Matter and measurements
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Matter and Measurements. Matter and Energy - Vocabulary. Chemistry Matter Energy Natural Law-(scientific law) Observation, Hypothesis, Theory, Law. States of Matter. Solids. States of Matter. Solids Liquids. States of Matter. Solids Liquids Gases. States of Matter. Change States

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Matter and Measurements

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Matter and measurements

Matter and Measurements


Matter and energy vocabulary

Matter and Energy - Vocabulary

  • Chemistry

  • Matter

  • Energy

  • Natural Law-(scientific law)

    • Observation, Hypothesis, Theory, Law


States of matter

States of Matter

  • Solids


States of matter1

States of Matter

  • Solids

  • Liquids


States of matter2

States of Matter

  • Solids

  • Liquids

  • Gases


States of matter3

States of Matter

  • Change States

    • heating

    • cooling


States of matter4

States of Matter

  • Illustration of changes in state

    • requires energy


Substances compounds elements and mixtures

Substances, Compounds, Elements and Mixtures

  • Substance

    • matter that all samples have identical composition and properties

  • Elements

    • Pure substances that cannot be decomposed into simpler substances via chemical reactions

    • Special elemental forms of atoms (diatomic)

      Elemental symbols

    • found on periodic chart


Substances compounds elements and mixtures1

Substances, Compounds, Elements and Mixtures


Substances compounds elements and mixtures2

Substances, Compounds, Elements and Mixtures

  • Compounds

    • Pure substances composed of two or more elements in a definite ratio by mass

    • can be decomposed into the constituent elements

      REVIEW

    • Element cannot be broken down

    • Compound can be broken down into its elements!


Substances compounds elements and mixtures3

Substances, Compounds, Elements and Mixtures

  • Mixtures

    • composed of two or more substances

    • homogeneous mixtures

      • Uniform throughout

      • Example: solutions

    • heterogeneous mixtures

      • Not uniform

      • Example: rocks


Matter and measurements

Classify the following substances as an element, compound or a mixture (homogeneous or heterogeneous). Which are pure substances?

  • Lightly scrambled egg

  • Water

  • Lava lamp

  • Seawater

  • Chicken noodle soup

  • Root beer

  • Sucrose (C12H22O11)


Separating mixtures

Separating Mixtures

  • Distillation


Separating mixtures1

Separating Mixtures

  • Chromatography

    paper


Chemical and physical properties

Chemical and Physical Properties

  • Extensive Properties - depend on quantity of material

    Ex. mass

  • Intensive Properties - do not depend on quantity of material

    Ex. boiling point


Chemical and physical properties1

Chemical and Physical Properties

  • Chemical Properties - chemical changes

    • Observed during change of material to new material

      • Iron rusting

  • Physical Properties - physical changes

    • No change to the identity of the substance

      • changes of state

      • density

      • color

      • solubility


Physical properties

Physical Properties

  • Density

    • mass / volumeintensive property

    • Mass and volume extensive properties

  • Solubility

    • Amount of substance dissolved in the solvent at a given temperature

      • Saturated solution

      • Unsaturated solution

      • Supersaturated solution


Identify the following as either a chemical or physical change

Identify the following as either a chemical or physical change.

  • Combination of sodium and chlorine to give sodium chloride.

  • Liquefaction of gaseous nitrogen.

  • Separation of carbon monoxide into carbon and oxygen.

  • Freezing of water.


Measurements in chemistry

Measurements in Chemistry

  • length meter m

  • volume liter l

  • mass gram g

  • time second s

  • current ampere A

  • temperature Kelvin K

  • amt. substance mole mol


Measurements in chemistry1

Measurements in Chemistry

  • mega M 106

  • kilo k 103

  • deka da 10

  • deci d 10-1

  • centi c 10-2

  • milli m 10-3

  • micro m 10-6

  • nano n 10-9

  • pico p 10-12

  • femto f 10-15


Units of measurement

Units of Measurement

  • Mass

    • measure of the quantity of matter in a body

  • Weight

    • measure of the gravitational attraction for a body

  • Length

    1 m = 39.37 inches

    2.54 cm = 1 inch

  • Volume

    1 liter = 1.06 qt

    1 qt = 0.946 liter


The use of numbers

The Use of Numbers

  • Exact numbers 1 dozen = 12 things

  • Accuracy

    • how closely measured values agree with the correct value

  • Precision

    • how closely individual measurements agree with each other


The use of numbers1

The Use of Numbers


The use of numbers2

The Use of Numbers

  • Exact numbers 1 dozen = 12 things

    • Counted numbers ex. 3 beakers

  • Significant figures

    • digits believed to be correct by the person making the measurement

  • Scientific notation

    • Way of signifying the significant digits in a number


Significant figures rules

Significant Figures - rules

  • leading zeroes - never significant

    0.000357 has three sig fig

  • trailing zeroes - may be significant

    must specify (after decimal – significant

    before decimal - ambiguous)

    1300 nails - counted or weighed?

    Express 26800 in scientific notation with

    4 sig figs3 sig figs2 sig figs


Significant figures rules1

Significant Figures - rules

  • imbedded zeroes are always significant

    3.0604 has five sig fig

    How many significant figures are in the following numbers?

