Lecture Presentation Chapter 1
Introduction to Matter and Measurement
Dr. Rajani Srinivasan
Tarleton State University
Our journey together in CHEM 105
Total chapters covered 12.5 ?????
Chapters 1-------12
Chapter 13 up till section 13.4
Present Chapter Contents
• Matter• Classification of Matter• Properties of Matter• Units of Measurements• Uncertainty in Measurement• Dimensional analysis
Chemistry
• Study of matter, its properties and behavior• Matter is the central to everything so
chemistry is the Central Science
Chemistry
Health Sciences
AgricultureGeology
Engineering
Matter
“Anything that has mass and occupies space”
Example : anything and everything around you is matter What is the basic building block of Matter ?
ATOMS
Matter
When two or more atoms combined together they are called
We will study about these later……………..Later
molecules
Classification of Matter
Matter
States of Matter Composition
Solid Liquid
gas
Element Compound
Mixture
States of Matter
SOLIDS LIQUIDS GAS
Have definite shapeDefinite volumeNon compressible
No definite shapeDefinite volumeNon compressible
No definite shapeNo Definite volume compressible
Element
Simple definition:Substances that cannot be further decomposed into simpler substances
Molecular levelSubstances made up of one kind of atom.
Example : Oxygen, Hydrogen, Iron etc.
Important facts about Elements
• A total of 118 elements are known• 90% of earths Crust is made up of only five
elements ( Oxygen, Silicon, Aluminum, Iron and Calcium)
• 90% of Human body is made up of (Oxygen , hydrogen and Carbon.
• Please refer to table 1.2 in the book which includes the names and symbols of common element.
Compounds
• Pure substances contains two or more substances
Molecular Level • Made up of two or more types of atoms
Example: Water (H2O), HCl etc.
Important facts
• In a compound each atom is present in a definite composition
• Example: irrespective of the source in water 11% of Hydrogen combines with 89% of Oxygen
“ This is called Law of constant composition or Law of definite proportion”
Mixtures
• Made of two or more substances but each substance retains its identity
• Example AIR – It’s a mixture of all the gasses but all the gasses retain their identity.
• Each substance that forms a mixture is called
“Component” of the mixture.
Types of Mixture
Mixture
Homogenous Heterogeneous
Air , salt , sugar etc. Rocks and woods
Having same composition, properties and appearance
Having different composition, properties and appearance
Physical
• Can be observed without Changing the composition and identity of the substance
Example: color. Odor , melting point , boiling point
Chemical Properties
• These properties describes the way substances completely change or react.
Example : Burning
“Chemical Change is a change in which substances changes into a completely new substances.”
Example : decomposition of water into Hydrogen and Oxygen , formation of water
Intensive Properties
• Properties of Matter which does not depend upon the amount of the examined substance.
• Can be used to identify a substance
Example: Color, odor, melting point etc.
Extensive Properties
• Properties which depend upon the amount of the substance examined
Example: Mass, length, volume etc.
Separation of Mixtures
• Based on the type of Mixture different methods are used for their separation
• Separation is based on the individual properties of the components in the mixture.
Gravity separation (filtration)
Magnetic Separation distillation
Chromatography
Heterogeneous mixture Homogenous mixture
Distillation
Used for separation homogenous mixture Principle : difference in boiling points between the components
Units of measurement
• Every quantitative measurement is represented by a number and a “Unit” depending upon the substance.
• Most commonly used System of measurements “ Metric” developed on France in the 18th century.
• Other type of system used is “English system” commonly used in United states .
SI Units
• “System International” Units derived from Système International d’Unités
• In 1960 an international agreement was reached to use same type of measurement system.
• A total of seven base units are used .
