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Basic Concepts of Matter

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Closely packed together with a definite ridged shape -Vibrate back and ... Femto- f 10-15. 18. Units of Measurement. Length. Measure of space in any direction ... – PowerPoint PPT presentation

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Title: Basic Concepts of Matter


1
Chapter 1
  • Basic Concepts of Matter

2
Classifying Matter
  • Matter
  • Anything that has mass and occupies space
  • Mass vs. Weight
  • Kinetic-Molecular Theory
  • All matter consists of extremely tiny particles
    in constant motion

3
States of Matter
  • Solid
  • -Closely packed together with a definite ridged
    shape
  • -Vibrate back and forth in a confined space
  • -the particles are not able to move past one
    another
  • Liquid
  • -arranged randomly with a definite volume
  • -fluid
  • -the particles are not confined in space and can
    move past one another
  • Gas
  • -no definite shape or volume
  • -fluid
  • -the particles are far apart and move very
    rapidly colliding with other particles and the
    container walls

4
Categorizing Matter
  • Elements
  • -cannot be decomposed into simpler form via
    chemical reactions
  • -found on periodic chart
  • -atoms are the smallest particle that retains the
    characteristic properties of the elements
  • Pure Substance
  • -consists of all the same substance (pure gold,
    distilled water, etc)
  • -have a set of unique properties that identifies
    it

5
Categorizing Matter
  • Chemical Compounds
  • -two or more elements in a definite ratio by mass
    with unique properties that separate them from
    the individual elements
  • -can be decomposed into the constituent elements
    by chemical reactions
  • -chemical compounds are held together by a
    chemical bound
  • Water hydrogen and oxygen
  • Carbon dioxide carbon and oxygen

6
Categorizing Matter
  • Mixtures
  • two or more pure substances in the same container
  • homogeneous mixtures (solution)
  • -uniform composition throughout
  • -single phase
  • -cannot be separated easily
  • heterogeneous mixtures
  • -nonuniform composition thoughout
  • -easily separated

7
Physical and Chemical Changes
  • Physical changes
  • changes in physical properties
  • -melting, boiling, and cutting
  • Chemical changes
  • changing one or more substances into one or more
    different substances (chemical reaction)
  • 2H2 O2 -gt 2H2O

8
Chemical and Physical Properties
  • Chemical Properties
  • observed during a chemical reaction (change in
    chemical composition)
  • -rusting, oxidation, burning
  • -chemical reactions
  • Physical Properties
  • observed without changing the substances
    composition
  • -allow for identification and classification
  • -density, color, solubility, melting point

9
Classification of Physical Properties
  • Extensive Properties
  • depend on the amount of substance present
  • -mass or volume
  • Intensive Properties
  • do not depend on the amount of substance present
  • -melting point, boiling point, density

10
Density
  • Describes how compact a substance is
  • Who discovered density?
  • Density mass/volume or D m/V

11
Density
  • Example Calculate the density of a substance if
    742 grams of it occupies 97.3 cm3.
  • 1cm3 1mL gt 97.3cm3 97.3mL

12
Density
  • Example You need 125 g of a corrosive liquid
    for a reaction. What volume do you need?
  • liquids density 1.32 g/mL

13
Units of Measure
  • Qualitative measures
  • Nonnumerical experimental observations describing
    the identity of a substance in a sample
  • Quantitative measures
  • Numerical experimental observations describing
    how much of a particular substance is in a sample
  • System International dUnites (SI)
  • measurement system used in the sciences based on
    the metric system

14
Math Review and Measurements
  • We make measurements to understand our
    environment
  • Human senses sight, taste, smell, hearing
  • Our senses have limits and are biased
  • Instruments an extension of our senses meter
    sticks, thermometers, balances
  • These are more accurate and precise
  • All measurements have units
  • METRIC SYSTEM vs. British System

15
SI units
  • Quantity Unit Symbol
  • length meter m
  • mass kilogram kg
  • time second s
  • current ampere A
  • temperature Kelvin K
  • amt. substance mole mol

16
Measurements in Chemistry
  • Name Symbol Multiplier
  • mega- M 106 (1,000,000)
  • kilo- k 103 (1,000)
  • deka- da 10
  • deci- d 10-1 (0.1)
  • centi- c 10-2 (0.01)

17
Measurements in Chemistry
  • Name Symbol Multiplier
  • Milli- m 10-3(0.001)
  • Micro- ? 10-6(0.000001)
  • Nano- n 10-9
  • Pico- p 10-12
  • Femto- f 10-15

18
Units of Measurement
  • Length
  • Measure of space in any direction
  • -derived unit cm
  • -standard length is a meter (m)

19
Units of Measurement
  • Volume
  • Amount of space occupied by matter
  • -derived unit mL or cm3 (cc)
  • -liter (L) is the standard unit

20
Units of Measurement
  • Time (t)
  • Interval or duration of forward events
  • -standard unit is the second (s)
  • Mass (m)
  • measure of the quantity of matter in a body
  • Weight (W)
  • measure of the gravitational
  • attraction (g) for a body (wm x g)

