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Chemistry 101 : Chap. 1 Matter and Measurement

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1 Chemistry 101 : Chap. 1 Matter and Measurement
What is Chemistry and Why we study it (2) Classification of Matter (3) Properties of Matter (4) Units of Measurement (5) Uncertainty in Measurement (6) Dimensional Analysis

2 The Study of Chemistry Chemistry: The study of the properties of matter and the changes that matter undergoes.  Matter : Physical material of the universe Anything that has mass and occupies space  Changes in Matter : Physical or Chemical changes Why Chemistry?  Chemistry is the central science  Chemistry is a practical science and has profound impact on our daily living

3 Macroscopic vs. Microscopic
Macroscopic World : Realm of ordinary-sized object. Things we can see with the naked eye.  (Sub)Microscopic World : Realm of atoms and molecules Carbon nanotube (10-9 m) Chemistry is the science that seeks to understand the properties and behavior of matter (macroscopic) by studying the properties and behaviors of atoms and molecules (microscopic)

4 Major Divisions in Chemistry
 Physical Chemistry (CHM321, CHM420)  Organic Chemistry (CHM211, CHM212)  Inorganic Chemistry (CHM 455, CHM546)  Analytical Chemistry (CHM235, CHM435)  Biochemistry (CHM365, CHM568) All divisions are interrelated and cannot be standing alone.

5 Classification of Matter: pure substance vs. mixture
Pure Substance: A sample of matter that has distinct properties and a composition that doesn’t vary from sample to sample (either element or compound) Elements: A pure substance that cannot be decomposed into simpler substances. The basic unit of an element is an atom. Nitrogen atom Nitrogen molecules Argon gas (atoms) Nitrogen gas (molecules)

6 Classification of Matter: pure substance vs. mixture
Compound : Substances that are composed of two or more elements. The basic unit of compound is a molecule Mixture : Combinations of two or more substances in which each substance retains its own chemical identity. nitrogen atom Two or more elements (compound) Two or more substances (mixture) hydrogen atom ammonium (molecule)

7 Elements  At the present time, there are 116 elements
Periodic Table of the Elements = H2, N2, O2, F2, Cl2, Br2, I2

8 Elements  Not all elements are equal…

9 Compounds  Most elements can interact (or react) with other elements
to form compounds Example: Combine hydrogen & oxygen to generate water Oxygen Hydrogen water However, elemental hydrogen and oxygen exist as diatomic molecules (H2 and O2) in nature. + O H2 2H2O

10 Mixture  Components: The substances making up a mixture
Homogeneous Mixture (solution) : Uniformly distributed throughout. (air, salt solution, sugar solution …) Heterogeneous Mixture : Do not have the same composition, properties and appearance throughout. (rock, wood …) Oil on water Air

11 Classification of Matter

12 Classification of Matter
 Example 14 K gold (2) Orange Juice (3) A cup of coffee (4) Mud

13 Separation of Mixture Separate a mixture into its components by taking advantage of the difference in their properties Filtration : Separation is based on the size of particles in the mixture. Filtration is used with heterogeneous mixtures

14 Separation of Mixture  Distillation : Separation is based
on the boiling points of the components in the mixture. Distillation is typically used with homogeneous solutions. Water changes its states from gas to liquid

15 Separation of Mixture  Chromatography : Separation is based on the solubilities of the components in the mixture. It is normally used with homogeneous mixture. Paper chromatography

16 Classification of Matter:
states of matter States of matter: A sample of matter can have three physically different states  Gas : Indefinite volume and indefinite shape (depends on the volume and shape of its container)  Liquid : definite volume, but indefinite shape.  Solid : definite volume and definite shape Pure substance can have any state depending on the temperature and pressure

17 Three States of Water

18 Properties of Matter Physical properties : They can be measured without changing the identity and composition of the substance Ex. color, order, density, boiling point…  Chemical properties : They describe the way a substance can change or react Ex. flammability, solubility, …

19 Physical vs. Chemical Properties
 Example : Zinc (Zn) silver-grey metal melting point: 420oC reacts with oxygen to form Zinc oxide (ZnO) density (25oC) = 7.13 g/cm3 generates hydrogen when dissolved in sulfuric acid

20 Properties of Matter Extensive properties : Properties that depend on
the quantity of a sample. Ex. Volume :  +  =  V1 + V2 = V1 + V2  Intensive properties : Properties that are independent on the quantity of a sample Ex. Temperature :  +  =  T T T

21 Extensive vs. Intensive Properties
 Example : Boiling/melting point (bp/mp) Mass Density Pressure

22 Changes of Matter Physical changes : Phase changes, but it is
still H2O (no change in its composition) Chemical changes : Aluminum (Al) reacts with Bromine (Br2). (A substance is transformed into a chemically different substance: AlBr3)

23 Units of Measurement : SI Unit
Système International (SI) d’Unités International agreement on the metric units for the uses in science (1960)

24 Units of Measurement : Prefixes
 Prefixes : They are used to indicate decimal fractions or multiples of various units. A Megabyte of memory : 106 bytes of memory Femtochemistry : chemistry that occurs on the time scale of second check out (prof. Zewail’s homepage)

25 Length and Mass Length : 1 meter (m) = 100 cm
Mass : 1 kilogram (kg) = 1000 g Metric to English conversion 1 m = yard 1 cm = inch 1 kg = lb Check out NOTE: Mass and weight are not the same thing. Mass is an intrinsic property of matter, but weight depends on the gravity.

