CHEMISTRY

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CHEMISTRY Chapter 1 2

• Matter, Measurements, and Calculations

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Chapter 1 Section 1

• Objectives
• Define chemistry
• List examples of branches of chemistry
• Compare and contrast basic research, applied
  research, and technological development

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What objects in this room are related to
chemistry?

• Plastics
• Fabrics
• Clothes
• Cooking oil
• Motor oil
• Make-up
• Radio
• Batteries
• Computers

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Chemistry in our daily lives.

• Antibiotics
• Food
• Transportation
• Sports
• Farming
• Military
• Industry

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Chemistry

• Study of the composition and properties of matter
  and the changes that matter undergoes
• What something is made of
• What is the internal arrangement

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Chemical

• Any substance that has a definite composition

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6 Main Branches of Chemistry

1. Organic substances containing C
2. Inorganic substances other than organic
3. Biochemistry living things
4. Physical chemistry changes of matter
5. Analytical chemistry id components of materials
6. Theoretical chemistry use math and computers to
   understand chemical behavior

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All branches involve some type of research.

• Basic research to increase knowledge
• how and why
• Applied research to solve problems
• Technological development production and use of
  products
• - lags behind discoveries
• - application of knowledge

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Review and Assignment

• 1. Define chemistry
• 2. List examples of branches of chemistry
• 3. Compare and contrast basic research, applied
  research, and technological development
• Assignment WS 1-1

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Quiz

1. Name two branches of chemistry.
2. List two ways that chemistry affects our daily
   lives.
3. Definition of chemistry.

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Chapter 1 - Matter
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Chapter 1 Section 2

• Objectives
• Distinguish between a mixture and a pure
  substance.
• Define what matter is.

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Matter

• anything that has mass and occupies space
• includes almost everything
• exceptions are light, heat, and sound
• properties are used to measure matter
• ex. mass
• Mass measure of quantity of matter
• - not affected by temp, location, or any other
  factor

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Demo.

• Mass vs. matter
• What caused the change in mass?
• Is air matter?

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Matter (cont.)

• Classified into 2 groups
• 1. pure substances
• 2. mixtures
• Pure substance matter that has the same
  properties throughout
• ex. element or compound

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Pure Substances

• Element substance that cannot be broken down by
  ordinary chemical change
• - only 1 type of atom
• - symbols abbreviated w/1 or 2 letters
• - can be an allotrope
• allotrope one of a number of different
  molecular forms of an element in the same
  state
• Compound substance made up of 2 or more
  elements chemically combined
• - can be broken down by chemical change
• - more than 1 type of atom

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Compounds

• Elements that make up a compound are combined in
  definite proportion by mass
• ex. 100 g water has 11.2 g H and 88.8 g of O
• 2. Chemical and physical properties of compound
  differ from those of its parts
• ex. water is liquid, H and O are gases
• 3. Compounds can be formed from simpler
  substances by chem change and can be broken down
  into simpler substances

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example

• 100 of water has 11.2 g H and 88.8 g O
• How many g of H is in a 120g sample of water?
• 120 g water 11.2 g H 13.4 g H
• 100 g water

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Mixtures

• - contain 2 or more substances that have
  different properties
• - vary in composition and properties from sample
  to sample
• ex. rock, wood, salt water
• Not chemically combined
• Can be separated by simple physical means
• ie. filtration, evaporation, distillation

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Formation of Mixtures

• A mixture can be formed 3 ways
• Element mixed w/1 or more other elements
• ex. carbon w/sulfur
• 2. Compound mixed w/ 1 or more other compounds
• ex. salt w/sugar
• 3. 1 or more elements mixed w/1 or more compounds
• ex. sulfur w/sugar

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Characteristics of Mixtures

• - retain properties of each of its parts
• ex. iron and sulfur
• - iron remains magnetic
• - composition can vary widely
• - can be homogeneous or heterogeneous

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Types of mixtures

• Homogeneous uniform composition throughout
• - called solutions
• ex. alloys, pop, air, coffee
• Heterogeneous not uniform throughout
• ex. concrete, soil, dry soup, spaghetti and meat
  balls

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Review and Assignment

• 1. Distinguish between a mixture and a pure
  substance.
• 2. Define what matter is.
• Assignment WS

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Chapter 1 Section 2

• Objectives
• 1. Distinguish between the physical properties
  and chemical properties of matter.
• 2. Classify changes of matter as physical or
  chemical.
• 3. Explain the gas, liquid, and solid states in
  terms of particles.

