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Chemistry_Unit 1 Chemistry_Unit 1 Presentation Transcript

  • Welcome to Chemistry! Unit 1
  • So, what is chemistry?
  • The study of matter , its structure, properties and composition, and the changes that matter undergoes.
  • So, why do I need to learn about chemistry?
  • Almost everything we come in contact with is made of materials created or enhanced by chemistry!! Examples… Plastics Clothing Medicine Cosmetics Hygiene Products
  • Branches of Chemistry Organic Inorganic Physical Analytical Biochemistry Theoretical
  • So, what are some careers in chemistry?
    • Law (Environmental or Patent)
    • Pharmacy
    • Space Exploration
    • Forensics
    • Engineering
    • Industry (Paper, Plastics, Ceramics)
    • Medicine
    • Teaching
    • Oceanography
  • Matter and Its Properties
    • Matter is anything that has mass and takes up space.
    • Mass is the measure of the amount of matter.
    • All matter has volume and mass!!
  • Building Blocks of Matter
    • Atom is the smallest unit of an element that maintains the identity of that element.
    • Element is a pure substance that can’t be broken down into simpler substances and is made of 1 type of atom.
    • Compounds are substances that can be broken down into simple substances; made of two or more elements.
  • Properties of Matter
    • Extensive properties
      • Depend on the amount of matter that is present.
      • Volume, mass, etc.
    • Intensive properties
      • Do not depend on the amount of matter that is present.
      • Boiling point, melting point, etc.
  • Properties (con’t)
    • Physical Properties
      • Characteristic that can be observed or measured without changing the identity of the substance.
      • Physical Changes do not involve changing the identity of the substance.
        • Melting, boiling, etc.
    • Chemical Properties
      • Relates to a substance’s ability to undergo changes that transform it into different substances.
      • Chemical Changes are when substances are converted into different substances.
  • Physical Changes
    • Changes of State
      • Solid – definite volume and shape
      • Liquid – definite volume but indefinite shape
      • Gas – neither definite volume or definite shape
  • Chemical Changes
    • Carbon plus oxygen forms carbon dioxide
    • (Charcoal burning…)
    • Reactants are the substances that react, and products are the substances formed by the chemical change.
  • Classification of Matter
    • Mixtures
    • Blend of 2 or more kinds of matter, each with its own identity and properties
    • Can be uniform (homogeneous) or not uniform (hetergeneous)
    • Pure Substances
    • Has a fixed composition
    • Has the same characteristics throughout
  • Elements
    • Elements are pure substances and arranged on based on their chemical properties on the periodic table.
  • Periodic Table
    • Vertical columns are groups
    • Horizontal rows are periods
    Group Period
  • So, how does one use chemistry to make discoveries? The Scientific Method!!!
  • The Scientific Method A logical approach to solving problems by observing and collecting data, formulating hypotheses, testing hypotheses, and formulating theories that are supported by data.
  • Scientific Method
    • State the problem clearly.
    • Gather information.
    • Form a hypothesis.
    • Test the hypothesis.
    • Evaluate the data to form a conclusion.
      • If the conclusion is valid, then it becomes a theory . If the theory is found to be true over along period of time (usually 20+ years) with no counter examples, it may be considered a law .
    • 6. Share the results.
  • Important Things to Consider…
    • Have a control , which is used to show that the result of an experiment is really due to what you are testing.
    • Know your variables , which are the factors that change in an experiment
      • Independent variable – what the experimenter changes
      • Dependent variable – changes because of the experiment
  •  
  • Rules and Tools of the Trade
  • Safety Rules
    • Don’t enter the lab without teacher!
    • You MUST wear safety goggles when conducting experiments!
    • Work in assigned place!
    • Wear lab apron and tie back long hair.
    • Keep lab table clear and clutter-free!
    • Don’t perform unauthorized experiments!
    • Don’t use flames without teacher’s permission!
  • Safety Rules (con’t)
    • No horseplay in the lab!
    • Don’t eat/drink in the lab!
    • Don’t taste chemicals.
    • Wash any chemical that comes in contact with your skin immediately and notify the teacher!
    • Wash your hands well when exiting the lab!
    • Check glassware for cracks!!
  • Safety Rules (con’t)
    • Stay in assigned areas.
    • Use proper techniques.
      • Point test tubes away for self/others when heating.
      • Don’t pour reagents back into bottles.
      • Dispose of materials properly.
      • Clean up spills and accidents.
  • Safety Rules (con’t)
    • 16. Report all accidents and problems to the teacher!!!
