Important Things to Note

1. Tardiness is considered an absence.

2. You are given 4 free absences before deductions start accruing. Attend class ALL the time.

3. 2 In-class Exams (one of which is the midterm) 25%

4. 1 Lab Exam close to the end of the semester (5%)

5. 1 COMPREHENSIVE Final Exam (25%)

6. 12 Quizzes ( 6 before midterms and 6 after) (25%) – lowest 2 quizzes per term are dropped (total of 4)

7. No make-ups on quizzes because of the drop rule.

8. You fail the lab portion, you fail the course.

9. Bring a scientific calculator to class EVERYDAY!

10.Engage/participate in class. You take out what you put in it.

11.Do NOT CHEAT on quizzes, exams, labs etc…

12. Treat everyone with respect!

13. I teach FAST. Keep up and study!

14.If you get offended easily, please take the class with a different teacher.

15.My class, my rules. I run it the way I see fit. Do not battle with me or give me attitude!

NOTE For Lab

• We MEET first in IC 419 for the pre-lab quiz and pre-lab instructions before performing the lab in IC 416.

• Bring lab notebook, lab manual, data sheet(s), calculator, apron and lab goggles.

Important Dates (Sections A and B)

• Exam 1 - Tuesday, September 26

• Exam 2 - Tuesday, November 7

• Lab Exam – Tuesday, November 14

• Final Exam – Monday, December 4 @ 10:15 a.m. – 12:15 p.m. in IC 111

Important Dates (Sections C and D)

• Exam 1 - Thursday, September 28

• Exam 2 - Thursday, November 9

• Lab Exam – Thursday, November 16

• Final Exam – Tuesday, December 5 @ 8:00 a.m. – 10:00 a.m. in IC 111

Chapter 1

The Study of Chemistry

Topics

• Introduction• Scientific Method• Classifications of Matter• Properties of Matter• Units of Measurement – Metric system• Temperature Conversion• Metric Conversion (Prefixes)• Accuracy vs. Precision• Significant Figures• Density

States of Matter

• Solid

• Liquid

• Gas

• Plasma

• Bose-Einstein Condensate

ATOM

• Is the simplest unit of matter.

Definitions

• Elements – can’t be decomposed further into simpler substances

- 118 elements presently - Lv (element 116 -

livermorium)• - Fl (element 114 - flerovium)• Compound – combination of 2 or more

elements• Pure substance – has distinct properties and

composition; does not vary from sample to sample (ex. Water, NaCl)

Definitions

• Mixtures – combinations of 2 or more substances (ex. sugar in water)

• 2 Types of Mixtures

• 1. Homogenous Mixtures (solutions) = 1 phase

• 2. Heterogeneous Mixtures = > 2 phases

SOLUTIONS

• Homogeneous mixtures are called SOLUTIONS.

Solution

• Solution – homogenous mixture

• A solution is not necessarily a liquid. Can be gas or solid.

Physical vs. Chemical Properties

• Physical properties – can be measured w/o changing identity and composition of substance (ex. Boiling pt.,freezing pt., color, odor, density, hardness)

• Chemical properties – describe how substance reacts or changes to form other compounds (ex. Flammability, toxicity)

Changes of State and Properties

• Physical changes – does not change composition of compound

• Chemical changes – converts to a different chemical substance

• Intensive Properties – independent of amt. (ex. Density, Temperature, Melting Pt)

• Extensive Properties – dependent on amt. (ex. Mass, Volume)

Units of Measurement

• Mass – grams; kilogram• Length – centimeter; meter• Volume – milliliter or cubic

centimeter (cm3)• Temperature – Celcius;

Kelvin

Temperature Conversions

• 0 oC = 273.15 K

• oF = 1.8 oC + 32

Things to Remember!

• 1 milliliter = 1 cc

• 1000 milliliter = 1 liter

• 0 oC = 32 oF = 273.15 K

Precision vs Accuracy

• Accuracy – when acquired value agrees with true value

• Precision – when acquired values exhibit reproducibility

Significant Figures

• More significant figures = more certainty

• Helps in determining how to round measured values and still precise

SIGNIFICANT FIGURES

• In counting and definitions, there are an infinite number of sig figs

• In measurements, the number of sig figs consists of all certain and the first uncertain digits

• Unit conversions do not determine # of sig. figs.

