Unit 1 - Matter

What you should know at the end of this unit.

____ Lab Safety ____ States of Matter

____ Lab Equipment

____ Properties of Matter ____ Classification of

____ SI Units and Prefixes Matter

____ Significant Figures ____ Physical Properties

____ Metric Conversions and Changes

____ Density ____ Chemical Properties

____ Density Calculations and Changes

____ Percent Error

____ Dimensional Analysis

What is Chemistry?

Chemistry is the

study of the

structure, function,

and properties of

matter and the

changes it

undergoes.

What are Chemicals?

Chemicals

are is any

substances

with a definite

composition.

Matter and its Properties

Matter is anything that has mass and

volume.

Example: A piece of chalk has matter but

sunlight does not.

Classification of Matter

Mixtures

Mixtures contain more than one

substance.

Their composition may vary from sample

to sample. Mixtures can generally be

separated by a physical change.

An element is a pure substance made of

only one kind of atom.

Example: Gold, Silver, Sodium, Fluorine

Compounds are substances composed

of atoms of two or more elements

chemically bonded in a specific ratio.

Example: Water is always H2O.

Homogeneous Mixtures

Homogeneous mixtures or solutions

are mixtures with uniform composition.

Heterogeneous Mixtures

Heterogeneous

Mixtures are

mixtures without

uniform

composition. You

can see the different

components that

make up the mixture

Methods for Separating Mixtures

Different methods can be used to separate the parts of a mixture. Most of these methods utilize physical properties.

One method would be filtration.

Distillation

Crude oil goes through fractional

distillation which utilizes differences in

boiling points to separate the different

petroleum products.

Distillation of Petroleum Products

Paper Chromatography

In paper

chromatography,

chemical

interactions

with the paper

make

compounds travel

at

different rates.

Mass is a

measure

of the

amount of

matter in

an object.

An atom is the

smallest

particle of a

chemical

element that

retains its

properties.

Volume is the amount of space occupied

by an object.

States of matter include solid, liquid, gas,

and plasma.

Plasma is a high energy state in which electrons have been knocked off the

atoms.

Example: Plasma is found in fluorescent light tubes and in the sun.

Properties are characteristics of matter

that can be observed.

Example: Water boils at 1000C

and vinegar reacts with baking soda to

form carbon dioxide.

Physical Properties

Physical properties are properties that can be

observed without changing the composition of

the substance.

Boiling Point

Melting Point Color

odor

Physical changes do not change the

identity of the substance.

Example: Cutting a board in half is a

physical change.

State (phase) changes are physical

changes from one state to another.

Chemical Properties

Chemical properties can be observed as

a substance changes into a different

substance.

In a chemical change, a new substance

will be formed.

Chemical Change

Chemical

changes are

chemical

reactions.

During a chemical

reaction one

substance is

changed into

another

substance.

Signs of a Chemical Reaction

1. Heat is given off or

absorbed.

2. Light is given off.

3. Gas or ppt formed

4. Color Change not due to

moisture loss or gain.

5. Odor change.

6. Change is not easily

reversed.

Properties of Substances

• Chemical properties

– Observed when substances take part in a chemical reaction –

• A change that converts it to a new substance.

• Physical properties (intensive, extensive (amount))

– Observed without changing the chemical identity of a substance.

• Melting point

• Boiling point

• Color

• Texture

• Density

Properties of Gold are:

Physical: Melting point of

1063 oC (intensive)

Color gold (intensive)

Amount in weight (extensive)

Chemical: Gold can be stored in air without

reacting chemically with oxygen

The International System (S.I.) is a set

of standard unit of measurement for

scientists throughout the world.

Quantity Unit ABBREV

Length meter m

Mass kilogram kg

Temperature kelvin K

Amount of

Substance

mole mol

S.I. Base Units

The Metric Number Line

We are mostly interested in measurements from the kilo

to milli – region of the number line. Use a device to help

you remember the order of the prefixes. Kind Hearted

Dads Make Dark Chocolate Milk

S.I. prefixes are added to the base units to increase or decrease their value by

powers of 10. Numerical

prefix Symbol Meaning Multiplier Exponential

Multiplier

kilo k thousand 1,000 1 X 103

hecto h hundred 100 1 X 102

deka da ten 10 1 X 101

deci d tenth 0.1 1 X 10-1

centi c hundredth 0.01 1 X 10-2

milli m thousandth 0.001 1 X 10-3

Converting within a specific quantity

requires moving the decimal place.

Example: 546 µm = .000546 m

Example: 0.00056 kL = 560 ml

Example: 1000 g = 1 kg

Uncertainty • Precision:

How closely individual measurements agree with

one another; the “fineness” of a measurement

• Accuracy:

How closely individual measurements agree with

the “true” value

34

Good accuracy

Good precision

Poor accuracy

Good precision

Poor accuracy

Poor precision

Significant Figures or Digits

The number of reliable digits in a

measurement based on accuracy of the

measuring instrument.

The last digit in the number may be an

estimated one.

Rules for Significant Digits

1. Numbers 1-9 are significant.

2. Zeros between other significant digits are

significant.

3. If a zero tells how well something is

measured it is significant.

4. If a zero just tells how big or how small a

number is it is NOT significant.