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Properties of Matter

• allow us to distinguish btwn substances
• characteristics of a substance
• what can be observed
• way that a substance behaves
• ex. color, taste, odor, gas, liquid, solid

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Properties (cont.)

• - can be extensive or intensive
• Extensive d/o amount of matter
• ex. volume, weight, mass, and E
• Intensive does not d/o amount of matter
• ex. melting point, boiling point, density, and
  conductivity

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Demonstration

• Properties
• - water and glycerin
• How do they compare?
• - look, feel, weight, flow
• - water and salt water
• How do they compare?
• - conductivity

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Physical Properties

• Can be observed or measured w/out changing the
  substance
• Can describe the substance
• Odor, taste, hardness, density, melting point,
  and boiling point
• Metals ductile (pulled into wire), malleable
  (hammered into sheets), luster (shine), good
  conductors

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Chemical Properties

• A transformation of a substance into a different
  one
• rusting, flammability, tarnishing, new substance
  formed

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Physical Change

• No new substance is formed
• CHANGE IN PHASE, pounding, grinding, cutting
• Changes of phase
• When a substance changes phase there is no change
  in composition
• Physically different, chemically the same
• Solid, liquid, or gas are the three states of
  matter

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States of Matter

• Solid definite volume and shape
• Particles are in fixed positions
• Held w/strong attractive forces
• Liquid definite volume and no definite shape
• Takes shape of container
• Particles can move past each other

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States of Matter (cont.)

• Gas neither definite volume nor definite shape
• Particles move easily and are very far apart
• Plasma high temperature state in which atoms
  lose their electrons

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Chemical Change

• One or more substance is changed to something new
• Rusting, burning, gas formed, digestion, heat or
  light added, explosion, color change, odor
  change, water formed

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Review and Assignment

• 1. Distinguish between the physical properties
  and chemical properties of matter.
• 2. Classify changes of matter as physical or
  chemical.
• 3. Explain the gas, liquid, and solid states in
  terms of particles.
• Assignment p. 18 and WS

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CHEMISTRY Chapter 1 Section 3

• Objectives
• Perform density calculations.
• Describe conservation of mass.

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Properties of Matter

• E is always involved in both physical and
  chemical changes
• Physical are not at noticable
• Chemical are more noticable
• Heat and light are given off

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Density

• is a physical property
• is always the same for a solid substance
• in gases and some liquids a change in temperature
  will change the density
• increase in temperature will decrease density
• D m/V

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Density problem

• Use the 5 steps in problem solving to solve the
  following problem.
• Lead has a mass of 22.7 g and its volume is 2.00
  cm3. What is its density?
• m 22.7 g V 2.00 cm3
• D m/V 22.7 g/2.00 cm3 11.4 g/ cm3

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Examples
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Conservation of Mass

• In reactions matter cannot be created or
  destroyed by a chemical change
• - mass stays the same, it may just change form

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Density Lab Results

• Group 1
• Group 2
• Group 3
• Group 4
• Group 5 -

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Review and Assignment

• 1. Perform density calculations.
• 2. Describe conservation of mass.
• Assignment WS and Density lab

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Chapter 2 - Sec.1

• Objectives
• Describe the purpose of the scientific method.
• Distinguish between qualitative and quantitative
  observations.
• Describe the steps to making a graph.
• Distinguish between inversely and directly
  proportional relationships.

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Scientific Method

• - a logical approach to solving problems
• 1. Make observations
• observe your surroundings
• 2. State the problem
• - stated as a question
• 3. Collect data
• 4. Form hypothesis
• - testable statement
• 5. Test hypothesis
• 6. Conclusion
• 7. Modify hypothesis and retest

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Observing

• Involves making measurements and collecting data
• Data can be qualitative or quantitative
• Qualitative non-numerical information
• - descriptive (the sky is blue)
• Quantitative numerical information - the mass
  is 25.7 grams

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Conclusion

• Can be explained by using models
• Model explanation of how phenomena occur or how
  things are related
• - visual
• - verbal
• - mathmatical

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Theory

• models may become part a theory
• Theory broad generalization that explains facts
  or phenomena
• - must be able to predict results
• ex. kinetic-molecular theory
• collision theory

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Controlled Experiments

• Use manipulated variable (independent)
• Use responding variable (dependent)
• One variable manipulated at a time
• Measurements are called data

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Making a Graph

• Shows results of an experiment in a meaningful
  pattern
• Dependent variable is on the vertical axis
• 1. Always include a title
• 2. Determine variables
• 3. Set up scale
• 4. Plot points
• 5. Draw best-fit line

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Oxygen obtained from electrolysis of water
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Relationships in graphs