Rules of Significant Figures

• 1. Non-zero integers always count.

• Ex. 1234.5 grams = 5 Sig. Figs.

• 2. Captive zeros are always significant.

• Ex. 100.3 grams = 4 Sig. Figs.

Rules of Significant Figures

• 3. Leading zeros are NEVER significant.

• Ex. 0.6780 grams = 4 Sig. Figs.

• 4. Trailing zeroes are significant ONLY if there is a decimal point

• Ex. 12.0 grams = 3 Sig. Figs• 120 grams = 2 Sig. Figs

Rules of Significant Figures

• 5. Exact numbers (obtained by counting) are infinite and do not determine the number of significant figures.

• Example: 4 cows = ?

Determine the # of Sig. Fig.

• 200.0• 1050• 3003• 0.0006• 10,000• 0.5

Rules of Significant Figures

• Multiplication/Division

–Answer will have the same # of sig figs as the value with the least # of sig figs• Ex: 3.8 x 200.0 = 2 Sig. Figs.

Rules of Significant Figures

• Addition/Subtraction

• Answer has the same # of decimal places as the number with the least # of decimal places

• Ex. 3.1 + 2.500 + 5.76 = 11.4

Order of Operations

• Parenthesis

• Multiplication/division

• Addition/subtraction

Rounding

• Look only to the right of the number you are rounding to:

• - If 5 or more, round up

• - If less than 5, round down

General Rule

• Carry ALL figures through to the end of a problem while keeping track of significant figures along the way. Round the final answer to the correct number of significant figures

Problem

• Indicate the number of sig. figs. in each of the following measured quantities:

• A. 358 kg

• B. 0.054 s

• C. 6.3050 cm

• D. 0.0105 L

• E. 7.0500 x 10-3 m3

Problem

• Round each of the following numbers to 4 sig. figs. And express the result in standard exponential or scientific notation.

• A. 102. 53070• B. 656, 980• C. 0.008543210• D. 0.000257870• E. - 0. 0357202

Problem

• Carry out the following operations and express the answer with the appropriate number of sig. figs.

• A. 12.0550 + 9.05

• B. 257.2 – 19.789

• C. (6.21 x 103)(1.1050)

• D. 0.0577 / 0.753

Prefixes in Metric System

• Mega - million• Kilo - 1,000• Hecto - 100• Deka - 10• ----- - 1 (liter, gram, meter)• Deci - 1/10 or 0.1• Centi - 1/100 or 0.01• Milli - 1/1000 or 0.001

Trick to Unit Conversions

• Arrange the values with units such that you can cancel what you do not want.

• You should only end up with the desired unit.

Tricks to Conversion

• Big to small - multiply– going down

• Small to Big - divide–Going up

Density

• Is the amount of mass in a unit volume of the substance

• Is affected by Temperature. – The higher the temp., the lower the density.

D = mass of substance = grams volume of substance mL or cm3

Density

• Density = massvolume

= grammL

Different ways of calculating volume

• I. For solids with regular shapes:

• A. For a cube: Vcube = s3

• B. For a rectangular solid, V = L x W x H

• C. For a cylinder: V= r2h

• D. For a sphere: V = 4/3 r3

Different ways of calculating volume

• II. For an Irregular Solid

• Water displacement

Different ways of calculating volume

• III. For a liquid

• Use of graduated cylinder, beaker, pipet or buret.

Problem

• A cube of osmium metal 1.500 cm on a side has a mass of 76.31 grams at 25 oC. What is its density in g/cm3 at this temperature?

Problem

• The density of titanium metal is 4.51 g/cm3 at 25 oC. What mass of titanium displaces 65.8 mL of water at 25 oC?

Problem

• The density of benzene at 15 oC is 0.8787 g/mL. Calculate the mass of 0.1500 L of benzene at this temperature.