SI units
• Length – meter- English unit = yard• Mass – kilogram English unit = Pound ( 1Kg=
2.2lb)• Temperature – SI unit is Kelvin English unit =
Fahrenheit
Temperature
• Degree of hotness or coldness is called temperature
• Always flows from temperature tohigher
lower
Temperature scales
Celsius or centigrade Kelvin Fahrenheit
Based on boiling (100 degree centigrade) and the freezing point of water (0 degree centigrade)
• It’s the SI unit • Zero kelvin is the
lowest attainable temperature =
- 273.15◦ C• Also called “Absolute
Zero”
• Commonly used in The USA
• Boiling point = 212◦ F• Freezing point = 32◦ F
K = ◦ C + 273.15 ◦ C = 5/9(◦ F-32) or ◦ F = 9/5 ◦ C + 32
Celsius or centigrade Fahrenheit
Derived units
• Those units which are derived from basic SI units or made of more than one basic SI units
1) Speed- – Ratio of distance travelled per unit time – SI units used distance= meter; time= seconds– Represented by
2) Volume- Length Cubed (L*L*L)
-SI unit used – m3, cm3,1L, etc.
ms
Derived units
3) Density
– mass of a substance per unit volume
- SI units used mass= g and Volume = cm3 or 1ml
- Represented byg
cm3
Uncertainty in measurement
Exact Numbers Inexact numbers
Numbers which can easily be counted Example = 12 dozens , 1000g in 1Kg
Numbers obtained by measurements usually include several errors like equipment errors, manual errors . Example: measure the length of a given object using ruler by many students
Precision accuracy
Uncertainty in measurements
Uncertainty in measurement
• Precision – Measure of how closely the individual measurements agree.
Average of several measured values gives us precision • Accuracy – how closely individual
measurement agree with the true value
Properly calibrated instrument gives us the accurate measurement
Significant figures
• All the digits of the measured quantity including the uncertain ones are called “Significant figures”.
• Usually there is a uncertainty in the last digit reported for any measured quantity.
Example: No. of significant figures in the reported measured mass
2.2g = 2
0.4g = 1
1.04g = 3
Rules 1. All nonzero digits are significant
613 has three sig figs
123456 has six sig figs
2. Zeroes between two significant figures are themselves significant.
5004 has four sig figs
602 has three sig figs
6000000000000002 has 16 sig figs!
Rules 3. Zeroes at the end of a number are significant if a decimal point is written in the number• 5.640 has four sig figs• 120000. has six sig figs• 120000 has two sig figs – unless you’re given additional information
in the problem
4.Zeroes at the beginning of a number are never significant.
• 0.000456 has three sig figs• 0.052 has two sig figs• 0.000000000000000000000000000000000052 also has two sig figs!• 2.30 x 10¯5 = 3 sigfig• 4.500 x 1012 = 4 sigfig
Rules
Exact numbers
Exact numbers, such as the number of people in a room, have an infinite number of significant figures.
There are exactly 12 inches in one foot. Therefore, if a number is exact, it DOES NOT affect the accuracy of a calculation nor the precision of the expression.
Example = 100 years in a century
Scientific notation
• Depending upon the scenario or type of experiment significant digits vary .
• To express the number of significant digits we use scientific notation
Example : 200 based on the question
Suppose
we need 2 sig figs = 2.0 * 10 2
We need 3 sig figs = 2.00 * 10 2
Rules
5) When addition or subtraction is performed, answers are rounded to the least significant decimal place.7.939 + 6.26 + 11.1 = 25.299 (this is what your calculator shows) your final answer is limited to one sig fig to the right of the decimal or 25.3 (rounded up).
Rules
6)When multiplication or division is performed, answers are rounded to the number of digits that corresponds to the least number of significant figures in any of the numbers used in the calculation
(27.2 x 15.63) 1.846 = 230.3011918 (this is what you calculator shows)
your final answer is limited to three sig figs, the answer is 230. (rounded down)
Dimensional analysis
• Used to convert one unit to other• Arrive at a proper unit at the end of the
problem• Correct use of conversion factors• Provides systematic way to solve the problem
and detect errors
Conversion factor
• It is a fraction whose numerator and denominator are same quantities expressed in different units (e.g., 1 in. = 2.54 cm)
1 in.
2.54 cm
2.54 cm
1 in.or
Example
• For example, to convert 8.00 m to inches,– convert m to cm– convert cm to in.
8.00 m 100 cm
1 m
1 in.
2.54 cm 315 in.
Conversion using Volume
• density of gold = 19.3 g/cm3
• Volume 2 in3
Find = mass in grams of gold conversion factor required = in3 to cm3
We know 1in = 2.54 cm Therefore 13in3= (2.54)3 cm3 = 16.39 cm3