matter and inmdwom

8/11/2019 matter and inmdwom

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Matter

Matter: Anything that occupies space and has mass.


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Physical Properties

Physical Properties: They can be measured and observed withoutchanging the composition or identity of a substance.

Examples

Odor, Color, Volume, Matter, Density, Melting Point, BoilingPoint


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A Further Breakdown: Extensive

vs. Intensive Physical Properties Extensive Properties: depend on amt of substance (mass,

volume)

Intensive Properties: do NOT depend on amt of substance

(melting point, boiling point)


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Chemical Properties

Properties in which there is a change in composition

Reactivity, flammability, etc.

Subdivided into physical and chemical changes


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Physical

Changes

Physical Change: change in physical properties

Examples

Ice melting, water boiling


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Chemical

ChangesChemical

Changes: Forming new substance(s)

Examples

Rusting of nails, digestion of food in our stomach, the growth of

grass


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PracticeClassify the following as a physical or chemical change or

physical or chemical property:

(a) Gallium metal melts in your hand (and in your mouth).

(b) A Page is White.

(c) Copper sheet acquires a green color over the years.

(d) Milk turns sour.

(e) Wax is melted over a flame.

(f) Propane gas is flammable.


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Pure Substances:

Elements and Compounds Element: A substance that cannot be separated into simpler

substances by chemical means.

Example

Gold and?

Compound: A substance composed of atoms of 2 or moreelements chemically united in fixed proportions.

Example

Sodium Chloride and?
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Mixtures

Mixture: A combination of 2 or more substances in which the

substances retain their identity though no longer seen.

Examples

Air, Soft Drinks, Wine, Coffee, Water pumped from the Earth.

Can you think of anymore?

They can be separated into pure substances:

Elements and/or Compounds.

They can converted into two or more pure substances.
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Mixtures

Homogeneous Mixture: The composition of the mixture, aftersufficient stirring, is the same throughout the solution. Ahomogeneous mixture is called a solution. It has one layer.

Ex: Salt dissolved in water.

Heterogeneous Mixture: The individual components of amixture remain physically separated and can be seen asseparate components. It has more than one layer.

Ex: A glass full of oil and water or sand in a bucket of water.
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PracticeClassify the following as a pure substance, a homogeneous

mixture (solution) or a heterogeneous mixture:

(a) Soda

(b) Kool-Aid

(c) Oil and Vinegar

(d) Common Table Salt (Sodium Chloride)

(e) A vein of gold embedded in quartz


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Separation of Mixtures

Distillation: is the process of vaporizing a liquid in a boiling pot

and then condensing (gas liquid) it again where it will

collect in another vessel.

Used to separate water from dissolved materials (solid or

liquid)

Used to make moon-shine; i.e., separate ethanol from

impurities


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Simple Distillation


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Separation of Mixtures

Filtration: the process of causing a liquid-solid heterogeneous

mixture to encounter a porous barrier so that the liquid passes

through. The solid is left behind.

The liquid that passes through is called the filtrate.

The remaining solid is the residue, or filter cake.

There are two purposes for filtrations:

(1) to remove solid impurities from a liquid.(2) to separate solid products from a liquid.
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Scientific Notation

Handling Numbers Associated with Measurements

Scientific Notation: Expresses a number as a product of a numberbetween 1 and 10 and the appropriate power of 10.

These numbers are very large and very small. They arecumbersome

Example: 702,400,000,000,000,000,000

0.00000000000000000000768
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Using Scientific Notation

1. Any number can be represented as the product of a numberbetween 1 and 10 and a power of 10 (either positive ornegative).

2. The decimal point should be placed with a one non-zero

number to its left.3. The power of 10 depends on the number of places the

decimal point is moved and in which direction.

4. If the decimal point is moved to the left, the power of 10 ispositive. If the decimal point is moved to the right, thepower of 10 is negative.


