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PPT – Basic Concepts of Matter PowerPoint presentation | free to download - id: 1d0af4-ZDc1Z

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Chapter 1

- Basic Concepts of Matter

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

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

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

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

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

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

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

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

Density

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

Density

- Example Calculate the density of a substance if

742 grams of it occupies 97.3 cm3. - 1cm3 1mL gt 97.3cm3 97.3mL

Density

- Example You need 125 g of a corrosive liquid

for a reaction. What volume do you need? - liquids density 1.32 g/mL

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

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

SI units

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

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)

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

Units of Measurement

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

Units of Measurement

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

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

Heat and Temperature

- Heat (q) vs. Temperature (T)
- 3 common temperature scales
- all use water as a reference
- -Fahrenheit (F)
- -Celsius (C)
- -Kelvin (K)

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

Temperature Scales

Temperature Scales

Fahrenheit to Centigrade Relationships

- Example Convert 211 oF to degrees Celsius.

Example Express 548 K in Celsius degrees.

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

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

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

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?

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.)

Significant Figures

- Addition Subtraction rule
- Answer retains the smallest decimal place value

of the addends.

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!!!

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

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

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

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

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

Dimensional Analysis

- Example Express 627 milliliters in gallons.

1 liter 1.06 qt 1 qt 0.946 liter

Practice on your Own

- 1kg 2.20 lbs
- Convert 25 g to lbs
- Convert 1 mL to Liters
- Convert 20 meters to cm

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

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

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