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John Dalton, 1766-1844

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Antoine Lavoisier, 1743-1794. Joseph Priestly, 1766-1844. Marie Curie, 1867-1934. Dmitri Mendeleev, 1834-1907. John Dalton, 1766-1844. What is Matter? Matter : Anything that occupies space and has mass Energy: Ability to do work, accomplish a change Physical States of Matter

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slide1

Antoine Lavoisier, 1743-1794

Joseph Priestly, 1766-1844

Marie Curie, 1867-1934

Dmitri Mendeleev, 1834-1907

John Dalton, 1766-1844

slide2

What is Matter?

Matter: Anything that occupies space and has mass

Energy: Ability to do work, accomplish a change

Physical States of Matter

Gas: Indefinite volume, indefinite shape, particles far away from each other

Liquid: Definite volume, indefinite shape, particles closer together than in gas

Solid: Definite volume, definite shape, particles close to each other

slide3

Properties of Matter

Property: Characteristic of a substance

Each substance has a unique set of properties identifying it from other substances.

Intensive Properties: Properties that do not depend on quantity of substance

Examples: boiling point, density

Extensive Properties: Properties that depend on or vary with the quantity of substance

Examples: mass, volume

slide4

Physical Properties: Properties of matter that can be observed without changing the composition or identity of a substance

Example: Size, physical state

Chemical Properties: Properties that matter demonstrates when attempts are made to change it into new substances, as a result of chemical reactions

Example: Burning, rusting

slide5

Changes in Matter

Physical Changes: Changes matter undergoes without changing composition

  • Example: Melting ice; crushing rock

Chemical Changes: Changes matter undergoes that involve changes in composition; a conversion of reactants to products

Example: Burning match; fruit ripening

slide6

Classifying Matter

  • Pure substance: Matter that has only 1 component; constant composition and fixed properties
  • Example: water, sugar
    • Element: Pure substance consisting of only 1 kind of atom (homoatomic molecule)
  • Example: O2
    • Compound: Pure substance consisting of 2 or more kinds of atoms (heteroatomic molecules)
  • Example: CO2
slide7

Mixture: A combination of 2 or more pure substances, with each retaining its own identity; variable composition and variable properties

  • Example: sugar-water
    • Homogenous matter: Matter that has the same properties throughout the sample
    • Heterogenous matter: Matter with properties that differ throughout the sample
  • Solution: A homogenous mixture of 2 or more substances (sugar-water, air)
slide8

Measurement Systems

Measurement: Determination of dimensions, capacity, quantity or extent of something; represented by both a number and a unit

Examples: Mass, length, volume, energy, density, specific gravity, temperature

Mass vs. Weight

Mass: A measurement of the amount of matter in an object

Weight: A measurement of the gravitational force acting on an object

slide9

Density: mass divided by volume; d = m/v

Specific gravity: density of a substance relative to the density of water

slide12

Unit of Length

  • Meter = basic unit of length, approximately 1 yard
    • 1 meter = 1.09 yards
  • Kilometer = 1000 larger than a meter
  • Centimeter = 1/100 of a meter
  • 100 cm = 1 meter
  • Millimeter = 1/1000 of a meter
  • 1000 mm = 1 meter
slide13

Unit of Mass

  • Gram: basic unit of mass
  • 454 grams = 1 pound
  • Kilogram: 1000 times larger than a gram
  • 1 Kg = 2.2 pounds
  • Milligram: 1/1000 of a gram
  • Unit of Volume
  • Liter: basic unit of volume
  • 1 Liter = 1.06 quarts
    • 1 Liter = 10 cm x 10 cm x 10 cm
  • 1 liter = 1000 cm3
  • 1 ml = 1 cm3 (1 cc)
slide14

Unit of Energy

  • Joule: Basic unit of energy
  • calorie: amount of heat energy needed to increase temperature of 1 g of water by 1oC
  • 1 cal = 4 joules
  • Nutritional calorie = 1000 calories = 1 kcal = 1 Calorie
  • Units of Temperature

Fahrenheit: -459oF (absolute zero) - 212oF (water boils)

Celsius: -273oC (absolute zero) - 100oC (water boils)

Kelvin: 0K (absolute zero) - 373 K (water boils)

slide16

Converting Celsius and Fahrenheit:

oC = 5/9 (Fo - 32) oF = 9/5 (oC) +32

Converting Celsius and Kelvin:

oC = K - 273 K = oC + 273

Scientific Notation and Significant Figures

Scientific notation: a shorthand way of representing very small or very large numbers

