b) compound – a substance of definite composition of elements

b) Compound – a substance of definite composition of elements. It can be decomposed into two or more simpler substances by simple chemical changes but not by physical means.

Example: SMORES

1 smore contains 2 graham crackers, 1 marshmallow, and 3 chocolates

6 smores contain: graham crackersmarshmallowschocolates

Ratio of crackers to marshmallows to chocolates = 2:1:3

1

126

18

1 smore contains 2 graham crackers, 1 marshmallow, and 3 chocolates

Ratio of crackers to marshmallows to chocolates = 2:1:3

2

49 graham crackers23 marshmallows69 chocolates

23 smores!

Chapter 2 Measurements and Units

1. Metric System

mass – gram, g (1 g = 0.0022 pounds)

length – meter, m (1 m = 39.37 inches)

volume – liter, L (1 L = 1.06 Quarts)

temperature – Celcius or centigrade, oC Freezing point of water = 0oC Boiling point of water = 100oC Energy – calorie, cal

1 cal = energy required to change 1 g of water 1oC Time – second, s

2. Prefix

pico p 10-12 = 0.000000000001(one-trillionth)

centimeter (cm) 1cm = 0.01 m = 1/100 m

millimeter (mm) 1 mm = 0.001 m = 1/1000 m

milliliter (mL) = 0.001 L = 1/1000 L

kilometer (km) 1 km = 1000 m

kilogram (kg) 1 kg = 1000 g

kilocalorie (kcal) 1 kcal = 1000 cal

megabite (MB) 1MB = 1,000,000 bites

gigabite (GB) 1 GB = 1,000,000,000 bites

Copyright © Houghton Mifflin Company. All rights reserved. 2–4

Figure 2.3 A cube 10 cm on a side has a volume of 1000 cm3, which is equal to 1 L.

1 c.c.

3. Accuracy and Precision Accuracy - the correctness of a measurement

Precision – reproducibility of measurements.

4. Significant figures

Copyright © Houghton Mifflin Company. All rights reserved. 2–6

Figure 2.5 The scale of a measuring device determines the magnitude of the uncertainty for the recorded measurement.

4. Significant figures

1.345 significant figures

0.2300 significant figures

0.00230

23.400500

5. Exponential notation (scientific notation)

23000000

2 significant figures

2.300 x 107

? significant figures

0.000000230 = 2.30 x 10-7

?

2.300 x 10-7 ?

= 2.3 x 107

Perfect units, exact numbers and inexact numbers

An exact number is a number whose value has no uncertaintyassociated with it – a number that arises when you count items or when you define a unit.

1 ft = 12 in1 hr = 60 min1 cm = 1/10 m½, ¼, etc.

Perfect units Exact numbers

1 kg = 2.205 lbInexact number

Infinite significant figures.

Infinite significant figures.

The value of an inexact number has a degree of uncertainty.

All measurements are inexact

Copyright © Houghton Mifflin Company. All rights reserved. 2–6

Figure 2.5 The scale of a measuring device determines the magnitude of the uncertainty for the recorded measurement.

3.8

3.75

6. Significant figure in arithmetic a) addition and subtraction

2.34 + 30.6081

+ 30.6081 2.34

32.9481 32.95

b) Multiplication and division

1.32 x 4.011 = 5.29452= 5.29

3 significant figures

6.3 x 0.000834 = 0.0053

2 significant figures

24

38

1065.310102.41021.6100.2

7. Conversion of units

It is experimentally determined that 1 inch equals 2.54 centimeters, or 1 centimeter equals 0.394 inch.

incm

cmin

154.21

54.21

Example 1Convert 2.00 inches to centimeters.

Example 2Convert 9.05 cm to inches. Factor-Label Method

Example 36.82 cm = ? feet

Example 4How many milligrams are there in 4.00 pounds?

8. Density Density = mass/volume (g/mL)

Figure 2.7 (a) The penny is less dense than the mercury it floats on. (b) Liquids that do not dissolve in one

another and that have different densities float on one another, forming layers.

a)

b)

Example 12.0 g of metal occupies 0.40 mL. What is the density?

Example 2Density = 2.0 pounds/quart = ? g/mL

Specific Gravity – ratio of the density of a substance to that of water.

Density of lead = 11.3 g/mLspecific gravity of lead = 11.3

Example:

density of water = 1.00 g /mL

no unit

9. Temperature Figure 1.4

Change of 180oF = change of 100oC

oF oC

0o

100o

32o

212o

180o 100o

Change of 180oF = change of 100oC

Change of 1oF = change of 100180

oC = change of 59

oC

Change of 1oC = = change of 95

oF

Example: 50oF = ?oC oF oC

0o32o

33o

34o1o

oF oC

0o32o

212o 100o

Absolute Temperature – Kelvin, K

K = 273 + oC

examples

FC

CF

oo

oo

3259

95)32(

10. Heat energy and specific heat

Unit of heat energy – calorie (cal)1 cal = heat needed to raise 1 g of water 1oC

1 kcal = 1000 cal1 cal = 4.184 Joules (J)

SI system

Specific heat (SH) of a substanceSH = amount of heat needed to raise 1 g of that substance 1oC

SH = heat absorbed

Mass x change of T=

calm x DT

cal = SH x m x DT

cal = SH x m x DT

SH of water = 1.0 ( cal g . oC)

Example 1 : How much heat is needed to raise 30 g water from 20oC to 30oC?

Example 2

122 cal of heat is added to 20 g of methanol at 15oC. What is the final temperature?

SH of methanol = 0.61 cal/g . oC

Example 3What quantity of heat is required to raise the temperature of 50 mL of ethanol from 22.0oC to 25.0oC? The density of ethanol at this temperature is 0.80g/mL. The specific heat of ethanol is0.59 cal/g oC.

Substance in blood Typical range

Calcium 8.5 – 10.5 mg/dLSodium 3.10 – 3.33 mg/mL Potassium 137 - 200 mg/LCholesterol 105 – 200 mg/dLFasting glucose 70 – 110 mg/dL Total protein 6.0 – 8.0 g/dL