Matter and Measurement

http://www.skanschools.org/webpages/rallen/

Matter

Anything that has mass and volume (takes up space)

Matter

Pure substances(can NOT be separated by

physical means)

Each piece looks the same (PURE!)

Each piece has the exact same composition

Mixtures(CAN be separated by

physical means

Each piece is different (not pure)

Pure SubstancesCan NOT be separated

by chemical meansCan be separated by

chemical means, ONLY

Element (simplest form of matter)

Example:

Monoatomic= Na(One atom)

Diatomic= O2

(Two atoms)

Compound or Molecule(2 or more different elements chemically combined)

Example:

NaCl H2O(Table Salt) (Water)

Pure SubstancesParticle Diagrams

Elements Compound

MixturesSame Composition

ThroughoutDifferent Composition

Throughout

Homogenous Mixture

Uniform throughout-distinct pattern

Example: salt water, iced tea

Homo= same

Heterogeneous Mixture

Not uniform throughout- no pattern

Example: Italian dressing, concrete, soil, chocolate chip cookies

Hetero= different

Mixtures

Particle Diagrams

Heterogeneous Homogeneous

Properties of MatterPhysical properties are the constants about a

substance

Can use our senses to observe them

Do not require chemical analysis

Ex: melting point, color, texture

Properties of MatterExtensive Property a property that depends on

how much material you are dealing withEnergy, mass, heat

Intensive Property a property that does not depend on how much material you are dealing with (helps identity matter; a constant about a particular type of matter)

Melting point, boiling point, color, density, hardness, solubility

Properties of MatterChemical properties include behaviors

substances adhere to when they react with other substances

Examples

reacting hydrogen gas with oxygen gas results in a combustion reaction

Very reactive when in the presence of nonmetals

pH

Physical vs. Chemical Changes

Matter is always changing

Physical Change a change that does NOT alter the chemical properties of a substance

Ex. Cutting paper, phase change

Change in size or shape (same composition)

Ex. ice melting to become liquid

Physical vs. Chemical Changes

Chemical Change a reaction in which the composition of a substance is changed

Ex. rusting

Properties different composition

1. Signs of a chemical reaction

Color change

Bubbling/fizzing

Energy produce or consumed

Ex. firewood burning

Do NowChange of Matter Physical or

ChemicalBurning Toast

Making Ice Cubes

Lighting a Candle

Spoiling Milk

Making Kool Aid

Physical

Physical

Chemical

Chemical

Physical

Elements vs. CompoundsElement= formula that contains only one symbol

Compound = formula which contains 2 or more different symbols/elements

Separation of MatterSeparation Apparatus

Type of Separatio

n (physical

or chemical)

Description of

technique

What types of matter

will it separate

Filtration PHYSICAL Filtrate flows through filter paper, undissolved particles (solids) remain on the filter paper

Heterogeneous mixtures or mixtures involving more than one phase

(ex. Sand and water)

Separation of MatterSeparation Apparatus

Type of Separati

on (physical

or chemical

)

Description of

technique

What types of matter will it

separate

Evaporation

PHYSICAL Separate solute (dissolved solid) from solvent (liquid) by boiling solution

• Solute escapes

• Very limited precision

Homogeneous Mixture (solution)

Separation Apparatus

Type of Separatio

n (physical

or chemical)

Description of technique

What types of matter

will it separate

Distillation PHYSICAL Separate solute from solvent by boiling solution and recondensing in receiving flask (both solute and solvent captured)

Separate 2 or more liquids w/different boiling points

Homogeneous (can use to remove impurities from water)

Distillation

Separation Apparatus

Type of Separati

on (physical

or chemical

)

Description of

technique

What types of matter

will it separate

Chromatography

PHYSICAL Separates particles based on

1) Size2) Solubility

Homogeneous

Chemical separation requires reacting a sample with something else in order to turn it into a completely different compound

Scientific NotationMethod for expressing very

large or small numbers easily

Ex.6.02 x 1023 atoms = 1 mole

PracticeWrite the following numbers in scientific

notation

1. 34000000 =

2. 0.0000067 =

3. 25,864 =

3.4 x 107

6.7 x 10-6

2.5864 x 104

Measurements and the Metric System

In chemistry we measure matter using SI units

SI System International

SI Unit Base Units

SI Metric PrefixesTable C

SI Metric Prefixes

Ex. In the word kilometer, the root word (or base unit) is “meter” and the prefix is kilo.