    0.0124

    0.124

    1.240

    1240


Significant figures rules2

Significant Figures - rules

multiply & divide rule - easy

product has the smallest number of sig. fig. of multipliers


Significant figures rules3

Significant Figures - rules

  • multiply & divide rule - easy

    product has the smallest number of sig. fig. of multipliers


Significant figures rules4

Significant Figures - rules

  • multiply & divide rule - easy

    product has the smallest number of sig. fig. of multipliers


Practice

Practice

  • 142 x 2 =

  • 4.180 x 2.0 =

  • 0.00482 / 0.080 =

  • 3.15x10-2 / 2.00x105 =

  • 24.8x106 / 6.200x10-2 =


Practice1

Practice

  • 142 x 2 = 300

  • 4.180 x 2.0 =

  • 0.00482 / 0.080 =

  • 3.15x10-2 / 2.00x105 =

  • 24.8x106 / 6.200x10-2 =


Practice2

Practice

  • 142 x 2 = 300

  • 4.180 x 2.0 = 8.4

  • 0.00482 / 0.080 =

  • 3.15x10-2 / 2.00x105 =

  • 24.8x106 / 6.200x10-2 =


Practice3

Practice

  • 142 x 2 = 300

  • 4.180 x 2.0 = 8.4

  • 0.00482 / 0.080 = 0.060

  • 3.15x10-2 / 2.00x105 =

  • 24.8x106 / 6.200x10-2 =


Practice4

Practice

  • 142 x 2 = 300

  • 4.180 x 2.0 = 8.4

  • 0.00482 / 0.080 = 0.060

  • 3.15x10-2 / 2.00x105 = 1.58x10-7

  • 24.8x106 / 6.200x10-2 =


Practice5

Practice

  • 142 x 2 = 300

  • 4.180 x 2.0 = 8.4

  • 0.00482 / 0.080 = 0.060

  • 3.15x10-2 / 2.00x105 = 1.58x10-7

  • 24.8x106 / 6.200x10-2 = 4.00x108


Significant figures rules5

Significant Figures - rules

  • add & subtract rule - subtle

    answer contains smallest decimal place of the addends


Significant figures rules6

Significant Figures - rules

  • add & subtract rule - subtle

    answer contains smallest decimal place of the addends


Significant figures rules7

Significant Figures - rules

  • add & subtract rule - subtle

    answer contains smallest decimal place of the addends


Practice6

Practice

  • 416.2 – 10.18 =

  • 16.78 + 10. =

  • 422.501 – 420.4 =

  • 25.5 + 21.1 + 3.201 =

  • 42.00x10-4 + 1.8x10-6 =


Practice7

Practice

  • 416.2 – 10.18 = 406.0

  • 16.78 + 10. =

  • 422.501 – 420.4 =

  • 25.5 + 21.1 + 3.201 =

  • 42.00x10-4 + 1.8x10-6 =


Practice8

Practice

  • 416.2 – 10.18 = 406.0

  • 16.78 + 10. = 27

  • 422.501 – 420.4 =

  • 25.5 + 21.1 + 3.201 =

  • 42.00x10-4 + 1.8x10-6 =


Practice9

Practice

  • 416.2 – 10.18 = 406.0

  • 16.78 + 10. = 27

  • 422.501 – 420.4 = 2.1

  • 25.5 + 21.1 + 3.201 =

  • 42.00x10-4 + 1.8x10-6 =


Practice10

Practice

  • 416.2 – 10.18 = 406.0

  • 16.78 + 10. = 27

  • 422.501 – 420.4 = 2.1

  • 25.5 + 21.1 + 3.201 = 49.8

  • 42.00x10-4 + 1.8x10-6 =


Practice11

Practice

  • 416.2 – 10.18 = 406.0

  • 16.78 + 10. = 27

  • 422.501 – 420.4 = 2.1

  • 25.5 + 21.1 + 3.201 = 49.8

  • 42.00x10-4 + 1.8x10-6 = 4.2 x 10-3


More practice

More Practice

4.18 – 58.16 x (3.38 – 3.01) =


More practice1

More Practice

4.18 – 58.16 x (3.38 – 3.01) =

4.18 – 58.16 x (0.37) =


More practice2

More Practice

4.18 – 58.16 x (3.38 – 3.01) =

4.18 – 58.16 x (0.37) =

4.18 – 21.5192 =


More practice3

More Practice

4.18 – 58.16 x (3.38 – 3.01) =

4.18 – 58.16 x (0.37) =

4.18 – 21.5192 =

-17.3392

Round off correctly


More practice4

More Practice

4.18 – 58.16 x (3.38 – 3.01) =

4.18 – 58.16 x (0.37) =

4.18 – 21.5192 =

-17.3392

Round off correctly to 2 sig. figs

-17


Unit factor method dimensional analysis

Unit Factor MethodDimensional Analysis

  • simple but important way to always get right answer

  • way to change from one unit to another

  • make unit factors from statements

    1 ft = 12 in becomes 1 ft/12 in or 12in/1 ft

    3 ft = 1 yd becomes 3ft/1yd or 1yd/3ft


Unit factor method dimensional analysis1

Unit Factor MethodDimensional Analysis

  • simple but important way to always get right answer

  • way to change from one unit to another

  • make unit factors from statements

    1 ft = 12 in becomes 1 ft/12 in or 12in/1 ft

  • Example: Express 12.32 yards in millimeters.