    Complete the contract and return tomorrow for a homework grade!!!
  • Lab Equipment
  • Safety Equipment
  • Beaker
    • Thin, glass vessel
    • Holds and heats liquids
  • Bunsen Burner
    • Heating device
  • Ring Stand
    • Iron stand
    • Clamps and rings are placed on it
    • Holds apparatus for experiment
  • Crucible
    • Porcelain cup
    • Used to heat solids to a high temperature
  • Flasks
    • Thin, glass vessels
    • Used to hold/heat liquids
    Erlenmeyer Florence
  • Evaporating Dish
    • Porcelain dish that can be heated to a high temperature
  • Tongs
    • A device used to pick up hot objects
  • Funnel
    • Device that allows one to pour liquids through a small opening.
  • Mortar and Pestle
    • Used to grind solids into a powder
  • Pipestem Triangle
    • Device placed on a ring or tripod
    • Used to hold a crucible when it is heated
  • Test Tube
    • Small glass tube used in most chemical reactions
  • Test Tube Holder
    • Device used to safely hold a test tube as it is being heated.
  • Widemouth Bottle
    • Bottle used to collect gases
    • Can’t be heated!!
  • Graduated Cylinder
    • Used to measure liquids exactly
  • Triple Beam Balance
    • Use to determine the mass of solids
  • Test Tube Clamp
    • Attaches to a ring stand
    • Holds test tubes
  • How does one collect data or determine a result in chemistry?
  • Chemistry is a QUANTITATIVE science, meaning that we describe most things by using numbers!!!
  • Scientific Notation
    • Scientists often work with very large and very small values.
    • Example
    • The mass of the Earth
    • 6,000,000,000,000,000,000,000,000 kg
  • Scientific Notation (con’t)
    • To make numbers more manageable, scientists place numbers in a shortened form.
    • It is based on the exponential notation. The numerical part of a measurement is expressed as a number between 1 and 10 multiplied by a whole-number power of 10.
    • M x 10 n
  • Scientific Notation (Examples)
    • The mass of a softball is 180 grams or 1.8x10 2 g.
    • 2,000 meters can be written as 2x10 3 m.
    • 0.003 kilograms can be written as
    • 3x10 -3 kg.
  • Negative vs. Positive Exponents
    • To determine if the exponent is negative or positive, remember this…
    • Whole numbers will have positive integers.
    • Decimal numbers will have negative integers.
  • Practice
    • 3,000 m
    • 1,000,000 km
    • 0.009 cm
    • 0.00065 dm
  • Removing from Scientific Notation
    • To take a number OUT of scientific notation, simply move the decimal the same number of places denoted by the integer.
    • Negative integers move the decimal to the left.
    • Positive integers move the decimal to the right.
  • Practice
    • Examples
    • 1. 3.1 x 10 -2 dm = 0.031 dm
    • 2. 6.5 x 10 7 mm = 65,000,000 mm
    • Practice
    • 1. 7.8 x 10 5 m
    • 2. 9 x 10 -6 dm
  • Calculations with Scientific Notation
    • Multiplication
    • When multiplying numbers in scientific notation, multiply the first part of the number and ADD the exponents!
    • (2.0 x 10 2 )(4.0 x 10 3 ) =
    • 2.0 x 4.0 = 8.0
    • 2 + 3 = 5
    • 8.0 x 10 5
  • Calculations with Scientific Notation
    • Division
    • When dividing numbers in scientific notation, divide the first part of the number and SUBTRACT exponents.
    • 8.0 x 10 5
    • 2.0 x 10 3
    • 8.0/2.0 = 4.0
    • 5 – 3 = 2
    • 4.0 x 10 2
  • Significant Figures/Digits Valid Digits/Figures
  • Rules for Significant Figures
    • Digits other than zero are always significant.
    • Examples
    • 96 g = 2 significant
    • 61.4 g = 3 significant
    • 0.52 g = 2 significant
    • One or more final zeros used after the decimal point are always significant.
    • Examples
    • 4.72 g = 3 significant
    • 4.7200 km = 5 significant
    • 82.0 m = 3 significant
    • Zeros between two other significant digits are always significant.
    • Examples
    • 5.029 m = 4 significant
    • 306 km = 3 significant
    • Zeros used solely for spacing the decimal point are not significant. The zeros are placeholders only.
