Matter and Measurement

Unit 1

Introduction Matter Physical/Chemical Properties and Changes Extensive/Intensive Properties Scientific Notation Metric System SI units Conversions Density Significant Figures Uncertainty

Unit Overview

The study of matter and the changes it undergoes.

Major divisions Inorganic Compounds of elements

other than carbon

Organic Compounds of carbon

Biochemistry Compounds of living matter

Physical Theory and concepts

Analytical Methods of analysis

What is Chemistry?

Differing Views

We can explore the MACROSCOPIC world — what we can see —

to understand the PARTICULATE worlds we cannot see.

We write SYMBOLS to describe these worlds.

A Chemist’s View

Matter has mass and occupies space

What is not matter?◦ Everything in the universe is either matter or

energy

What in the world isn’t matter?

Classification of Matter Matter

Pure Substance

Mixture

Element Compound Homogeneous

Heterogeneous

Iron CO2 Juice Trail

Mix

Pure Substances Element

◦ Cannot be converted to a simpler form by a chemical reaction.

◦ Example hydrogen and oxygen

Compound◦ Combination of two or more elements in a

definite, reproducible way.

◦ Example water - H2O

Mixtures

• A combination of two or more pure substances.

◦ Homogeneous - Looks the same throughout

◦ Heterogeneous - Does not look the same throughout

Which are homogeneous or heterogeneous?

◦ Blood Skittles “T-Bone” steak

◦ Orange Juice Vegetable Soup Salad Dressing

Qualitative and Quantitative Analysis

Qualitative analysis is data that is observed Colors, textures, smells, tastes, appearance, etc

Quantitative analysis is data that can be measured

Length, height, area, volume, mass, speed, time, temperature, humidity

Physical properties can be measured or observed

color density odor melting point taste boiling point

Chemical properties describe matter’s ability to change into another substance

ability to burn reactivity ability to decompose

Properties

Changes Physical changes do not change identity of

substance tearing melting freezing boiling grinding cutting

Chemical changes change identity of substance

burning reacting combusting

Extensive and Intensive Properties Extensive properties

Depend on the quantity of sample measured.

Example - mass and volume of a sample.

Intensive propertiesIndependent of the sample size.Properties that are often characteristic of the substance being measured.

Examples - density, melting and boiling points.

Scientific Notation

• Method to express really big or small numbers.

◦ Format is Mantissa x Base Power

We just move the decimal point around

Decimal part oforiginal number

Decimalsyou moved

Really Big Numbers If a number is larger than 1• The original decimal point is moved X places to

the left.

•The resulting number is multiplied by 10X.

•The exponent is the number of places you moved the decimal point.

1 2 3 0 0 0 0 0 0. = 1.23 x 108

Really Small Numbers If a number is smaller than 1• The original decimal point is moved X places to

the right.

•The resulting number is multiplied by 10-X.

•The exponent is the number of places you moved the decimal point.

0. 0 0 0 0 0 0 1 2 3 = 1.23 x 10-7

Using Your Calculator Most calculators use scientific notation

when the numbers get very large or small.

How scientific notation is displayed can vary.

It may use x10n

or may be displayed using an E.

They usually have an Exp or EE◦ This is to enter in the exponent. +

-1

/

x

0

2 3

4 5 6

7 8 9

.

CE

EE

log

ln

1/x

x2

cos tan

English units ◦ Still commonly used in the United States

Examples: pound, inch, foot, cup, pint

Why English system not used in chemistry ◦ Very confusing and difficult to keep track of the

conversions needed◦ Vary in size so you must memorize many

conversion factors

Measurements in Chemistry

Changing the prefix alters the size of a unit.

Metric prefixes

Prefix Symbol Factor

mega M 106 1 000 000

kilo k 103 1 000

hecto h 102 100

deka da 101 10

base - 100 1

deci d 10-1 0.1

centi c 10-2 0.01

milli m 10-3 0.001

micro µ 10-6 0.000001

nano n 10-9 0.000000001

pico p 10-12 0.000000000001

SI - System International - systematic subset of the metric system

Physical Quantity Name AbbreviationMass kilograms kgLength meters mTime seconds sTemperature Kelvin KAmount mole molElectric Current Ampere ALuminous Intensity candela cd

SI Units

Give you ability to convert between units

Problem solving technique (factor label method)

1. Write what you know2. Game plan3. Set up units4. Conversion factors5. Solve

Conversions

Convert 26 gallons to cups◦ Answer: 416 cups

Convert 18 miles to centimeters◦ Answer: 2.9×106 cm

Conversion practice

Allow you to convert between metric prefixes

Write what you know Set up units Bigger unit gets the “1” Smaller unit is 10x where “x” is how many places

apart the two units are

Example◦ Convert .25 kg to mg.

