1. Chem 1151: Ch. 1 Matter, Measurements and Calculations

2. What is Chemistry? Chemistry is how matter is organized/reorganized/changed at the molecular level. Matter is everything that isn’t nothing. Matter and energy are two sides of the same coin. Chemistry is often called the central science because it is an essential component of the natural and life sciences.

3. Chemistry and Astronomy Elemental composition of stars can be determined by different wavelengths of visible light emitted. When starlight passes through a planets atmosphere, certain frequencies of light disappear because they are absorbed by compounds in the atmosphere. http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

4. Elements are created from the fusion reactions inside stars. Hydrogen burning – 4 protons become alpha particle (helium nucleus). Helium burning - 3 alpha particles form 12 C then 16 O. Carbon and oxygen burning produce 28 Si, 24 Mg, 32 S, and other elements. Each of these requires more heat than the fusion reaction before it. Uranium is last naturally-occuring element as a result of solar activity. http://www.geophysics.rice.edu/department/faculty/sawyer/ESCI324/esci324_spr_06_lec2.ppt#866,19,Slide 19 Source of Elements

5. Chemistry and Geology Geochemistry: Study of the chemical composition of the earth Chemical transformations in solids – Ex. Polymorphism. – How limestone becomes marble. http://geology.com/rocks/limestone.shtmlhttp://geology.com/rocks/limestone.shtml; http://www.italartworld.com/

6. Polymorphism Graphite: Each Carbon is covalently bonded to 3 other carbons in ring Mohs scale hardness: 1-2 Diamond: Each carbon is bonded to 4 other carbons Mohs scale hardness: 10 T+P Value: $0.10 Value: $1000.00

7. Chemistry and Biology You reach a certain level in biology where processes can only be understood in terms of chemistry. Chemistry in biology explains: – Why your adrenaline levels increase when you are afraid or excited – Why a body fails to produce insulin (diabetes) – Why cells become cancerous – Neurotransmitter (e.g., dopamine, norepinephrine) imbalances that can produce: Euphoria when you have a few beers or fall in love Depression http://cs.fit.edu/~wds/classes/cse5255/cse5255/davis/text.html

8. Centrality of Chemistry Biological processes are driven by chemistry, and the chemical processes are driven in turn by energy processes (physics). Life can only exist because different elements exist and can form different compounds. Different elements exist because fundamental particles can be organized differently to form the different elements inside stars.

9. Matter, Mass and Weight Matter is anything that has mass and occupies space. Mass is a measurement of the amount of matter in an object. Mass is independent of the location of an object. An object on the earth has the same mass as the same object on the moon. Weight is a measurement of the gravitational force acting on an object. Weight depends on the location of an object. An object weighing 1.0 lb on earth weighs about 0.17 lb on the moon. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

10. Properties of Matter Physical Properties: These can be observed/measured without changing the composition of matter – Appearance – Texture – Color – Odor – Shape – Melting Point (MP) – Boiling Point (BP) – Density (D) – Solubility Physical Changes of Matter: – Smashing brick with a hammer – Freezing water or melting ice

11. Properties of Matter Chemical Properties: These can be observed when you try to change matter from one form to another, i.e., you have different arrangements of atoms or molecules – Reactivity – Flammability – Combustibility Chemical Changes of Matter: – Setting a piece of paper on fire – Reaction between an acid and a base

12. Particulate Model of Matter All matter is made up of tiny particles called molecules and atoms. MOLECULES A molecule is the smallest particle of a pure substance that is capable of a stable independent existence. ATOMS Atoms are the particles that make up molecules. Fundamental Particles?Fundamental Particles? Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

13. Molecule Classifications 1.By number of atoms:  Diatomic molecules contain two atoms.  Triatomic molecules contain three atoms.  Polyatomic molecules contain more than three atoms. 2.By type of atoms:  Homoatomic molecules contain same kind of atoms  Heteroatomic molecules contain 2 or more kinds of atoms Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

14. Molecule Classifications (continued) 3.By chemical and physical properties: Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Constant composition Fixed physical and chemical properties Formed by chemical interactions of atoms Composition can vary Physical and chemical properties can vary Compounds retain their chemical identities and can be separated Mixtures Pure Substances

15. Molecule Classifications (continued) 3.By chemical and physical properties: Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Matter MixturesPure Substances HeterogeneousHomogeneous (Solutions) ElementsCompounds More than 1 kind of atom (heteroatomic) Can be divided into simpler compounds or elements (complex sugar  simple sugar  atoms) H 2 O, CO 2, CO, NaCl Only 1 kind of atom (homoatomic) Simplest pure substances Can be divided but chemically same O 2, H 2, Au Properties of a sample vary by location Oil and water Properties of a sample same throughout Sugar and water