Examples: 3 x 102, 2.5 x 10-4

slide17

The exponent is the number of places the decimal must be moved from its original position in the number to its position when the number is written in scientific notation

  • If the exponent is positive, move the decimal to the right of the standard position
  • Example: 4.50 x 102 450
    • 3.72 x 105372,000
  • If the exponent is negative, move the decimal to the left of the standard position
  • Example: 9.2 x 10-3 .0092
slide18

Practice with Scientific Notation

50,000 = 5.0 x 104 300 =

.00045 = 4.5 x 10-4 .0005 =

3.00 x 102

5 x 10-4

slide19

Significant Figures

  • Significant Figures: Numbers in a measurement that reflect the certainty of the measurement, plus one number representing an estimate
  • Example: 3.27cm
  • Rules for Determining Significance:
  • All nonzero digits are significant
  • Zeroes between significant digits are significant
  • Example: 205 has 3 significant digits
  • 1,006 has
  • 10,004 has

4 sig. figs.

5 sig. figs.

slide20

Leading zeroes are not significant

  • Example: 0.025 has 2 significant digits
  • 0.000459 has 3 significant digits
  • 0.0000003645
  • Trailing zeroes are significant onlyif there is a decimal point in the number
  • Examples: 1.00 has 3 significant figures
  • 2.0 has 2 significant digits
  • 20 has
  • 1500
  • 1.500

4 sig. figs.

1 sig. fig.

2 sig. figs.

4 sig. figs.

slide21

Calculations and Significant Figures

Answers obtained by calculations cannot contain more certainty (significant figures) than the least certain measurement used in the calculation

Multiplication/Division: The answers from these calculations must contain the same number of significant figures as the quantity with the fewest significant figures used in the calculation

Example: 4.95 x 12.10 = 59.895

Round to how many sig. figs.?

Final answer:

3

59.9

slide22

Addition/Subtraction: The answers from these calculations must contain the same number of places to the right of the decimal point as the quantity in the calculation that has the fewest number of places to the right of the decimal

Example: 1.9 + 18.65 = 20.55

How many sig. figs.required?

Final answer:

Rounding Off

Rounding off: a way reducing the number of significant digits to follow the above rules

1

20.6

slide23

Rules of Rounding Off:

Determine the appropriate number of significant figures; any and all digits after this one will be dropped.

If the number to be dropped is 5 or greater, all the nonsignificant figures are dropped and the last significant figure is increased by 1

If the number to be dropped is less than 5, all nonsignificant figures are dropped and the last significant figure remains unchanged

  • Example: 4.287 (with the appropriate number of sig. figs. determined to be 2)
      • 4.287 4.3
slide24

We only use significant figures when dealing with inexact numbers

Exact (counted) numbers: numbers determined by definition or counting

Example: 60 minutes per hour, 12 items = 1 dozen

Inexact (measured) numbers: numbers determined by measurement, by using a measuring device

Example: height = 1.5 meters, time elapsed = 2 minutes

practice
Practice:

Classify each of the following as an exact or a inexact number.

A. A field is 100 meters long.

B. There are 12 inches in 1 foot.

C. The current temperature is 20o Celsius.

D. There are 6 hats in the closet.

Inexact

Exact

Inexact

Exact

slide26

Calculating Percentages

percent = “per hundred”

% = (part/total) x 100

Example: 50 students in a class, 10 are left-handed. What percentage of students are lefties?

% lefties = (# lefties/total students) x 100

= 10/50 x 100

= .2 x 100

= 20%

slide27

Practice Using and Converting Units in Calculations

  • Sample calculation: Convert 125m to yards.
  • Write down the known or given quantity (number and unit)
  • 125 m
  • Leave some blank space and set the known quantity equal to the unit of the unknown quantity
    • 125 m = yards
    • Multiply the known quantity by the factor(s) necessary to cancel out the units of the known quantity and generate the units of the unknown quantity
    • 125 m x 1.09 yards/1 m = yards
slide28

Once the desired units have been achieved, do the necessary arithmetic to produce the final answer

  • 125 x 1.09 yards /1 = 136.25 yards
  • Determine appropriate amount of sig. figs. and round accordingly
  • Fewest sig. figs. in original problem is 3 (from 125), so final answer is 136 yards
slide29

Accuracy vs. Precision

Error: difference between true value and our measurement

Accuracy: degree of agreement between true value and measured value

Uncertainty: degree of doubt in a measurement

Precision: degree of agreement between replicated measurements

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