Answer: 2.5x105 mg

Metric Conversions

Ratio of mass to volume of matter

Common units are g / cm3 or g / mL.

Example: what is the density of 5.00 mL of a fluid if it has a mass of 5.23 grams?

d = mass / volume d = 5.23 g / 5.00 mL d = 1.05 g / mL

Example 2: What would be the mass of 1.00 liters of this sample?

Density

Density = Mass

Volume

cm3 = mL cm3 = mL

Measuring Mass Mass - the quantity of matter in an object. Weight - the effect of gravity on an object.

Since the Earth’s gravity is relatively constant, we can interconvert between weight and mass.

The SI unit of mass is the kilogram (kg). However, in the lab, the gram (g) is more commonly used.

Measuring Volume Volume - the amount of space that an object

occupies.

• The base metric unit is the liter (L).

• The common unit used in the lab is the milliliter (mL).

• One milliliter is exactly equal to one cm3.

• The derived SI unit for volume is the m3 which is too large for convenient use.

Relative Densities of Elements

Significant Figures Method used to express accuracy and

precision.

You can’t report numbers better than the method used to measure them.

67.2 units = three significant figures

Significant Figures The number of significant digits is

independent of the decimal point.

255 25.5 2.55 0.255 0.0255

These numbersAll have three

significant figures!

Significant Figure Rules

Leading zeros are not significant.

Leading zeroLeading zero

Captive zeroCaptive zero

Trailing zeroTrailing zero

0.421 - three significant figures

4012 - four significant figures

114.20 - five significant figures

Zeroes between non-zeros are significant.

Trailing zeros are significant ONLY IF there is a decimal point in the number.

How many significant figures are in the following?

Significant Figures

1.0070 m 5 sig figs

17.10 kg 4 sig figs

100,890 L 5 sig figs

3.29 x 103 s 3 sig figs

0.0054 cm 2 sig figs

3,200,000 2 sig figs

Significant Figures and Calculations

123.45987 g+ 234.11 g 357.57 g

805.4 g- 721.67912 g 83.7 g

Addition and subtraction Report your answer with the same

number of digits to the right of the decimal point as the number having the fewest to start with.

Significant Figures and Calculations

Multiplication and division.Report your answer with the same

number of digits as the quantity have the smallest number of significant figures.

Example. Density of a rectangular solid.25.12 kg / [ (18.5 m) ( 0.2351 m) (2.1m) ]

= 2.8 kg / m3

(2.1 m - only has two significant figures)

After calculations, you may need to round off

◦ If the first insignificant digit is 5 or more, you round up

◦ If the first insignificant digit is 4 or less, you round down

Rounding Off

A properly written number in scientific notation always has the proper number of significant figures.

Scientific Notation and Significant Figures

0.00321 = 3.21 x 10-3

Three SignificantFigures

Three SignificantFigures

Measured and Exact Numbers

In science, all of our numbers are either measured or exact.

Exact - Infinite number of significant figures.1 foot = exactly 12 inches

Do not count toward significant figures

Measured - the tool used will tell you the level of significance and varies based on the tool.

When using a measuring tool, record the numbers that are certain and add a guess number

Significant Figures in Measurement

Types of Error Systematic• Errors in a single direction (high or low).

•Can be corrected by proper calibration or running controls and blanks.

Random•Errors in any direction.

•Can’t be corrected. Can only be accounted for by using statistics.

Uncertainty in Measurement

All measurements contain some uncertainty.• We make errors• Tools have limits

Accuracy How close to the true value

Precision How close to each other

Neither accurate nor precise

Precise but not accurate

Precise AND accurate

Types of Error

◦ Instrument not ‘zeroed’ properly◦ Reagents made at wrong concentration

◦ Temperature in room varies ‘wildly’◦ Person running test is not properly trained

Random

Systematic