Matter and Scientific Measurement Page 16

Matter and Scientific Measurement Page | 16

Unit 1: Matter and Scientific Measurement

(Link to Prentice Hall Text: Chapters 2, 3 and 4)

Name:______

Due Date / Assignments / Page Number, Problem Numbers
Assignment 1 / Memorize the conversion between symbols and names for elements 1-38, 47, 53-56, 78-80, 82, 85-88.
Assignment 2: Identifying Mixtures; Physical/Chemical Changes / 47: 30, 32, 34, 35
Assignment 3: Separating Mixtures / 47: 38, 41
Assignment 4: Calculations with Significant Digits / 78: 41, 44
Assignment 5: Calculations with Density / 79: 61

A. Matter, Change and The Governing Law of Chemistry

Chemistry is the study of how “stuff” (matter) changes. Chemists are particularly interested in how matter changes at the particle level, on a scale that is much too small to see. Because the scale on which chemists seek to understand change cannot be seen, chemists often perform experiments on an observable scale and then use their evidence to make inferences about change at the particle level.

Chemists point to two types of change: physical and chemical.

Physical change:

Examples:

Chemical change:

Examples:

Evidence of a chemical change:

Exercise

Classify the following as being either a chemical or a physical change:

1. Sodium chloride dissolves in water. ______

2. Hydrochloric acid reacts with sodium hydroxide to produce a salt, water and heat. ______

3. A pellet of sodium is sliced in half. ______

4. Water is heated and changed to steam. ______

5. Food is digested. ______

6. Starch molecules are formed from smaller glucose molecules. ______

7. Ice melts. ______

8. Plant leaves lose water through evaporation. ______

9. A red blood cell placed in distilled water swells and bursts. ______

10. The roots of a plant absorb water. ______

11. Iron rusts. ______

12. A person cools by sweating. ______

13. A match burns. ______

The Law of Conservation of Matter states: ______

______

The Law of Conservation of Energy states: ______

______

Practice with the Law of Conservation of Matter

Put a (P) in the blank if the statement is probable. Put a (I) in the blank if the statement is improbable.

1. If a fire is hot enough, it could convert garbage into pure energy. ______

2. When wood burns in a fireplace, the ashes have the same mass as the wood before it was burned. ______

3. When you exercise and burn calories, fat is converted to energy and the mass disappears. ______

4. When a tree grows, it converts the sunlight into the mass that becomes the bark, the trunk, the branches and the leaves. ______

5. When you were a baby, developing in utero, you grew because energy in your mother’s body was converted into flesh. ______

B. Four Properties of Matter

Physical:______

Chemical:______

Extensive: ______

Intensive: ______

C. Types of Matter

Atom

Molecule

Pure Element

Pure Compound

Mixtures

Heterogeneous Mixtures

Homogeneous Mixtures

Exercise

Classify each of the following substances as a pure substance (P) or as a mixture (M). If the substance is pure, classify it as an element (E) or a compound (C). If the substance is a mixture, classify it as a heterogeneous (HT) or homogeneous (HO) mixture.

