Chapter 2: Measurement and Calculations… Section 2-1: Scientific Method (pg29-31) will not be...

Chapter 2: Measurement and Calculations…

Section 2-1: Scientific Method (pg29-31) will not be explicitly covered but used

throughout this entire class…

1

Section 2-2: Units of Measurement & Metric review

and summary…Pgs 33-39

2

The Fundamental Metric System (SI) Units…some of them

Physical Quantity Name Abbreviation

Mass kilogram kg

Length meter m

Time second s

Temperature Kelvin K

Electric Current Ampere A

Amount of Substance mole mol

Luminous Intensity candela cd

SI Units

4

SI (a.k.a. Metric system) PrefixesCommon to Chemistry

Prefix Unit Abbr. ExponentKilo k 103

Deci d 10-1

Centi c 10-2

Milli m 10-3

Micro 10-6

We will use these units primarily in this class, but these are only a few of the units within the SI (Metric System)

5

KNOW THIS CHARTKNOW THIS CHART

(BIGGER) (smaller)mega-

(M)kilo-

(k)BASE UNIT

metergramLiter

deci-(d)

centi-(c)

milli-(m)

micro-(µ)

Even though there aren’t any prefixes that we need to use there, it’s important to remember

where the BLANK spaces are on this table.

Even though there aren’t any prefixes that we need to use there, it’s important to remember

where the BLANK spaces are on this table.

Group work/discussion…

• Get a white board and marker for the group…

7

Density (review hopefully!...) Check out this picture on the right... •Which layer has the highest density?•Which layer has the lowest density?•Imagine that the liquids have the following densities:

– 10g/cm3.– 3g/cm3.– 6g/cm3.– 5g/cm3.

•Which number would go with which layer?

Density…• Density is defined as the ratio of

mass (matter) to volume.• Density (D) is usually expressed

D = m/v• Units are most commonly g/ml or

g/cm3 , but can be in ANY units of mass to volume!

• Density is a characteristic physical property that does not change with size of sample (BUT some conditions can affect density value…like what?)

9

Conversion Factors (pg 40-42)

• We’ll cover this soon…not yet.

10

(p. 44 - 57)

Section 2-3: Using Scientific Measurements

11

A. Accuracy vs. Precision• Accuracy - how close a measurement is to the

accepted value

• Precision - how close a series of measurements are to each other

ACCURATE = CORRECT

PRECISE = CONSISTENT12

B. Percent Error

• Indicates accuracy of a measurement

100literature

literaturealexperimenterror %

your value

accepted value13

B. Percent Error• A student determines the density of a substance to

be 1.40 g/mL. Find the % error if the accepted value of the density is 1.36 g/mL.

100g/mL 1.36

g/mL 1.36g/mL 1.40error %

% error = 2.9 %

14

C. Significant Figures

• Indicate precision of a measurement.

• Recording measurements in Sig Figs…– Sig figs in a measurement include the known digits

plus a final estimated digit

2.35 cm

15

C. Significant Figures

• Counting Sig Figs (textbook Table 2-5, p.47)

– Zeros that are in between any digits 1-9 are ALWAYS counted as significant

• Ex: 105 = 3 sig figs

1208.5 = 5 sig figs

– Count all numbers (1 though 9) EXCEPT…

• Leading zeros -- 0.0025

• Trailing zeros without a decimal point -- 2,500

16

4. 0.080

3. 5,280

2. 402

1. 23.50

C. Significant Figures

Counting Sig Fig Examples

1. 23.50

2. 402

3. 5,280

4. 0.080

4 sig figs

3 sig figs

3 sig figs

2 sig figs17

C. Significant Figures

• Calculating with Sig Figs– Multiply/Divide - The # with the fewest sig figs

determines the # of sig figs in the answer.

(13.91g/cm3)(23.3cm3) = 324.103g

324 g

4 SF 3 SF3 SF

18

C. Significant Figures

• Calculating with Sig Figs (con’t)– Add/Subtract - The # with the lowest decimal value

determines the place of the last sig fig in the answer.

3.75 mL

+ 4.1 mL

7.85 mL

224 g

+ 130 g

354 g 7.9 mL 350 g

3.75 mL

+ 4.1 mL

7.85 mL

224 g

+ 130 g

354 g 19

C. Significant Figures

• Calculating with Sig Figs (con’t)– Exact Numbers do not limit the # of sig figs in the answer.

• Counting numbers: 12 students• Exact conversions: 1 m = 100 cm• “1” in any conversion: 1 in = 2.54 cm

20

C. Significant Figures

5. (15.30 g) ÷ (6.4 mL)

Practice Problems

= 2.390625 g/mL

18.1 g

6. 18.9 g

- 0.84 g18.06 g

4 SF 2 SF

2.4 g/mL2 SF

21

D. Scientific Notation (and standard notation)

• PROPER Sci. notation uses format: M x 10n

– M is larger than 1 and smaller than 10 (1.000 to 9.999)

• To converting into Sci. Notation:

– Move decimal until there’s 1 digit to left of decimal (COUNTING places moved tells you about the exponent)

– Large numbers (>1) positive exponentsSmall numbers (<1) negative exponent

– Only include sig figs. (significant figures)

65,000 kg 6.5 × 104 kg

22

D. Scientific Notation

• 2,400,000 g

• 0.00256 kg

• 7 10-5 km

• 6.2 104 mm

Practice Problems

2.4 106 g

2.56 10-3 kg

0.00007 km

62,000 mm23

D. Scientific Notation

• Calculating with Sci. Notation

(5.44 × 107 g) ÷ (8.1 × 104 mol) =

5.44EXPEXP

EEEE÷÷

EXPEXP

EEEE ENTERENTER

EXEEXE7 8.1 4

= 671.6049383 = 670 g/mol = 6.7 × 102 g/mol

Type on your calculator:

24

E. Proportions

• Direct Proportion

Inverse Proportion

xy

xy

1

y

x

y

x 25