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Examples

Express 685,000 in scientific notation:

The decimal point must be moved fiveplaces to the left

Thus, the decimal point has one non-zero number to its left

6.85 x 105

Express 0.00000663 in scientific notation:

The decimal point must be moved sixplaces to the right

Thus, the decimal point has one non-zero number to its left

6.63 x 10-6

Try these:

809,000,000,000

0.0000000006


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Fundamental SI Units

Units: The units part of a measurement tells us what quantity is

being used to represent the results of the measurement.

SI = Systeme Internationale (French)

Physical Quantity Name of Unit Abbreviationmass kilogram kg

length meter m

time second s

temperature kelvin Kamount of substance mole mol


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Measurements of

Length, Volume, and Mass Length: Measurement of how long a thing is from end to end. The SI base unit of length is the meter (m).

Volume: Amount of 3-D space occupied by a substance.

Its SI derived unit is m3.

Another common unit of volume is the liter (l).

Mass: Quantity of matter present in an object.

The SI base unit of mass is the kilogram (kg).

Prefixes can be used for all units: i.e., milligram, milliliter, millimeter


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Prefixes used with SI UnitsPrefix Symbol MeaningTera T 1 x 1012

Giga G 109

Mega M 106

Kilo k 103

Deca D 101

deci d 10-1

centi c 10-2

milli m 10-3

micro m 10-6

nano n 10-9

pico p 10-12


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The Use of Prefixes

1 dL = 1 x 10-1L = 0.1 L

1 mg = 1 x 10-3g = 0.001 g

1 km = 1 x 103

m = 1000 m


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Uncertainty in Measurement

Measurements

3.00 cm 3.01 cm 3.02 cm

Notice that the first two digits are the same.

These are called the certain numbers.

The third digit is estimated and can vary.

It is called an uncertain number.

Give the certain and uncertain numbers in the following

measurements:

2.509 kg 1.0596 L


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Precision & Accuracy

Precision: How well measurements agree with

one another

Accuracy: agreement of measurement with

accepted (book) value


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Practice

A 5-page package of high quality printing

paper had its length measured in inches. The

measurements obtained were:

11.003, 11.003, 11.004, 11.003, 11.003

The cover says its length is 11.003 inches.

Do you have good or bad precision?

What about your accuracy: good or bad?


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More Practice

Five blank writable CDs had the same pieceof music burned on to them. The original CDsaid that the track was two minutes and thirty-

three seconds (233) long. However, the length of the track on the burned

CDs was the following:

215, 215, 215, 215, 215

Do you have good or bad precision?

What about your accuracy: good or bad?


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Significant Figures

Significant Figures: Numbers recorded in ameasurement.

(All the certain numbers+the first uncertain

number)

The more significant figures (sig figs) in a

measurement the greater the precision. 32.0 is less precise than 32.000000


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Guidelines for Using

Significant Figures Nonzero Integers: Any digit that is not zero is significant.

Example

894 has _________ significant figures.

2.341 has _________ significant figures.

id li f i


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Significant Figures Leading Zeros: Zeros to the left of the first nonzero digit are not

significant.

They are used to indicate the placement of the

decimal point.

Example

0.07 has __________ significant figures.

0.0000048 has __________ significant

figures.

G id li f i


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Significant Figures Captive Zeros: Zeros between nonzero digits are significant.

Example

707 has ___________ significant figures.

50,001 has __________ significant figures.

G id li f i


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Significant Figures Trailing Zeros:

If a number is greater than 1, then all the zeros written to

the right of the decimal point count as significant figures.

Example3.0 has __________ significant figures.


4.042 has __________ significant figures.7.0000 has __________ significant figures.

8,500 has __________ significant figures.

G id li f U i


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Significant Figures Leading, Captive, and Trailing Zeros:

If a number is less than 1, then only the zeros thatare at the end of the number, and zeros that are

between nonzero digits are significant.Example

0.070 has ___________ significant figures.

0.4006 has ___________ significant figures.0.00520 has __________ significant figures.



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Significant Figures Exact Numbers:

They are assumed to have an unlimited number of

significant figures.


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Significant Figures Numbers With Trailing Zeroes And No Decimal Point:

For numbers that do not contain decimal points, the measurement issaid to be ambiguous.