1. Sodium (Na)______

2. Water (H2O) ______

3. Soil ______

4. Coffee______

5. Oxygen (O2)______

6. Alcohol (CH3CH2OH)______

7. Carbon Dioxide (CO2)______

8. Cake Batter______

9. Air (Mixture of N2, O2, CO2 and Ar)______

10. Soap______

11. Iron (Fe) ______

12. Salt Water______

13. Ice Cream______

14. Nitrogen (N2) ______

15. Eggs______

16. Blood______

17. Table Salt (NaCl) ______

18. Methane (CH4) ______

19. Gold (Au) ______

20. Palladium (Pd) ______

D. Separation of Mixtures

(1) Filtration

a. Solid in a Liquid

b. Solid in a Gas

c. Liquid in a Liquid

(2) Distillation

a. Solid in a Liquid

b. Liquid in a Liquid

(3) Chromatography

a. Separating Liquids

E. Units

Quantity / Symbol / Unit Name / Unit Symbol
Length / l / meter / m
Mass / m / gram / g
Time / t / second / s
Temperature / T / Kelvin / K
Amount of substance / n / mole / mol
electric current / I / Ampere / A
luminous intensity / Iv / Candela / cd
Quantity / Symbol / Unit Name / Unit / Derivation
Area / A / square meters / m2 / A = l X w
Volume / V / cubic meters / m3 / V = l X w X h
Density / d / grams per mL / g/mL / d = m/V
Molar mass / MM / grams per mole / g/mol / MM = m/n
Concentration / c / moles per liter / M / M = mol/V
Molar volume / Vm / liters per mole / L/mol / Vm = V/mol

F. Density

Density -

You must show ALL work and circle your final answer.

1. 100 grams of a liquid completely fill a 200 mL bottle. What is the density of the liquid?

2. A solution has a density of 1.50 g/mL. How many grams are needed to obtain 10.0 mL of solution?

3. If a block of copper measures 2.00 cm x 4.00 cm x 5.00 cm and weighs 356 grams, what is its density in g/mL?

4. The density of mercury is 13.6 g/mL.

a. what is the mass of 8.20 mL of mercury?

b. what volume would 120 grams of mercury occupy?

5. A student pipets 5.00 mL of ethanol into a flask weighing 15.25 grams. She finds that the mass of the flask plus ethanol = 19.17 grams. Calculate the density of ethyl alcohol.

6. A chemist needs 2.00 g of a liquid compound, which has a density of 0.718 g/mL. If the compound costs $5.67 per mL, how much will a 2.0 gram sample cost?

G. Accuracy and Precision

Definitions:

Accuracy –______

Precision – ______

Precision versus Accuracy:

Look at each target and decide whether the “hits” are accurate, precise, both accurate and precise, or neither accurate nor precise: (Note: An accurate “hit” is a bulls eye!)

Expressing Errors in Measurement:

Scientists often express their uncertainty and error in measurement by giving a percent error. The percent error in chemistry is defined as:

Answer the following three questions. Pay attention to significant figures, and show your work!

1. While doing a lab, a student found the density of a piece of pure aluminum to be 2.85 g/cm3. The accepted value for the density of aluminum is 2.70 g/cm3. What was the student's percent error?

2. A student measured the specific heat of water to be 4.08 J/g • Co. The literature value of the specific heat of water is 4.18 J/g • Co. What was the student’s percent error?

H. Significant Digits in Measurements

Significant figures – These are all the digits you know for sure, plus one place that is an estimate.

Uncertainty – Limit of precision of the reading (based on your ability to estimate the final digit). See examples below.

Rules for zeros: All zeros count except placeholder zeros – these are the ones that disappear when you write the number in scientific notation. Examples:

93,000,000 = 9.3 x 107 2 sf’s

0.000372 = 3.72 x 10-4 3 sf’s

0.0200 = 2.00 x 10-2 3 sf’s

Exercise

For each of the following, write the scale reading, then the number of significant figures in the reading.

Reading SF’s

1. 1.

2. 2.

3. 3.

4. 4.

5. 5.

6. 6.

7. 7.

8. 8. -

I. The Atlantic Pacific Rule For Determining Significant Digits

(1) / Pacific – "P" is for decimal point is present. If a decimal point is present, count significant digits starting with the first non-zero digit on the left.

Examples: / (a) 0.004703 has 4 significant digits.
(b) 18.00 also has 4 significant digits.

(2) / Atlantic – "A" is for decimal point is absent. If there is no decimal point, start counting significant digits with the first non-zero digit on the right.

Examples: / (a) 140,000 has 2 significant digits.
(b) 20060 has 4 significant digits.

Imagine a map of the United States. If the decimal is absent count from the Atlantic side. If the decimal point is present, count from the Pacific side. In both cases, start counting with the first non-zero digit.

