Chapter 1 Chemistry: Matter and Measurement

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Honors Chemistry Unit 1 Notes: Matter and Measurements

Objectives:

1. Students will become familiar with the meanings of various introductory concepts such as the meanings of the words: chemistry, substance, element, compound, atom, molecule, heterogeneous and homogeneous mixtures, and matter.

2. Students will know the difference between intensive and extensive physical properties and chemical properties, between physical and chemical changes.

3. Students will review the basic steps and vocabulary of the scientific method.

4. Students will be familiar with the SI units for mass, length, temperature, time and volume and the various metric prefixes and conversion problems.

5. Students will be able to identify the characteristics of precision and accuracy in a set of data.

6. Students will demonstrate proficiency in using scientific notation, and the rules of significant figures in computations involving data.

7. Students will use the unit-conversion method of solving problems.

8. Students will become familiar with the use of density in problem solving.

Review: Scientific Method

1. Hypothesis:

2. Experiment:

3. Data:

a. Data is always ________________.

4. Scientific Law

a. Scientific laws explain WHAT is observed.

1. Example of a scientific law:

5. Scientific Theory

b. Scientific theories explain WHY something is observed.

1. Example of a scientific theory:

6. Scientific Method—Review—

a.

b.

c.

d.

e.

f.

1.1 Introduction and Types of Matter

A. General terms we will use throughout this course (some are review)

1. Chemistry—definition

Chemistry deals with the properties and reactions of substances.

2. Matter is anything that has ___________and occupies ________________.

It exists in three phases: _____________, ____________and ____________.

Fill in the table below about these three phases:

|Phase |How would you describe shape? |How would you describe volume? |

| | | |

| | | |

| | | |

3. Matter can be classified into ________categories—

• ____________________________, each of which has a __________ composition and a ______________set of properties.

• ____________________, composed of __________or more substances.

4. Pure substances are either _____________or _______________, whereas mixtures can be either ___________________or _______________________.

B. Elements

1. An element is _________________________________________________.

a. How many elements are there?

b. How is an element identified?

C. Compounds

1. A compound is a _______________________________________________.

a. What elements does water contain?

b. What elements are contained in methane, acetylene and naphthalene?

c. Compounds have fixed ____________________. That is, a given compound contains the same _________________ in the same _______________________.

d. Are the properties of compounds same/different from the elements they contain?

e. Give an example to support your answer to the question in d):

f. Name two methods that can be used to resolve compounds into their elements:

D. Mixtures

1. Mixtures—definition:

2. There are two types of mixtures:

a. Homogeneous—definition:

Another name for a homogeneous mixture is a __________________.

b. Heterogeneous—definition:

3. Name 3 different laboratory methods that can be used to separate the components of a mixture:

END OF NOTES REQUIRED IN HOMEWORK #1

1.3 Scientific Measurements

A. Measurement Systems

Chemistry is a ___________________________science.

This means that experiments and calculations almost always involve _____________________ _________________.

Scientific measurements are expressed in the __________________________.

This is a _______________based system in which all of the units of a particular

quantity are related to each other by factors of _______________.

B. Prefixes (see handout)

You will need to memorize all of the prefixes (factors, names and abbreviations from

109 (Giga-) to 10-9 (nano)!

One example of a memory device:

C. Instruments and Units

1. LENGTH: The standard unit of length in the metric system is the _______________, which is a little larger than a _______________.

USING THE PREFIXES WITH LENGTH:

2. VOLUME: The common units of volume in chemistry are the _______________________ and the ______________________________________.

The common instruments for measuring volume in chemistry are: _______________________ and the _____________.

Note that 1 cm3 = 1 mL (We will use this exact conversion factor throughout the year, so you will need to memorize it!)

3. MASS: The common unit of mass in chemistry is the ________________ (used in lab).

****______________ is a measure of the amount of matter in an object; _______________ is a measure of the gravitational force acting on the object. Chemists often use these terms interchangeably.****

4. TEMPERATURE: Temperature is the factor that determines the __________________________.

a. Comparison of Fahrenheit, Celsius, and Kelvin scales

ENGLISH UNITS (INCHES, °F, ETC.) ARE NEVER USED TO TAKE MEASUREMENTS IN LAB!

b. Conversion formulas for temperature:

Example 1L.1 A baby has a temperature of 39.8oC. Express this temperature in oF and K.

5. ESTABLISHMENT OF THE INTERNATIONAL SYSTEM OF UNITS (SI) –

See Appendix 1 on p. 635 in your book)

SI UNITS AS ESTABLISHED BY THE SI:

LENGTH—

VOLUME—

MASS—

TEMPERATURE--

D. Uncertainties in Measurements

1. Precision vs. Accuracy

Definitions:

Precision – how close answers are to _________________________ (_________________________)

Accuracy – how close answer is to ___________________________________

(_________________________________________________________)

2. Percent Error—used to calculate accuracy of results

Equation

Ex1.9 A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is

2.70 g/mL. What is the percent error for the student’s results?

E. Scientific Notation - see handout

Significant Figures (sig figs)

Why are they important?

Numbers in math: vs. Numbers in chemistry:

A. Significant figures – all the digits in a measurement that are known with certainty plus a ________________________________________________________________________

B. Rules for Counting Sig Figs

Rule #1:

Rule #2:

Rule #3:

Rule #4:

You must memorize the rules and learn to use them or lose points throughout the year!

Example 1L.2 State the number of significant figures in the following set of measurements:

o 1.2304 mm

o 1.23400 cm

o 1.200 x 105 mL

o 0.0230 m

o 0.02 cm

o 8 ounces = 1 cup

o 30 cars in the parking lot

C. General Rounding Rule: When a number is rounded off, the last digit to be retained is increased by one only if the following digit is 5 or greater.