16. Measurements and Units Scientific measurements consist of a number and a standard metric unit. Measurements are made using measuring devices (e.g. rulers, balances, graduated cylinders, etc.). Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

17. Measurements and Units Scientific measurements consist of a number and a standard metric unit.

18. Matter MixturesPure Substances HeterogeneousHomogeneous (Solutions) ElementsCompounds Oil and Water Sugar Water More than 1 kind of atom (heteroatomic) Can be divided into simpler compounds or elements (complex sugar  simple sugar  atoms) H 2 O, CO 2, CO, NaCl Only 1 kind of atom (homoatomic) Simplest pure substances Can be divided but chemically same O 2, H 2, Au Constant composition Fixed physical and chemical properties Formed by chemical interactions of atoms Composition can vary Physical and chemical properties can vary Compounds retain their chemical identities and can be separated MixturesPure Substances

19. Temperature Scales The three most commonly-used temperature scales are the Fahrenheit, Celsius and Kelvin scales. The Celsius and Kelvin scales are used in scientific work. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

20. Temperature Conversions Readings on one temperature scale can be converted to the other scales by using mathematical equations. Converting Fahrenheit to Celsius. Converting Celsius to Fahrenheit. Converting Kelvin to Celsius. Converting Celsius to Kelvin. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

21. Scientific Notation Scientific notation provides a convenient way to express very large or very small numbers. Numbers written in scientific notation consist of a product of two parts in the form M x 10 n, where M is a number between 1 and 10 (but not equal to 10) and n is a positive or negative whole number. The number M is written with the decimal in the standard position.

22. Scientific Notation (continued) STANDARD DECIMAL POSITION The standard position for a decimal is to the right of the first nonzero digit in the number M. SIGNIFICANCE OF THE EXPONENT n A positive n value indicates the number of places to the right of the standard position that the original decimal position is located. A negative n value indicates the number of places to the left of the standard position that the original decimal position is located.

23. Scientific ↔ Standard Notation Converting from scientific notation to standard numbers 1.1 x 10 2 = 1.1 x 10 x 10 = 1.1 x 100 = 110Decimal  1.1 x 10 -2 = 1.1/ (10 x 10) = 1.1/100 = 0.011  Decimal Converting Exponents When you move the decimal (l or r), the exponent will be equal to number of places you moved the decimal.

24. Standard to Scientific Notation 60023.5  345.233  -345.233  0.00345  0.10345  1.42  6.00235 × 10 4 3.45233 × 10 2 -3.45233 × 10 2 3.45 × 10 -3 1.0345 × 10 -1 1.42 × 10 0

25. Scientific to Standard Notation 7.932 × 10 5  4.01 × 10 4  3.220 × 10 2  5.673 X 10 1  3.142 X 10 -3  7.6 X 10 -6  4.5655 X 10 -4  793200 40100 322.0 56.73 0.003142 0.0000076 0.00045655

26. Math Operations with Scientific Notation Multiplication Division Addition/Subtraction Convert numbers to the same exponents (5.00 x 10 2 ) + (6.01 x 10 3 ) = (0.500 x 10 3 ) + (6.01 x 10 3 ) = (5.00 x 10 2 ) + (60.10 x 10 2 ) = (5.00 + 60.10) x 10 2 = (65.10 x 10 2 ) = 6510 (6.01 x 10 3 ) - (5.00 x 10 2 ) = (6.01 x 10 3 ) - (0.500 x 10 3 ) = (60.10 x 10 2 ) - (5.00 x 10 2 ) = (60.10 - 5.00) x 10 2 = (55.10 x 10 2 ) = 5510

27. Examples of Math Operations Multiplication a. (8.2 X 10 -3 )(1.1 X 10 -2 ) = (8.2 X 1.1)(10 (-3+(-2)) ) = 9.02 X 10 -5 Not using SigFigs b. (2.7 X 10 2 )(5.1 X 10 4 ) = (2.7 X 5.1)(10 2+4 ) = 13.77 X 10 6 Now change to Scientific Notation 1.377 X 10 7 Division a. 3.1 X 10 -3 = (3.1/1.2)(10 -3-2 ) = 2.6 X 10 -5 1.2 X 10 2 b. 7.9 X 10 4 = (7.9/3.6)(10 4-2 ) = 2.2 X 10 2 3.6 X 10 2 Adding/Subtracting 3.05 X 10 3 + 2.95 X 10 3 = (3.05 + 2.95)(10 3 ) = 6.0 X 10 3