• Directly proportional if dividing one by the
  other gives you a constant value
• If one increases so does the other
• If started at point (0,0)
• Inversely proportional if their product is
  constant
• If one increases the other decreases
• Produce a curve

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Review and Assignment

• Describe the purpose of the scientific method.
• Distinguish between qualitative and quantitative
  observations.
• Describe the steps to making a graph.
• Distinguish between inversely and directly
  proportional relationships.
• Assignment graphing WS

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Quiz

1. List three steps of the scientific method.
2. List two steps in making a graph.

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Chapter 2 Sec.2

• Objectives
• 1. Distinguish between a quantity, a unit, and a
  measurement standard.
• 2. Name SI units for length, mass, time, volume,
  and density.
• 3. Distinguish between mass and weight.

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Measurements

• Basic part of science
• Make observations more meaningful
• Needs to be more than just a number or quantity
• Need a common system of units
• For consistency
• Measure your desk w/anything you have available

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SI System

• The International System of Units
• Used in all science
• A standard
• Based on 10
• Makes it easier to convert from one unit to
  another

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SI System (continued)

• 7 base units
• 1. Length meter (m)
• 2. Mass kilogram (kg)
• 3. Time second (s)
• 4. Amount mole (mol)
• 5. Temperature Kelvin (K)
• 6. Electric Current ampere (amp)
• 7. Luminous intensity candela (cd)

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Weight vs. mass

• Mass quantity of matter
• - how much space it takes up
• - measured w/a balance
• - unit kg
• Weight F gravity pulls on matter with
• - measured w/spring scale
• - unit Newton
• On the moon will our weight or mass stay the same?

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SI Prefixes

• You must know these.
• Kilo- 1000
• Deca 10
• Base unit (m, s, L)
• Centi 1/100 or 0.01
• Milli 1/1000 0r 0.001

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Derived Units

• combination of base units
• Examples
• - Area m2
• - Volume m3
• - Density kg/m3
• - Newton m?kg/s2

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Derived Units (cont.)

• Area determined by multiplying 2 lengths
• Volume determined by multiplying 3 lengths for
  a solid
• - for liquids unit is cm3 or mL
• 1 mL 1 cm3

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Review and Assignment

• 1. Distinguish between a quantity, a unit, and a
  measurement standard.
• 2. Name SI units for length, mass, time, volume,
  and density.
• 3. Distinguish between mass and weight.
• Assignment WS 2-2 and p. 42 1-3

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Quiz

1. What is the base SI unit for mass?
2. Kilo ______
3. Centi _____
4. What is a derived unit?
5. 1 cm3 _____ mL

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Chapter 2 - Sec.3

• Objectives
• Distinguish between accuracy and precision.
• Determine the number of significant figures in
  measurements.
• Perform mathematical operations involving
  significant figures.

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Accuracy and Precision

• Accuracy closeness of a measurement to correct
  value
• Precision closeness of a set of measurements to
  each other
• Consistency
• Do not have to be correct
• d/o measuring instrument

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Bullseyes
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Significant Figures

• digits in a measurement that are know with
  certainty and one digit that is estimated
• CALCULATORS DO NOT KEEP TRACK OF SIGNIFICANT
  FIGURES

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Significant Figure Rules

• 1. Digits other than zero are ALWAYS significant
• ex. 61.4 3 sig. fig.
• 2. All zeros at the end of a number and to the
  right of the decimal with a preceding the
  decimal are ALWAYS sig
• ex. 4.7200 km 5 sig. fig.
• 3. Zeros used only for spacing are NOT
  significant
• ex. 7000 1 sig. fig.
• 20 1 sig. fig.
• 100.0 4 sig. fig.
• 4. Zeros between sig. fig are significant
• 5. Zeros in front of a non-zero are NOT sig.
• - dont count until you get to 1st non-zero from
  lf to rt
• 0.004 1 sig. fig.
• 0.0009 1 sig. fig.

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Significant Figures

• 1,000 _____ sig figs
• 100.0 _____ sig figs
• 0.00012340 _____ sig fig
• 10.0340 _____ sig fig

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Calculating w/Significant Figures

• Addition and Subtraction
• - use same of decimal places as the
  measurement w/the least decimal places
• ex. 2.098 3 DECIMAL places
• 6.2 1 DECIMAL place
• 8.298 round to 1 Decimal
• 8.3 is the final answer

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Adding and Subtracting

• 10.0 123 _____
• 23.456 23.0 _____
• 100.12 56.45 _____
• 1,000 12.234 _____

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Calculating w/sig. figs (cont.)