Example 1L.3 Solve the following problems and state the answers with the proper number of significant figures.

a. Calculate the area of an object with a length of 1.345 m and a width of 0.057 m.

b. Calculate the volume of the same object with a thickness of 3.40 x 10-2 m.

c. Calculate the density of a substance with a mass of 12.03 g and a volume of 7.0 mL.

D. Sig Fig Process for Addition/Subtraction

Step #1:

Step #2:

Examples:

E. Sig Fig Process for Multiplication/Division

Step #1:

Step #2:

Examples:

Example 1L.4 Express the answers below with the correct number of significant figures:

a. 129.0 g + 53.21 g + 1.4365 g =

b. 10.00 m - 0.0448 m =

c. 23.456 × 4.20 × 0.010 =

d. 17 ÷ 22.73 =

****When you are doing several calculations, carry out all the calculations to at LEAST one more sig fig than you need (I carry all digits in my calculator memory) and only round off the FINAL result.

Conversion of Units

A. Use of conversion factors (a.k.a. __________________________ or ________________ method)

Definition:

Set-up of Problems & Examples:

B. Example 1L.5 Calculate the following single step conversions: a. How many joules are equivalent to 25.5 calories if 1 cal = 4.184 joules? b. How many liters gasoline can be contained in a 22.0 gallon gas tank if 3.785 L = 1 gal?

C. Example 1L.6 The following multiple step conversion can be solved, knowing that

1 in = 2.54 cm. Convert the length of 5.50 ft to millimeters.

D. Multiple unit conversions. Example 1L.7 The average velocity of hydrogen molecules at 0oC is 1.69 x 105 cm/s. Convert this to miles per hour.

Example 1L.8 A piece of iron with a volume of 2.56 gal weighs 168.04 lbs. Convert this density to scruples per drachm with the following conversion factors: 1.00 L = 0.264 gal, 1.000 kg = 2.205 lb, 1.000 scruple = 1.296 g, 1.000 mL = 0.2816 drachm.

E. Area and Volume Conversions

Example 1L.14 Express the area of a 27.0 sq yd carpet in square meters.

Conversion factors needed:

Example 1L.15 Convert 17.5 quarts to cubic meters. (1 L = 1.057 qt, 1 ft3 = 28.32 L)

1.3 Properties of Substances

1. Every pure substance has its own unique set of _______________that serve to ____________________________________________________________.

2. Properties used to identify a substance must be __________________; that is, they must be independent of __________________.

_________________properties depend on the amount.

Classify the following as either intensive (I) or extensive (E) properties:

a. density

b. mass

c. melting point

d. volume

3. Chemical property:

o Example of a chemical property:

4. Physical property:

o Example of a physical property:

5. Physical vs. chemical change:

6. Another name for a chemical change is a ________________________________

7. Classify the following as either physical (P) or chemical (C) changes:

a. ice melting

b. gasoline burning

c. food spoiling

d. log of wood sawed in half

Density

1. Review

a. Definition:

b. Formula:

Sample Problem A piece of copper has a mass of 57.54 g. It is 9.36 cm long, 7.23 cm wide, and 0.95 mm thick. Calculate density (g/cm3).

Strategy

1. Get dimensions in common units.

2. Calculate volume in cubic centimeters.

3. Calculate the density.

Sample Problem: Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in grams? In pounds?

Strategy

1. Use density as a conversion factor to calculate mass (g) from volume.

2. Convert mass (g) to mass (lb)

Need to know conversion factor; 454 g = 1 lb

Ex1L.9 What is the density of Hg if 164.56 g occupy a volume of 12.1 cm3?

Ex1L.10 What is the mass of 2.15 cm3 of Hg?

Ex1l.11 What is the volume of 94.2 g of Hg?

Example 1L.12 Given the following densities: chloroform 1.48 g/cm3 and mercury 13.6 g/cm3 and copper 8.94 g/cm3. Calculate if a 50.0 mL container will be large enough to hold a mixture of 50.0 g of mercury, 50.0 g of chloroform and a 10.0 g chunk of copper.

Example 1L.13 How many kilograms of methanol (d = 0.791 g/mL) does it take to fill the 15.5-gal fuel tank of an automobile modified to run on methanol?

HOMEWORK ASSIGNMENTS ON NEXT PAGE ( ( (

HOMEWORK ASSIGNMENTS

Note: Homework exercises for this chapter are found either in the photocopied pages starting on p. 19 of this packet or on worksheets handed out in class. Additional worksheets could be given out in class to supplement these assignments!

Homework #1 – Read pages Section 1.1 in the photocopied pages from Masterton and Hurley starting on the next page of this packet (p. 19-23 of this packet). Fill in notes packet pgs. 2-4 as you read (beginning with the section entitled “Introduction and Types of Matter”). This material will NOT be covered in class, BUT you will be responsible for its content on tests/quizzes.

Homework #2 – Answer Questions 1, 3, 5, 7, 9 from p. 24 of this packet.

Homework #3 – Exponents and Scientific Notation Worksheet—Read entire worksheet and fill in answers to all numbered exercises.

Homework #4 – Answer Questions 11, 13, 15 from p. 24 of this packet.

Homework #5 – Complete BOTH Sig Figs Worksheet (handed out in class) and answer Questions 17, 19, 21, 27 from pages 24 & 25 of this packet.

Homework #6 – Answer Questions 33, 35, 39, 41,43 from p. 25 of this packet.

Homework #7 – Answer Questions 45, 47, 49, 51 from p. 26 of this packet.

Homework #8 – Answer Questions 59, 61, 65, 67, 69 from p. 26 of this packet.

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