26 Temperature Water freezing Water boiling Celsius scale (oC) 0 100
Fahrenheit scale (oF) oC = 5/9 (oF  32) oF = 9/5(oC) + 32 Kelvin : K = oC (exact) Absolute zero temperature : 0 K =  oC The lowest attainable temperature in our universe

27 Temperature (98.6 oF  32)5/9 = 37 oC 37 oC + 273.15 = 310.15 K
William Thomson Kelvin ( ) “On an Absolute Thermometric Scale” Philosophical Magazine, vol. 1 pp (1848) (98.6 oF  32)5/9 = 37 oC 37 oC = K

28 Derived Units Use the defining equation for the quantity of interest
and substitute the appropriate SI units  Volume: abc = (length)3 = m3 a b c In chemistry, we normally use smaller units. Liter : (10 cm)3 = 1 L = 1 dm3 = 10-3 m3 1 gal = 3.8 L (2) Milliliter = 1 mL = 10-3 L = 1 cm3 = 1 cc

29 Derived Units Density : The amount of mass in a unit volume of
substance SI unit of density  In chemistry, we typically use g/mL = g/cm3 = g/cc  Density depends on temperature  Don’t be confused about density and weight

30 Density, Volume and Mass
(1) 1.00  102 g of mercury occupies a volume of 7.36 cm3. What is the density of mercury? (2) The density of liquid methanol is g/mL. What is the volume of 65.0 g of liquid methanol? (3) The density of gold is g/cm3. What is the mass in gram of a cube of gold if the length of the cube is 2.00 cm?

31 Uncertainty in Measurement
We need to distinguish two different types of number in science  Exact Number : Defined number 1 dozen = 12, 1 m = 100 cm Counted number There are 120 students in the class.  Inexact Number : Numbers from measurement (human errors, machine errors..)

32 Precision and Accuracy
 Precision : How closely individual measurements agree with one another.  Accuracy : How closely individual measurements agree with the correct or “true” value. good precision poor accuracy good precision good accuracy poor precision poor accuracy

33 Significant Figures Measured quantities are generally reported in such a way that only the last digit is uncertain. mass of a dime = g Uncertain. Could be 6 or 4… (2) Sometimes,  sign is used to specify the uncertainty. mass of a dime =  g Significant Figures : All digits of a measured quantity, including the uncertain one. g  5 significant figures

34 Rules for Significant Figures
All non-zero digits are significant (2) Zeros at the beginning of a number are never significant  count the digits starting with the first non-zero digit has TWO significant figures (3) Zeros between non-zero digits are significant  has THREE significant figures (4) Zeros at the end of a number are significant.  has FOUR significant figures 2060 has FOUR significant figures 2.06 x 103 has THREE significant figures

35 Significant Figures in Calculation
The number with the fewest number of significant figures limits the certainty of the calculated quantity. Multiplication & Division : The final answer can have no more significant figures than the fewest number of significant figures in any number in the problem.  Addition & Subtraction : The final answer can have no more decimal places than the fewest number of decimal places in any number in the problem

36 Significant Figures in Calculation
Example 1: Area of a rectangle whose measured edge lengths are cm and 5.2 cm Area = (6.221 cm) x (5.2 cm) = cm2 = Only 2 significant figures Include only 2 significant figures Example 2 : Addition of three measured numbers 20.42 1.322 + 83.1

37 Significant Figures in Calculation
When calculation involves multiple steps… Retain at least one more extra digit (past the number of significant figure) in each step  When you use a calculator… Enter the numbers one after another (without worrying about significant figures) and rounding only the final answer

38 Significant Figures in Calculation
Example 3: [1255  (3.45  108)] = 863  [1255  372.6] = 863  882.4 = = Example 4: (  ) + (2813  12) = = 33890 = From calculator = =

39 Dimensional Analysis We carry units through all calculations. Units behave like numbers: they are multiplied together, divided into each other, or canceled. Example: How many inches are in 10 cm? Correct Wrong Advantages of dimensional analysis (1) It ensures that your answer has the correct unit (2) It makes it easier to find out possible errors

40 Unit Conversion Example: The speed of N2 in air at 25 oC is 515 m/s.
Conversion factor Example: The speed of N2 in air at 25 oC is 515 m/s. Convert the speed into mile/hour

41 Unit Conversion Example: The density of water is 1.00 g/mL.
What is the mass 1.00 gal of water in grams?

42 An example The density of gold is g/cm3. If 2.00 g of gold wire has 0.12 mm radius, how long the wire is?

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