28. Uncertainty and Error Whenever you measure anything, there is always some level of uncertainty. Systematic Error may occur when the instrument is consistently wrong. Random Error may occur due to inaccurate or inconsistent sample measuring or instrument reading. Ex. You step on a scale and the needle is approximately half-way between 150 and 151 lbs, but there are several reasons why you might not be able to correctly identify your weight. The needle moves around a little bit as you watch, so you can’t tell if it is exactly half-way (Random error). The scale is not calibrated properly, so your actual weight is closer to 140 than 150 (Systematic error).

29. Significant Figures Significant figures are the numbers in a measurement that represent the certainty of the measurement, plus one number representing an estimate. Q: When is a number NOT significant? A: Look at the zeros Leading zeros are NOT significant.0.00123 Confined zeros ARE significant.0.00103 Trailing zeros ARE significant, when decimal visible0.0012300 But NOT significant if no decimal12300 Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

30. Calculations with Significant Figures Product (multiplication) or quotient (division) must have same number of sig figs as value with the fewest number of sig figs. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Ex. 01 Ex. 02

31. Calculations with Significant Figures Sum (addition) or difference (subtraction) must contain the same number of places to the right of the decimal (prd) as the quantity in the calculation with the fewest number of places to the right of the decimal. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011 Ex. 01 Ex. 02

32. Calculations with Significant Figures Mixed or long Calculations Your final answer cannot have more certainty than your least certain measurement Calculate your answer first from raw data, then adjust for sig figs. (5.00 / 1.235) + 3.000 + (6.35 / 4.0)=4.04858... + 3.000 + 1.5875=8.630829...

33. Calculations with Significant Figures Rules for Rounding If the first nonsignificant figure to drop from your answer is ≥ 5, all nonsignificant figures dropped, last significant figure increased by 1. If the first nonsignificant figure to drop from your answer is < 5, all nonsignificant figures dropped, last significant figure stays the same. Exact Numbers Numbers with no uncertainty or are known values. Not used to determine sig figs N a = 1 mol = 6.02 x 10 23 π= 3.142 1m = 1000 mm I count 99 bottles of beer on the wall Reduced simple fractions KE = ½ mν 2

34. Dimensional Analysis (Factor-unit method) The factors used in the factor-unit method are fractions derived from fixed relationships between quantities An example of a definition that provides factors is the relationship between meters and centimeters: 1m = 100cm. This relationship yields two factors: Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

35. Factor Unit Method Examples A length of rope is measured to be 1834 cm. How many meters is this? Solution: Write down known quantity (1834 cm). Set known quantity = units of the unknown quantity (meters). Use factor (100 cm = 1 m), to cancel units of known quantity (cm) and generate units of the unknown quantity (m). Do the math. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

36. Factor Unit Method Examples Q: If an arrow shot from a bow travels 30 yards in 1 second, many cm does it travel in 4 seconds? Time = 4 s Rate = 30 yards/sec 1 yard = 3 feet 1 foot = 12 in. 1 in = 2.54 cm Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

37. Percentage The word percentage means per one hundred. It is the number of items in a group of 100 such items. PERCENTAGE CALCULATIONS Percentages are calculated using the equation: In this equation, part represents the number of specific items included in the total number of items. Seager SL, Slabaugh MR, Chemistry for Today: General, Organic and Biochemistry, 7 th Edition, 2011

38. Percentage Calculation A student counts the money she has left until pay day and finds she has $36.48. Before payday, she has to pay an outstanding bill of $15.67. What percentage of her money must be used to pay the bill? Solution: Her total amount of money is $36.48, and the part is what she has to pay or $15.67. The percentage of her total is calculated as follows:

39. Density and Specific Gravity Density is the ratio of the mass of a sample of matter divided by the volume of the same sample. Specific gravity is the ratio of the density of a compound relative to the density of water (D HOH = 1.0 g/cm 3 )

40. Density Calculation A 20.00 mL sample of liquid is put into an empty beaker that had a mass of 31.447 g. The beaker and contained liquid were weighed and had a mass of 55.891 g. Calculate the density of the liquid in g/mL. Solution: The mass of the liquid is the difference between the mass of the beaker with contained liquid, and the mass of the empty beaker or 55.891g -31.447 g = 24.444 g. The density of the liquid is calculated as follows: