Chapter 1 Chemistry: Matter and Measurement

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

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

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

6. Students will use the dimensional analysis method of solving problems.

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

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 reactionsof substances.

2. Matter is anything that has ______and occupies ______.

It exists in three phases: ______, ______and ______.

Fill in the table below about these three phases:

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 compoundcontains 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 particularquantity are related to each other by factors of ______

B. SI System: Units & Measuring

SI UNITS (Based on the metric system)

LENGTH –

MASS –

VOLUME –

TIME –

TEMPERATURE –

Any measurement should include two parts: a ______and a ______

• You should always report measurements with all readable numbers plus ______

• This ensures precision & accuracy!

Report all measurements with all readable numbers + only one estimated digit!

D. Uncertainties in Measurements

1. Precision vs. Accuracy

Precision – how close answers are to ______(______)

Accuracy – how close answer is to ______

(______)

2. Percent Error – used to calculate accuracy of results

Equation:

Example #1_A student reports the density of a pure substance to be 2.83 g/mL. The accepted value is

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

E. Scientific Notation - see handout

Significant Figures (sig figs)

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

Why are they important?

Numbers in math: vs.Numbers in chemistry:

1. For the purposes of significant figures, there are two major categories:
2. Nonzero digits:
3. Zero digits:

1. Rules for Counting Sig Figs
2. All ______are significant!

Ex.) 3269  ______significant figures

257  ______significant figures

1.234567  ______significant figures

1. Zeroes take three forms:
2. ______zeroes
3. ______zeroes
4. ______zeroes

1. Leading Zeroes = zeroes that come ______the nonzero digits in a number.
2. Leading zeroes are placeholders only and are ______considered significant.
3. Ex.)

1. Trapped Zeroes = zeroes ______two nonzero digits.
2. Trapped zeroes are ______.
3. Ex.)

1. Trailing Zeroes = zeroes that ______nonzero digits.
2. Trailing zeroes are only significant if there is a ______in the number.
3. Ex.)

1. Exact Numbers: Any number which represents a numerical count or is an exact definition has an infinite number of sig figs and is NOT counted in the calculations.
2. Ex.)

1. Practice: How many significant digits are in each of the following?

1.034 s3000 m

0.0067 g72 people

12 apples

1. Rules for Using Sig Figs with Addition/Subtraction
2. Step #1: Determine the number of ______in each number to be added or subtracted.
3. Step #2: Calculate the answer and then round the final number to the least number of ______from Step #1.
4. Examples:

1. Rules for Using Sig Figs with Multiplication/Division
2. Step #1: Determine the number of ______in each number to be multiplied or divided.
3. Step #2: Calculate the answer and then round the final number to the least number of ______from Step #1.
4. Examples:

Example #3_Express the answers below with the correct number of significant figures:

a. 129.0 g + 53.21 g + 1.4365 g =

1. 10.00 m - 0.0448 m =

1. 23.456 × 4.20 × 0.010 =

1. 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.k.a Dimensional Analysis)

A. Use of dimensional analysis(a.k.a. factor-label method, unit conversion, conversion factors, etc.)

Definition:

Set-up of Problems & Examples:

B. General Dimensional Analysis

Example #4_Calculate the following single step conversions:

a)How manyjoules are equivalent to 25.5 calories if 1 cal = 4.184 joules?

b)How many milliliters of gasoline can be contained in a 22.0 gallon gas tank if 3.785 L = 1 gal and there are 1000 mL in 1 liter?

C. Metric Conversions with Dimensional Analysis

The given powers of ten all refer back to the base unit. We are going to set up our conversions so that the 1 always goes with the larger unit and the power of ten is always positive!

These prefixes are based on powers of 10. What does this mean?

• For each prefix, every “step” is either ______or ______
• Examples:
• 1 dL = ______mL
• 1 hg = ______dg
• 1 Gm = ______μm

Example #5_

a)How many millimeters are in 1.25 kilometers?

b)Lake Erie contains 480,000 GL of water. How many microliters is this?

c)A person weighs 5.443 x 1010 μg. How many pounds is this equal to? (1 kg = 2.205 lbs.)

Example #6_Convert the length of 5.50 ft to millimeters (1 in = 2.54 cm)

D. Dimensional Analysis with Derived (Compound Units)

______are units that are composed of multiple base units.

• A.k.a. compound units

Examples:

Example #7_The average velocity of hydrogen molecules at 0oC is 1.69 x 105 cm/s. Convert this to miles per hour.

Example #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 #9_Express the area of a 27.0 sq yd carpet in square meters.

(1 yard = 3 feet, 1 foot = 0.3048 meters)

Example #10_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:

• Example of a chemical property:

4. Physical property:

• Example of a physical property:

5. Physical vs. chemical change:

6. Another name for a chemical change is a ______

Example #11_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

a. Definition:

b. Formula:

Example #12_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).

Example #13_Mercury (Hg) has a density of 13.6 g/cm3. What is the mass of 95 mL of Hg in pounds? (Need to know conversion factor; 454 g = 1 lb)

Example #14_What is the density of Hg if 164.56g occupy a volume of 12.1cm3?

Example #15_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 #16_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?

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