• Multiplication and Division
• - use same sig. fig. as the measurement w/the
  least sig. fig.
• ex. 2.38 3 sig. fig
• x 9.0 2 sig. fig
• 21.42 round to 2 sig. Fig
• 21 is the final answer

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Multiplying and Dividing

• 100.0 x 10 _____
• 34.56 x 23.45 _____
• 12.045 x 34.008 _____
• 50.04 x 23 _____

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Review and Assignment

• Distinguish between accuracy and precision.
• Determine the number of significant figures in
  measurements.
• Perform mathematical operations involving
  significant figures.
• Assignment WS 2-6 and sig fig WS

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Quiz

• How many significant figures are in the following
  numbers?
• 1. 8,000 _____
• 2. 100.01 _____
• 3. 0.00056_____
• 4. 4500.10 _____
• 5. What is precision?

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Chapter 2 - Sec.3 Day 2

• Objectives
• Perform mathematical operations involving percent
  error.

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Percent Error

• Observed value based on lab measurements
• True value based on generally accepted
  references
• Error exists in any measurement
• d/o measurer, instrument, conditions

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Percent Error

• error true value obs. value x 100
• true value
• Example
• atomic mass of Al 28.9 g
• measured mass 27.0 g
• What is the error?
• 28.9 g 27.0 g x 100 7.00
• 28.9 g

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Review and Assignment

• Perform mathematical operations involving percent
  error.
• Assignment WS 2-5 and error WS

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Quiz

• How many significant figures are in the following
  numbers?
• 1. 8,104 _____
• 2. 100.01 _____
• 3. What does error tell us?
• 4. What is accuracy?
• 5. What is precision?

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CHEMISTRY Chapter 2 Sec.3 Day 3

• Objectives
• Use dimensional analysis to convert measurements.
• Convert measurements into scientific notation.
• Perform mathematical operations using exponents.

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Problem Solving Rules

• Write down what is known.
• - mass 346 g volume 34.6 cm3
• 2. Write down unknown.
• - density ?
• 3. Write the equation to use.
• D m/V
• 4. Fill in knowns.
• D 346 g/34.6 cm3
• 5. Solve for unknown and label.
• D 200 g/cm3
• 6. Check your work.

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Dimensional Analysis

• - use with conversion factors to change from one
  unit to another
• Steps
• convert 2550 m to km
• 1. Determine conversion factor
• - 1000 m to 1 km
• 2. Set up T-bars
• 3. Write given in first box

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Dimensional Analysis (cont.)

• 4. Write conversion factor in 2nd box
• - unit on bottom matches unit of given
• 5. Matching labels cancel
• - if 1 from conversion factor is on top divide
• - if 1 from conversion factor is on bottom
  multiple

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Scientific Notation

• Used to represent very large or very small
  numbers
• There are two parts
• Basic form is M x 10n
• M is a number
• n is a number representing how many places to
  move the decimal

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Scientific Notation (cont.)

• If n is negative, your number is a decimal
• If n is positive, your number is a large number
• Examples
• 60,000,000 6 x 107
• 0.000005 5 x 10-6
• 125,000 1.25 x 105

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Scientific Notation (cont.)

• Write the following in scientific notation.
• 1,000,000,000
• 23,456
• 0.0005678
• 0.034
• 14,239.1

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Scientific Notation (cont.)

• Write the following in long hand.
• 1 x 10-9
• 3.5 x 105
• 7.123 x 10-3
• 5 x 102
• 4.56 x 10-2

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Multiplication w/exponents

• Step 1
• Multiply coefficients
• Step 2
• Add exponents
• ex. (2 x 102) (2.5 x 105) 5 x 107

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Division w/exponents

• Step 1
• Divide coefficients
• Step 2
• Subtract exponents
• ex. (5 x 10-2) (1.0 x 107) 5 x 10-9

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Addition Subtraction w/exponents

• All numbers must be written in the same power of
  10
• ex. 5.8 x 103 2.16 x 104
• - change to
• 0.58 x 104 2.16 x 104 2.74 x 104

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Scientific Notation sig figs

• All numbers in front of the x 10 are significant
• ex. 2.00 x 102 3 sig fig
• 2 x 102 1 sig fig

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Scientific Notation calculators

• 5.44 x 107/8.1 x 104
• 5.44 (EE or exp) 7 / 8.1 (EE or exp) 4
• 6.7 x 102

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Review and Assignment

• Use dimensional analysis to convert measurements.
• Convert measurements into scientific notation.
• Perform mathematical operations using exponents.
• Assignment p. 57 1-7 and WS