The mass of an atom of Carbon-12 is exactly 12 atomic mass units (amu's).

Carbon-12, Carbon-13, and Carbon-14 are isotopes of the element, carbon.

The weight on the periodic table (atomic weight) is an average based upon a natural sample.

CarbonÕs atomic wt. is 12.011. 12.011 amuÕs is the average weight of one atom of carbon.

6.022 x 1023 formula units (AvogadroÕs Number) of a substance is called a mole.

There are 6.022 x 1023 amu's per gram and there are also 6.022 x 1023 formula units per mole.

EXAM Practice: Calculate the average mass of a magnesium atom in grams.

Calculate the number of atoms in one-millionth of a gram of magnesium metal.

How many atoms are contained in 0.325 mole of magnesium?

How many moles of calcium are present in 94.1 grams of calcium?

Answers: 4.037 x 10-23 g atom-1, 2.477 x 1016 atoms, 1.96 x 1023 atoms, 2.35 mol

THE LAW OF DEFINITE PROPORTIONS

(Constant Composition) is based on the observation that ÒDifferent samples of a pure compound always contain the same elements in the same proportions by mass.Ó

The formula of a substance indicates its chemical composition (molecule, if molecular; ion ratio, if ionic). Cation - positive ion Anion - negative ion

Introduction to Naming Compounds

You will need to know the names and formulas given in Table 2-2 on page 51 of your textbook, Figure 2-6 on page 53, and Table 2-3 on page 55. We will add to this list as we go through this course. Again, I suggest flashcards.

Formula weight (mass) of a substance is the sum of the atomic weights of the elements in the formula unit, each taken the number of times the element occurs. A formula weight is also the molecular weight if the compound exists as molecules and not as ions.

The formula weight of butane, (C4H10) is:

The formula weight of ammonium sulfate, (NH4)2SO4 is:

One mole of a substance is 6.022 x 1023 formula units (formulas or molecules) of the substance and it has a definite mass as shown by the formula of the substance.

One mole of butane is grams of C4H10.

One mole of ammonium sulfate is grams.

Calculate the number of C4H10 molecules in 75.0 g of butane.

Calculate the number of oxygen atoms in 56.2 grams of magnesium perchlorate, Mg(ClO4)2.

What percentage by mass of aluminum sulfate, Al2(SO4)3, is aluminum?

Answers: 58.123(or 58.0 since in this lecture course, in our calculations, we usually use atomic weights expressed to only their first decimal place), 132.0, 58.0 g, 132.0 g, 7.8 x 1023, 1.2 x 1024, & 15.8%

* A millimole (mmol) is 1/1000 of a mole.

Nonstoichiometric compounds do not obey the Law of Definite Proportions

(Ni0.97O, TiO1.7-1.8, Cu1.7S)

SIMPLEST (OR EMPIRICAL FORMULAS)

Example: A compound contains 42.1% sodium, 18.9% phosphorus, and 39.0% oxygen by mass. What is its simplest (or empirical) formula?

Steps followed in obtaining formula from percent composition:

1. Find number of moles of each element in a given sample (use 100g if no sample mass is given).

2. Divide the number of moles of each by the smallest number of moles.

3. Multiply by the smallest whole number that will convert each of the molar amounts into a whole number.

Example: A compound contains 85.6% carbon and 14.4% hydrogen by mass. If its molecular weight is 42, its formula is:

REMEMBER!!! EMPIRICAL FORMULA - smallest whole number ratio of atoms in a compound. Molecular formula will be the empirical formula multiplied by a whole number (perhaps by one).

Determination of Molecular Formulas

Key Terms are listed at the end of chapters. It is important that you understand each term. Again, I suggest that you prepare flash cards.

SAMPLE TEST QUESTIONS

How many moles of nitrogen atoms are there in 1.27 moles of (NH4)2SO4?

(a) 1.27 mol (b) 2.54 mol (c) 7.60 x 1023 mol (d) 167 mol (e) 11.16 mol

How many molecules are in 1.0 kg of H2O?

(a) 3.3 x 1025 (b) 1.1 x 1025 (c) 8.7 x 1023 (d) 1.8 x 1027 (e) 5.6 x 1023

A compound is 50.0% S and 50.0% O by mass. What is the simplest formula of the compound?

(a) SO (b) SO3 (c) SO4 (d) SO2 (e) SO42-

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