How many significant digits are in each of the following numbers?
a) 45.67
b) 4095
c) 30,000,000
d) 45.00043
e) 20
f) 1.0
g) 4
h) 34.8700
i) 0.0034500600
Calculations shouldn't have more precision than the least precise measurement. This leads to 2 Rules for Calculations:
(A) / For addition and subtraction: The answer should not have more places past the decimal than the number with the least places past the decimal.

Example: / 1.2 + 12.348 = 13.5 / Not 13.548

(B) / For multiplication and division: The answer should not have more significant figures than the number with the least amount of significant figures.

Example: / 502 x 3.6 = 1800 / Not 1807.2

These last 2 rules can be called the Many-Places rule. For multiplication/division, how many significant figures is important. For plus/minus, number of places is important.

I. How many significant digits are in the following numbers?

1. 3.4069

2. 0.56

3. 0.00890

4. 25,000

5. 14.987

II. Write each of the above numbers with one significant digit.

1. ______

2. ______

3. ______

4. ______

5. ______

III. Multiply or divide, according to the problem given below. Make sure that your answer contains the correct amount of significant digits.

1. 5.67 cm × 6 cm =

2. 0.004090 mm × 12.4 mm =

3. 34 m × 1 m =

4. 30,000 m ÷ 9.008 =

5. 5.67miles ÷ 8.07 =

IV. Add or subtract, according to the problem given below. Make sure that your answer contains the correct amount of significant digits.

1. 4.56 cm + 6.0 cm =

2. 0.00089 m + 3.4 m =

3. 4.5 m – 0.897 m =

4. 25,000 m/s – 349.00 m/s =

5. Tricky! 67.8 °C -0 °C =

H. Factor-Label Method of Dimensional Analysis

Name / Symbol / Size
nano / n / 1 × 10-9 _ = 1 n_
micro / μ / 1 × 10-6 _ = 1 μ _
milli / m / 1 × 10-3 _ = 1 m_
centi / c / 1 × 10-2_ = 1 c_
base
kilo / k / 1 × 103 _ = 1 k_
Mega / M / 1 × 106_ = 1 M_

Start with a number fact, such as 4.1 cm or 0.075 mL. Examine the units of the desired answer. Multiply your fact with the factor . The starting units cancel out and you end up with the desired units. Some conversions require more than one factor. For example, we do not typically convert directly from kg to μg. So, the best approach is to convert from kg to g (the base unit) then from g to μg.

Remember, even though we write factors with x signs, we multiply by the numerators and divide by the denominators.

Exercise

Use the factor-label method to make the following conversions. Remember to use the appropriate number of significant figures in your answer.

1. 74 cm x = meters

2. 8.32 x 10-2 kg x = grams

3. 55.5 mL x = cm3

4. 0.00527 cal x = kilocalories

5. 9.52 x 10-4 m x = micrometers

6. 41.0 mL x = liters

7. 6.0 x 10-1 g x = mg

8. 8.34 x 10-9 cg x = g

9. 5.0 x 103 mm x = m

10. 1 day x x x = seconds

11. 5 x 104 mm x x = km

12. 9.1 x 10-13 kg x x = ng

13. 1 year x x = hours (approximately)

14. 4.22 cL x x = mL

15. 1 mile x x = inches

Exercise

Use the factor-label method to make the following conversions. Remember to use the appropriate number of significant figures in your answer.

1. How many nickels could you trade for 250 yen? $1 = 150 yen.

2. Your school club sold 600 tickets to a chili supper. The chili recipe for 10 persons requires 2 teaspoons of chili powder? How many teaspoons of chili powder will you need altogether?

3. How many cups of chili powder will you need? Three teaspoons (tsp) equal one tablespoon (TBS) and 16 tablespoons equal 1 cup.

4. How many seconds in a year? (Assume 365 day in a year.)

5. Chloroform is a liquid once used for anesthetic purposes. What is the volume of 5.0 g of chloroform? The density of chloroform is 1.49 g/mL.

6. How many inches long is a football field? (A football field is 100 yards.)

7. How many m3 is 4.6 cm3? Express your answer in scientific notation.

8. How many mg is 59.0 kg? Express your answer in scientific notation.