Examples of Determining Significant

Figures

1. Any nonzero digit is significant.

457 cm 29 cm

2. Any zero between nonzero digits is

significant.

1005 kg 807 kg

3. Any zero at the “beginning” of a number is

not significant; it’s a place holder.

0.0026 Å 0.41 Å 37

= 3 SF = 2 SF

= 4 SF = 3 SF

= 2 SF = 2 SF

4. Any zero at the “end” of a number and after the decimal point is significant.

0.05000 K 3000 K 5. If a number ends with a decimal point, assume

that all digits are significant. 7000. J 20. J 6. For exact numbers (e.g. 4 beakers) and those

used in conversion factors (e.g. 1 inch = 2.54 cm), there is no uncertainty in their measurement. Therefore, IGNORE exact numbers when finalizing your answer with the correct number of significant figures.

• For more practice: http://lectureonline.cl.msu.edu/~mmp/applist/sigfig/sig.htm 38

= 4 SF = 1 SF

= 4 SF = 2 SF

When adding or subtracting significant figures:

Line up the decimals, the last place that is significant in both numbers is where you draw a line. Add or subtract the numbers, then look at the number just past the line if it is bigger or equal to 5 round up if not just drop after the line.

Calculating with Sig Figs

1. Addition & subtraction: a sum or difference may be

no more precise than the least precise measurement.

Consider the fewer number of decimal places.

Ex: 15.047 g + 4.12 g = ?

41

15.047

+ 4.12

19.167

Ex: 25,040 mL + 37,200 mL = ?

25040

+ 37200

62240

→ 19.17 g

62,200 mL

or 6.22 x 104 mL

When multiplying or dividing

significant figures:

Count how many significant figures in each of the numbers being used. Then use the smallest amount of significant figures of all the numbers when reporting the answer.

Calculating with Sig Figs

2. Multiplication & division: a

product or quotient may be no

more significant than the least

significant measurement.

Consider the fewer number of

significant figures.

Ex: 3.000 x 4.00 = 43

12.0

(4 SF) x (3 SF) = (3 SF)

4. Series of operations: keep all non-significant digits during the intermediate calculations, and round to the correct number of SF only when reporting an answer.

Ex: (4.5 + 3.50001) x 2.00 =

44

(8.00001) x 2.00 = 16.0002 → 16

Significant Figure Practice

Site #1 http://www.sciencegeek.net/Chemistry/taters/Unit0Sigfigs.htm

Site #2 http://chemistry.csudh.edu/lechelpcs/sigfigurescsn7.html

Site #3 http://lectureonline.cl.msu.edu/~mmp/applist/sigfig/sig.htm

Operations and Sig Fig’s Practice http://www.teacherbridge.org/public/bhs/teachers/Dana/SigFigOperations.html

Quiz http://antoine.frostburg.edu/chem/senese/101/measurement/sigfig-quiz.shtml

Conversion Factors

• It is important to know how to use conversion

factors.

• This is also called factor labeling or dimensional

analysis.

• How can you convert miles to kilometers?

• How can you convert a quart of

something to liters and reverse?

You need to use conversion units!

Dimensional Analysis or Factor Labeling

A conversion factor is a fraction that I

a ratio equivalent to 1.

Example: 1000 m / 1 km

1 km / 1000 m

Factor – label method or dimensional

analysis is a problem – solving method

that uses conversion factors to change

unit.

• Lets take the first miles to kilometers.

• It I had to travel 75.4 miles, how many kilometers did I travel?

• From the conversion chart we know 1 mile = 1.609 kilometers

• Now set it up as a conversion factor (a fraction)

1.609 kilometers then multiply it by how many miles

1 mile

75.4 miles 1.609 km =

1 mile

= 121 kilometers (notice the 3 sig. fig.’s)

• The second question was how many liters are

there in a quart? Lets find out how many liters

are in 1.000 quarts. Notice the four sig. fig.’s!

• Again look up any conversion factors that you

may have.

• We are in luck the chart says 1 L = 1.057 quart

• Make the factor label needed. We want to get

rid of the quarts so that unit is placed on the

bottom of the fraction.

1 L

1.057 quart

Start with what you are given and use

the fraction.

1.000 quarts 1 liter =

1.057 quarts

= 0.9461 liters

Notice the four sig. fig.’s!

Density

• Is the ratio of mass to volume.

• D = m

V

• Often given in g/mL

Example 1:

A piece of metal is 1.5 cm X 1.5 cm X 1.5 cm.

It has a mass of 30.17 g. Calculate its density.

D = m / V

D = 30.17 g _

1.5 cm(1.5 cm)(1.5 cm)

D = 8.9 g/cm3

Example 2:

What is the volume of a 17.8 g piece of

cobalt that has a density of 8.90 g/cm3 ?

D = m / V

V = m/D

V = 17.8 g

8.90 g/cm3

V = 2.00 cm3 = 2.00 mL

Dimensional Analysis • What is the volume (in3) of a 0.500 lb

sample of Pb? (d = 11.34 g/mL)

• Here many conversions are needed.

Don’t forget the units!

55

3in 22.1219485.1

0.500 lbs Pb |453.59 g

| 1 lb

| 1 mL |

| 11.34 g |

1 cm3 |

1 mL |

1 in3 =

16.4 cm3

Pb