    • Examples
    • 7000 g = 1 significant
    • 0.00783 kg = 3 significant
  • Arithmetic with Significant Digits
    • Addition and Subtraction
    • - Lease precise value
    • - Example
    • 24.686 + 2.343 + 3. 21 = 30.239
    • = 30. 2
    • Division and Multiplication
    • - Fewest digits
    • - Example
    • 36.5 m/3.414 s = 10.69 m/s = 10.7 m/s
  • Learning Check
    • What are some U.S. units that are used to measure each of the following?
    • A. length
    • B. volume
    • C. weight
    • D. temperature
  • Solution
    • Some possible answers are
    • A. length - inch, foot, yard, mile
    • B. volume - cup, teaspoon, gallon, pint, quart
    • C. weight - ounce, pound (lb), ton
    • D. temperature -  F
  • Standards of Measurement
    • When we measure, we use a measuring tool to compare some dimension of an object to a standard.
    For example, at one time the standard for length was the king’s foot. What are some problems with this standard?
  • SI measurement
    • Le Système International d‘Unités
    • Among countries with non-metric usage, the U.S. is the only country significantly holding out . The U.S. officially adopted SI in 1966.
  • SI Base Units Quantity Symbol Unit Abbreviation Length l Meter m Mass m Kilogram kg Time t Second s Temperature T Kelvin K Amt. of Substance n Mole mol Electric Current I Ampere A Luminous Intensity I v Candela cd
  • Mass vs. Weight
    • Mass: Amount of matter (grams, measured with a BALANCE)
    • Weight: Force exerted by the mass, only present with gravity (pounds, measured with a SCALE)
    Can you hear me now?
  • Derived Units Combination of SI base units Area Volume Density
  • SI Prefixes Table 5.  SI prefixes Factor
  • Converting Among Units
    • There are two ways to convert among units:
    • Moving the decimal
    • Factor-label method
  • Moving the Decimal
    • 100 cm  m
    • Step 1
    • Look at the unit that your problem is stated in and the unit that your answer is to be put in
    • cm  m
    • Step 2
    • Determine if you are going from a large unit to a small unit OR a small unit to a large unit.
    • cm  m
    • Small unit  Large unit
    • Step 3
    • Determine the way the decimal will move.
    • If you are moving to a R educed unit, move R ight.
    • If you are moving to a L arger unit, move L eft.
    • Cm  m
    • Small  L arge
    • Move L EFT!
    • Step 4
    • Determine the number of places the decimal must move!
    • Use the SI Prefixes-Table 2 (p. 35)
    • 1 centimeter = .01 meter
    • OR
    • 100 centimeter = 1 meter
    • The decimal will move the number of 0’s, which is two!
    • Step 5
    • Move your decimal!
    • 100 cm = 1 m
  • Practice
    • 10000 dm  m
    • 100 m  km
    • 10 km  m
    • 10 km  cm
  • Factor-Label Method
    • 16 m  mm
    • Step 1
    • Look at the units and where you are starting and where you are finishing.
    • m  mm
    • Step 2
    • Write down the conversion factor(s).
    • 1000 mm  1 m
    • Step 3
    • Step up a problem:
    • ALWAYS start with what you are given !
    • Then, add in conversion factor(s).
      • 16 m x 1000 mm =
      • 1 m
    • Step 4
    • Cancel out like values.
      • 16 m x 1000 mm =
      • 1 m
    • Step 5
    • Run through your calculator (or brain).
    • 16 m x 1000 mm = 16,000 mm
    • 1 m
    • Practice :
    • 58 ns  s
    • 9270 mm  m
    • 12.3 ks  s
    • 15.5 s  ks
  • How do you measure up?
  • Reading a Meterstick
    • . l 2 . . . . I . . . . I 3 . . . .I . . . . I 4 . . cm
    • First digit (known) = 2 2.?? cm
    • Second digit (known) = 0.8 2.8? cm
    • Third digit (estimated) between 0.05- 0.08
    • Length reported = 2.75 cm
    • or 2.74 cm
    • or 2.76 cm
    Let's Try It!!
  • Stating a Measurement
    • In every measurement there is a
    • Number followed by a
    • Unit from a measuring device
    • The number should also be as precise as the measurement!
  • Three targets with three arrows each to shoot. Can you hit the bull's-eye? Both accurate and precise Precise but not accurate Neither accurate nor precise How do they compare? Can you define accuracy and precision?
  • Accuracy vs. Precision
    • Accuracy
      • How close a measurement is to the true correct value for the quantity
    • Precision
      • How close a set of measure-ments for a quantity are to one another, regardless of whether the measurements are correct