Kilo = 1000

1 km = 1000 m

Conversion FactorsA mathematical expression that relates two units

that measure the same type of quantity

1 min = 60 sec

1000 g = 1 kg

1 L = 1000 mL

Conversion FactorsKilo Hecta Deca base unit deci centi

milli g = gram

m = meter

L = liter

Practice:

3 g = kg 3000 7 m = mm0.007

Dimensional AnalysisWhen you are required to solve a problem with

mixed units, or to convert from one set of units to another

Ex. How many minutes are there in the month of October?

60 minutes x 24 hours x 31 days =

1 hour 1 day

44, 640 minutes

Accuracy vs. PrecisionAccuracy how close your results are to the

desired valueEx. Hitting bulls eye when you’re aiming for it

For most experiments, ACCURATE means +/- 5% from the expected value

Precision how close your results are to one another; how repeatable your results are; consistency/grouping

There are two kinds of numbers in the world:

exact:

example: There are exactly 12 eggs in a dozen.

example: Most people have exactly 10 fingers and 10 toes.

inexact numbers:

example: any measurement.

If I quickly measure the width of a piece of notebook paper, I might get 220 mm (2 significant figures). If I am more precise, I might get 216 mm (3 significant figures). An even more precise measurement would be 215.6 mm (4 significant figures).

Significant FiguresAka Sig Figs

A method for handling UNCERTAINTY in all measurements

This arises due to the fact that we have different equipment with different degrees of ACCURACY

Sig figs are associated with MEASURED VALUES

EXACT NUMBERS do NOT COUNT when determining sig figs

Ex. Atomic masses on the periodic table Conversions 1 in = 2.54 cm

Significant FiguresThe Atlantic/Pacific Method

Determine if a decimal point is present

If yes think “P” for present P = Pacific Coast1. Start at the first nonzero number2. Count all the way to the Atlantic-NO

EXCEPTIONS

If no think “A” for absent A = Atlantic Coast1. Start at first nonzero numberCount all the way to the Pacific-NO EXCEPTIONS

Significant Figures Rules1) ALL non-zero numbers (1,2,3,4,5,6,7,8,9) are

ALWAYS significant.

2) ALL zeroes between non-zero numbers are ALWAYS significant (CAPTIVE Zeros)

Ex. 40.7 L or 87,009 km

3) For numbers less than one, all zeros to the left of the 1st nonzero number are NOT SIGNIFICANT (Leading Zeros)

0.009587 m or 0.0009 kg

Significant Figures Rules4) Zeros at the end of a number and to the right

of a decimal point are SIGNIFICANT (Trailing Zeros)

85.00 g or 9.07000000 L

5) Zeros at the end of a whole number may be significant or not. If there is a decimal after the last zero, they are significant. If there is no decimal after the end zeros, they are NOT significant

2000. m or 2000 m

Significant Figure Rules6) Exponential digits in scientific notation are

not significant 1.12x106 has three significant digits (1, 1, and 2.)

Significant Figures

Number # Sig Figs

48,923 5

3.967 4

900.06 5

0.0004 ( 4 x 10-4) 1

8.1000 5

501.040 6

3,000,000 ( 3 x 106) 1

10.0 ( 1.00 x 101) 3

Using Sig Figs in Calculations

General Rule your final answer must be expressed in the lowest amount of significant figures that were originally given to you

Operation Rule Examples

Multiplication/Division

Perform operation as normal & express answer in the least # of sig figs that were given to you

12.257 x 1.162 =

14.242634

Answer 14.24

Addition/Subtraction

Line decimal points up; round final answer to lowest decimal place

3.9 5 2.8 79+ 213.6 220.4 29

Measuring Matter

Mass vs. Weight

Mass Relationship Weight

How much matter

something has

Directly proportional:

As mass increases weight

increases

Depends on gravity (force

pulling an object toward earth)

Measuring MatterChemistry deals mainly with mass

We have the same mass on earth and the moon. The different forces of gravity on each cause us to weigh more on earth than on the moon

Measuring MatterVolume the amount of space an object takes up

Techniques:Liquids use graduated cylinder, burette (beaker, flask)

Regular Solids measure dimensions and use ( l x w x h)

Irregular Solids displacement method

Measuring MatterDensity amount of mass in a given space;

RATIO of mass to volume

Formula: D = mass Volume

mV