Unit factor method

Unit Factor Method


Unit factor method1

Unit Factor Method


Unit factor method2

Unit Factor Method

  • Example: Express 323. milliliters in gallons


Unit factor method3

Unit Factor Method

  • Express 323. milliliters in gallons.


Unit factor method4

Unit Factor Method

  • Example: Express 5.50 metric tons in pounds. 1 metric ton = 1 Megagram


Unit factor method5

Unit Factor Method

  • Example: Express 5.50 metric tons in pounds.


Unit factor method6

Unit Factor Method

  • area is two dimensional

  • Example: Express 4.21 x 106 square centimeters in square feet


Unit factor method7

Unit Factor Method

  • area is two dimensional

    express 4.21 x 106 square centimeters in square feet


Unit factor method8

Unit Factor Method

  • area is two dimensional

    express 4.21 x 106 square centimeters in square feet


Unit factor method9

Unit Factor Method

  • area is two dimensional

    express 4.21 x 106 square centimeters in square feet


Unit factor method10

Unit Factor Method

  • volume is three dimensional

  • Example: Express 3.61 cubic feet in cubic centimeters.


Unit factor method11

Unit Factor Method

  • volume is three dimensional

  • Example: Express 3.61 cubic feet in cubic centimeters.


Percentage

Percentage

  • Percentage is the parts per hundred of a sample.

  • Example: A 500. g sample of ore yields 27.9 g of sulfur. What is the percent of sulfur in the ore?


Percentage1

Percentage

  • Percentage is the parts per hundred of a sample.

  • Example: A 500. g sample of ore yields 27.9 g of sulfur. What is the percent of sulfur in the ore?


Derived units density

Derived Units - Density

  • density = mass/volume

  • What is density?

  • Example: Calculate the density of a substance if 123. grams of it occupies 57.6 cubic centimeters.


Derived units density1

Derived Units - Density

  • density = mass/volume

  • What is density?

  • Example: Calculate the density of a substance if 123. grams of it occupies 57.6 cubic centimeters.


Derived units density2

Derived Units - Density

  • density = mass/volume

  • What is density?

  • Example: Calculate the density of a substance if 123. grams of it occupies 57.6 cubic centimeters.


Derived units density3

Derived Units - Density

  • Example: Suppose you need 175. g of a corrosive liquid for a reaction. What volume do you need?

    • liquid’s density = 1.02 g/mL


Derived units density4

Derived Units - Density

  • Example: Suppose you need 175. g of a corrosive liquid for a reaction. What volume do you need?

    • liquid’s density = 1.02 g/mL


Derived units density5

Derived Units - Density

  • Example: Suppose you need 175. g of a corrosive liquid for a reaction. What volume do you need?

    • liquid’s density = 1.02 g/mL


Heat temperature

Heat & Temperature

  • heat and T are not the same thing

    T is a measure of the intensity of heat in a body

  • 3 common T scales - all use water as a reference


Heat temperature1

Heat & Temperature

MPBP

  • Fahrenheit 32oF 212oF

  • Celsius 0oC 100cC

  • Kelvin 273 K 373 K


Relationships of the 3 t scales

Relationships of the 3 T Scales


Relationships of the 3 t scales1

Relationships of the 3 T Scales


Relationships of the 3 t scales2

Relationships of the 3 T Scales


Heat and temperature

Heat and Temperature

  • Example: Convert 111.oF to degrees Celsius.


Heat and temperature1

Heat and Temperature

  • Example: Convert 111.oF to degrees Celsius.


Heat and temperature2

Heat and Temperature

  • Example: Express 757. K in Celsius degrees.


Heat and temperature3

Heat and Temperature

  • Example: Express 757. K in Celsius degrees.


The measurement of heat

The Measurement of Heat

  • SI unit J (Joule)

  • calorie

    1 calorie = 4.184 J

  • English unit = BTU


Synthesis question

Synthesis Question

  • It has been estimated that 1.0 g of seawater contains 4.0 pg of Au. The total mass of seawater in the oceans is 1.6x1012 Tg, If all of the gold in the oceans were extracted and spread evenly across the state of Georgia, which has a land area of 58,910 mile2, how tall, in feet, would the pile of Au be?

    Density of Au is 19.3 g/cm3. 1.0 Tg = 1012g.


Group activity

Group Activity

  • On a typical day, a hurricane expends the energy equivalent to the explosion of two thermonuclear weapons. A thermonuclear weapon has the explosive power of 1.0 Mton of nitroglycerin. Nitroglycerin generates 7.3 kJ of explosive power per gram of nitroglycerin. The hurricane’s energy comes from the evaporation of water that requires 2.3 kJ per gram of water evaporated. How many gallons of water does a hurricane evaporate per day?


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