Example

700: 1, 2, or 3 sig figs?Use Scientific Notation: 7x102has one sig fig.

7.0x102has two sig figs.

7.00 x 102 has three sig figs.

(How many significant figures are in 701? Do you need adecimal pt?)

R di Off N b


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Rounding Off Numbers:

Rules for Rounding Off*We like to reduce our number to fewer digits.*

1. If the digit to be removed is less than 5, then the precedingdigit stays the same. When rounding off, use only the firstnumber to the right of the last significant figure. Do notround off sequentially.

Example

8.934 rounds off to _________ if we only want 2 sig. figs.

R di Off N b


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Rounding Off Numbers

Rules for Rounding Off2. If the digit to be removed is equal to or greater than 5, then the

preceding digit is increased by 1. When rounding off, useonly the first number to the right of the last significant figure.Do not round off sequentially.

Example

8.627 rounds off to ________ if we only want 3 sig. figs.

0.425 rounds off to ________ if we only want 2 sig. figs.

R l f U i Si ifi t


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Rules for Using Significant

Figures in Calculations Addition and Subtraction:

In the answer, the number of sig figs to the right of the decimal pointare determined by the lowest number of sig figs to the right of thedecimal point given by the measurements.

The measurement is said to be limiting. It limits the number ofsignificant figures in the result.

Example

90.442 + 1.1 = 91.542 Rounded Off to 91.5

3.000 - 0.10 = _________ Rounded Off to __________

1081 - 7.25 = _________*For Addition and Subtraction, the decimal points are

counted as sig figs.*

R l f U i Si ifi t


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Figures in Calculations Multiplication and Division:

The number of sig figs is determined by the originalnumber that has the smallest number of sig figs.

The measurement is said to be limiting. It limits the

number of sig figs in the result.

Example

(2.7)x(3.5029) = 9.45783 Rounded Off to 9.5

(7.85)/(124.6) = _____ Rounded Off to ____________*For Multiplication and Division, the wholemeasurements sig figs are counted.*

R l f U i Si ifi


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Figures in Calculations What about:

Order of operations!

Follow the add/sub sig figs for each operation

Then divide, following division sig fig rules

Thus, 7.85 + 11.1 = 19.0

And 124.64 = 121

Therefore, 19.0/121 = 0.157

(7.85 + 11.1)= ?

(124.6 - 4)

P bl S l i d


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Problem Solving and

Dimensional Analysis How do we convert from one unit of measurement to another?

We do this via conversion factors.

For instance:

1 dollar = 100 pennies

Both represent the Same Amount of Money

Conversion factorsallow us to carry out conversions between

different units that mean the same quantity.

They are not taken into sig fig consideration.

Found on A-11 thru A-13.

P bl S l i d


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Problem Solving and

Dimensional AnalysisConvert 57.4 m into mm

Convert 6.1 dm into km

Convert 8.1 m2to cm2

31000mm57.4m 57.4 x 10 mm1m

P bl S l i d


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Problem Solving and

Dimensional AnalysisConvert 1.06 in. into cm

Convert 23.80 L into gal

Convert 7.62 g/mL into oz./gal


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Comparing Temperature Scales


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Temperature Conversions

Converting Between the Kelvin and Celsius

Scales

ToC+ 273.15 = TK

Converting between the Fahrenheit and

Celsius Scales

ToF= 1.80(ToC) + 32


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Temperature Conversions

Convert 172 K to oC.

Convert 41.2oC to oF.

Convert 239.05 oF to K.


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Density

Density: Amount of matter present in a given

volume of substance

Density = mass/volume = g/mL

Not to be confused with weight!


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Example

The volume of a liquid

in a graduated cylinder

is 24.00 ml, and weighs

36.0 grams. What is thedensity of this liquid? m 36.0 g gD = = 1.50

V 24.00 mL mL


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Practice

Mercury has a density of 13.6 g/ml. What

volume of mercury must be taken to obtain

100 grams of the metal?

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matter and inmdwom