Matter, Energy and Measurement

Bettelheim, Brown, Campbell and Farrell

Chapter 1

Introductory Material

• Definitions

• Review of basic ideas and skills

• Some examples and practice problems

• Matter: Anything that has mass and takes up space (has volume)

– Mass: measure of amount of matter• Does not change with location

– Weight: measure of pull of gravity on matter• Can change with location

• Chemistry is the science that deals with matter: the structure and properties of matter and the transformations of matter from one form to another

– Chemical change: Matter changes its identity (formula)

• Also called chemical reaction• Old substances used up and new ones formed

– Physical change: Matter keeps its identity• Ex: Dissolving substance or changing state (phase)

between solid, liquid and gas

Scientific Method

Observation

Law

Theory

Experimentation

Direct observations (facts)about the behavior of matter

Summarizes and explains a wide range of observations

A unifying principle that explains abody of observations and the laws basedon them; suggests new experiments

If experiments contradict the theory, the theory may have to be modified or discarded

• Law is summary of observations

• Theory explains those observations– Hypothesis is preliminary

statement of theory)– Model uses pictures or

everyday objects to explain observation

– Always subject to revision as needed

Exponential (Scientific Notation)

• Used to represent very big or very small numbers – uses “powers of 10”

• Examples– Large number: 1234 = 1.234 x 103 (positive

exponent)

– Small number: 0.000456 = 4.56 x 10-4 (negative exponent)

• Write the following in exponential form:

2572

0.00452

• Write the following in decimal form

1.4 x 10-2

1.74 x 103

Significant Digits (Figures)

• Can only have as many digits in measurement as can be read with measuring device – Last digit is “uncertain” (estimated)

• Calculations can only have number of digits justified by measurements

• See Appendix II to review rules for sig figs

Copyright © 2001 The McGraw-Hill Companies, Inc. Permission required for reproduction or display.

Significant FiguresFigure TA 1.2

Significant figures - all digits in a number representing data or results that are known with certainty plus one uncertain digit.

How many sig figs are there in:

1. The number 1.020

2. The number 1020

3. The product of 12.2 x 4.0230

4. The sum of 12.2 + 4.0230

5. The sum of 1.2 x 10-3 + 1.4 x 10-2

Table 1.1, p.7

Table 1.2, p.8

Metric System Unit Prefixes

Metric System Unit Prefixes

giga

mega

kilo

deci

centi

milli

micro

nano

109

106

103

10-1

10-2

10-3

10-6

10-9

G

k

d

c

m

n

Prefix ValueSymbol

M

Conversions factors between English and Metric Systems

Length

1 in.1 m1 mile

===

2.54 cm39.37 in.1.609 km

Mass

1 oz1 lb1 kg

===

28.35 g453.6 g2.205 lb

Volume

1 qt1 gal1 L

===

0.946 L3.785 L33.81 fl oz

1 L = 1.057 qt

Taken from Table 1.3, BBM

Table 1.3, p.8

Unit ConversionsFactor-Label Method

• Conversion factorConversion factor– Ratio, including units, used as a multiplier to

change from one system or unit to another– For example, 1 lb = 453.6 g– Example:Example: convert 381 grams to pounds

– Example:Example: convert 1.844 gallons to milliliters

1.844 gal x 3.785 L1 gal

x 1000 mL1 L

= 6980 mL

381 g x 1 lb453.6 g

= 0.840 lb

Unit ConversionsFactor-Label Method

27 cm = ? in 1 in = 2.54 cm

• Which unit goes on top?– The unit you want to end up with

• Which unit goes on bottom?– The unit you start with

Temperature Scales• Fahrenheit (F):Fahrenheit (F): freezing point of water is 32°F

and the boiling point of water is 212°F

• Celsius (C):Celsius (C): freezing point of water is 0°C and boiling point of water is 100°C

• Kelvin (K):Kelvin (K): zero is the lowest possible temperature; also called the absolute scale– degree is the same size as Celsius degree– K = °C + 273

°F =95

°C + 32_ °C =59

(°F - 32)_

Fig. 1.6, p.11

Convert the following:

25 oC to K

25 oC to oF

67 oF to oC

The Three States of Matter• GasGas

– has no definite shape or volume– fills whatever container it is put into– is highly compressible

• LiquidLiquid– has no definite shape but a definite volume– is slightly compressible

• SolidSolid– has a definite shape and volume– is essentially incompressible

Density

• Density:Density: the ratio of mass to volume

– Commonly used units are g/mL for liquids, g/cm3 for solids, and g/L for gases.

– Example:Example: If 73.2 mL of a liquid has a mass of 61.5 g, what is its density in g/mL?

d = md = densitym = massV = volumeV

d = mV

= 61.5 g73.2 mL

= 0.840 g/mL

Specific Gravity

• Specific gravity:Specific gravity: the density of a substance compared to water as a standard– because specific gravity is the ratio of two

densities, it has no units (it is dimensionless)– Example:Example: the density of copper at 20°C is

8.92 g/mL. The density of water at this temperature is 1.00 g/mL. What is the specific gravity of copper?

= 8.928.92 g/mL1.00 g/mL

Specific gravity =

Energy• Energy:Energy: the capacity to do work

– may be either kinetic energy or potential energy– the calorie (cal) is the base metric unit

• Kinetic energy:Kinetic energy: the energy of motion– KE increases with velocity– KE increases with mass

• Potential energy:Potential energy: the energy an object has because of its position; stored energy

KE =12_ mv2

Energy• Kinetic energy

– Mechanical energy, light, heat, and electrical energy

• Potential energy is – Chemical energy

• Stored in chemical substances• Energy released in chemical reactions

• The law of conservation of energy– Energy cannot be created or destroyed

• Energy can only be converted from one form to another

Heat and Temperature

• Heat is a form of energy

– Heating refers to the energy transfer process when two objects of different temperature are brought into contact

– Heat always flows from the hotter object to the cooler one until the two have the same temperature

Energy Units

• 1 calorie (cal) is the amount of heat needed to raise 1 g of water 1oC

• I kcal = 1000 cal (calorie)

• 1 cal = 4.184 J (joule)

Specific Heat

• Specific heat:Specific heat: the amount of heat necessary to raise the temperature of 1 g of a substance by 1°C.

WaterIceSteamIronAluminumCopperLead

WoodGlassRockEthanolMethanolEther

1.000.480.480.110.220.0920.038

0.420.220.200.590.610.56

Acetone 0.52

Substance SubstanceSpecific Heat(cal/g •°C)

Specific Heat(cal/g •°C)

Specific Heat• The amount of heat change when matter

is heated or cooled is given by the following equation

• Example:Example: how many calories are required to heat 352 g of water from 23°C to 95°C?

Amount of heat = 1.00 calg • °C

x 352 g x (95 - 23)°C

= 2.5 x 104 cal = 25 kcal

Amount of heat = specific heat x mass x change in temperature

= SH x m x (T2 - T1)