1 kg 1000g 1 kg 2.2 lbs 1 g
1000mg
21
Heat and Temperature
  • Heat (q) vs. Temperature (T)
  • 3 common temperature scales
  • all use water as a reference
  • -Fahrenheit (F)
  • -Celsius (C)
  • -Kelvin (K)

22
Temperature Reference Points
  • Melting Point Boiling Point
  • of water of
    water
  • 32 oF 212 oF
  • 0.0 oC 100 cC
  • 273 K 373 K
  • Body temperature 37.0 oC or 98.6 oF
  • 37.2 oC and greatersick
  • 41 oC and greater, convulsions
  • lt28.5 oC hypothermia

Fahrenheit Celsius Kelvin
23
Temperature Scales
24
Temperature Scales
Fahrenheit to Centigrade Relationships
  • Example Convert 211 oF to degrees Celsius.

Example Express 548 K in Celsius degrees.
25
Precision and Accuracy
Accuracy how closely measured values agree with
the correct value Precision how closely
individual measurements agree with each other
Precise
Accurate
Both
Neither
26
Mathematics in Chemistry
  • Exact numbers (counted numbers)
  • 1 dozen 12 things
  • Measured Numbers
  • Use rules for significant figures
  • Use scientific notation when possible
  • Significant figures
  • digits in a measured quantity that reflect the
    accuracy of the measurement
  • -in other words, digits believed to be correct by
    the person making the measurement
  • Exact numbers have an infinite number of
    significant figures
  • 12.000000000000000 1 dozen

27
Significant figures (numbers/digits)
  • Why use significant numbers?
  • -Calculators give 8 numbers
  • -People estimate numbers differently
  • -Dictated by the precision (graduation) on your
    measuring device
  • -In the lab, the last significant digit is the
    digit you (the scientist) estimate
  • Scientists have develop rules to help determine
    which digits are significant

28
Rules for Significant Figures
  • 1. All Nonzero numbers are significant!!!
  • 2. Leading zeroes are never significant
  • 0.000357
  • 3. Imbedded zeroes are always significant
  • 3.0604
  • 4. Trailing zeroes may be significant
  • - You must specify significance by how the number
    is determined or even written
  • 1300 nails - counted or weighed?
  • 1.30000 How many significant figures?

29
Significant Figures
  • Multiplication Division rule
  • The product retains the number of significant
    figures that corresponds to the multiplier with
    the smallest number of significant figure (sig.
    fig.)

30
Significant Figures
  • Addition Subtraction rule
  • Answer retains the smallest decimal place value
    of the addends.

31
Scientific Notation
  • Express answers as powers of 10 by moving the
    decimal place right (-) or left ()
  • Use of scientific notation is to remove doubt in
    the Significant Figures
  • 2000 ? 2 x 103
  • 15000 ? 1.5 x 10?
  • 0.004 ? 4 x 10-3
  • 0.000053 ? __.__ x 10?
  • In scientific notation, zeros are given if they
    are significant!!!
  • 1.000 x 103 has 4 significant figures
  • 2.40 x 103 has ? significant figures

Key to Sig. Figs Locating the decimal and
deciding when to count the zeros!!!
32
Review 2
  • Units of Measure
  • -length
  • -volume
  • -time
  • -mass
  • -weight
  • Heat vs. Temperature
  • -three temperature scales
  • -temperature conversions
  • Precision vs. Accuracy
  • Significant Figures
  • Scientific Notation

33
Conversion Factors
  • Length
  • 1 m 39.37 inches
  • 2.54 cm 1 inch
  • Volume
  • 1 liter 1.06 qt
  • 1 qt 0.946 liter
  • See Text for more conversion factors

34
Conversion Factors
  • Why do conversions?
  • -Scientists often must convert between units
  • Conversion factors can be made for any
    relationship of units
  • -Use known equivalence to make a fraction that
    can be used to convert from one unit to the
    other

35
Dimensional Analysis
  • 1 inch 2.54 cm
  • Use the ratio to perform a calculation so the
    units will divide out
  • Example Convert 60 inches to centimeters

36
Dimensional Analysis
  • Example Express 9.32 yards in millimeters.
  • 3 ft 1 yard
  • 1 ft 12 in or
  • 1 in 2.54 cm
  • 100 cm 1 m
  • 1000 mm 1 m

37
Dimensional Analysis
  • Example Express 627 milliliters in gallons.

1 liter 1.06 qt 1 qt 0.946 liter
38
Practice on your Own
  • 1kg 2.20 lbs
  • Convert 25 g to lbs
  • Convert 1 mL to Liters
  • Convert 20 meters to cm

39
Area length x width
Dimensional Analysis
  • Area is two dimensional thus units must be in
    squared terms
  • Express 2.61 x 104 cm2 in ft2

40
Volume length x width x height
Dimensional Analysis
  • Volume is three dimensional thus units must be in
    cubic terms
  • Express 2.61 ft3 in cm3
  • this volume is used in medical measurements--cc

41
Percentage
  • Percentage is parts per hundred of a sample
  • x100
  • Example A 335 g sample of ore yields 29.5 g of
    iron. What is the percent of iron in the ore?

g of substance total g of sample
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