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CHEMISTRY Mrs. Adams Room 601.

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Presentation on theme: "CHEMISTRY Mrs. Adams Room 601."— Presentation transcript:

1 CHEMISTRY Mrs. Adams Room 601

2 Today’s TO DO: 9/7/10 Introduction Agenda Book Syllabus & Expectations
Notebooks Hand out Set up Journal #1 Homework 

3 Notebook My gift to you…I expect you to BRING it and USE it E-V-E-R-Y DAY! Set Up Leave 1st 4 sheets blank for Table of Contents Date Title Format pg # What is Chemistry? Journal 1 if same as above, leave blank Notes Starting with 5th sheet as pg.1, number each page in lower outer corner

4 Journal #1 – What is Chemistry?
Go to pg.1 of your notebook Write title Journal #1 – What is Chemistry? on top w/date on right, Write the following prompt: What do you think of when you hear the word chemistry? List 10 things that you think are part of chemistry. Leave a blank line Answer

5 Homework: Real Life Application of Chemistry
Find a current article showing an application of chemistry to our lives Write a summary in your own words that includes main ideas of the article your thoughts about the article 1-2 questions you still have that were NOT answered in the article If done on computer, me the article as well as your summary (GREEN ) If hand written, also provide a printout of the article DUE: Friday, Sept. 10th

6 Journal #2 – Self Reflection
How do you think your friends and family would describe you? What are your goals & expectations in this class? What do you hope to accomplish in the next 5 years? What do you hope to accomplish in the next 15 years?

7 Lab Safety Read the Flinn Lab Safety Directions
To reinforce the concepts in your assigned section, your table will create one of the following to present to the class: Skit Poem Cartoon Song Poster or Ad Your presentation should try to convince us that your rules are THE MOST IMPORTANT rules to follow in this class.

8 Lab To do a lab really well, have your report done well in advance.
In Other Words… Know WHAT you’re supposed to do BEFORE you do it. Know WHY you’re doing something BEFORE you do it.

9 Lab BEFORE every lab, you will address: Background information
Materials & procedures Safety concerns Hypothesis Creation of data tables

10 Lab AFTER every lab you will Organize your data into charts and graphs
Analyze your data Accept or reject your hypothesis Discuss your conclusion & possible source(s) of error

11 Scientific Method A way to answer questions about the world based on observations and experiments. INQUIRE Ask questions OBSERVE Not always visual EXPERIMENT Changing a variable to determine UNDERSTAND Where does your Hypothesis fit? Where does your Conclusion fit?

12 Scientific Method in Action
CHEMOTHERAPY OIL SPILLS FOOD INDUSTRY COSMETIC INDUSTRY For each: Explain the general idea of why scientific method applies Write a SPECIFIC observation or question AND hypothesis you could test

13 Measurements Always uncertain Instruments never flawless
Some estimation always required Example: A ruler

14 Numbered lines = centimeters Smaller lines = 0.1 cm = 1 mm
Reliable Smaller lines = 0.1 cm = 1 mm Any point between each line Must be estimated Not reliable Ex: Line above the ruler = _____ cm Written as _______________ cm 2.45 2.45  +/- 0.01

15 Reliability 2 ways to check numbers Precision Accuracy
repeat measurement test against standard Precision how close repeated measurements are to each other Accuracy how close measurements are to standard or accepted value

16 Precision vs. Accuracy

17 Sample Problem In 3 separate trials, Sara calculates the density of water to be 0.88g/ml 0.87g/ml Is she precise ? Yes, all close together Is she accurate? No, accepted value is 1 g/ml Density = mass/volume

18 Significant Digits Sig Figs
Number of digits within a value that are considered significant with respect measurement validity Follows Pacific-Atlantic Rule

19 Decimal PRESENT - PACIFIC
Start at far left of number (like pacific ocean is on far left of US)  Start counting first Non Zero number End at rightmost digit (including zeros)

20 Examples 34.067g 0.0007458ml 0.009070g ____ sig figs ____ sig figs

21 Decimal ABSENT – ATLANTIC
far right of number (like the Atlantic Ocean is at far right of US) Start counting first Non Zero number End at leftmost digit

22 Examples 2030cm ____ sig figs 2007dm 19,000,000,000g 3 4 2

23 Practice Problems 0.0026701m 19.0550kg 3500V 1,809,000L ____ sig figs
4

24 Sig Fig's in Calculations
Exact numbers or conversions do not count as sig figs Ex: Speed of light ~ 300,000,000 m/s Can have infinite # of sig figs and must be specified Sample – Speed of light expressed to 3 significant digits = 3.00 x 108

25 Multiplication & Division
Answer must have same # of sig figs as lowest sig figs found within problem Ex: x = = 58.0 (3) (5) (3)

26 Example: Volume = length x width x height
Find the volume an object with sides m x 1.34m x 13.22m (5 sig figs) (3) (4) m (10) m3 (5) 192.7 m (4) 193 m (3)

27 Addition & Subtraction
Largest uncertainty determines number of sig figs Answer will have lowest sig figs to the right of the decimal from numbers in problem Ex: =

28 More Examples 34.50g + 3.2345g + 671.1g + 25.345g 2092 ml – 147.54 ml

29 Practice Problems 6.15m x 4.026m = 12.7km / 3.0 =
150ml ml + 209ml ml = (35.6L + 2.4L) / 4.803  = 2.542m x (16.408m m) = 24.8 m2 4.2 km ~ 440 ml 7.91 L 31.85 m

30 Journal # 3 – Sig Figs What are the rules for significant figures?
Be sure to include those for addition/subtraction and multiplication/division

31 Scientific Notation One digit to the left of the decimal
# of digits to right of decimal is determined by sig fig rules Example 19,000,000 ml 2 sig figs 1.9 x 107 ml Example g 4 sig figs 4.569 x 10-4 g 

32 Scientific Notation Practice
3.27 x 104   32,700 = 1,024,000 = = = 1.024 x 106    x 10-3    3.901 x 10-9   

33 Knowing Equations: Density
Density = mass /volume D=m/v Know formula & manipulate w/algebra OR know graphic below M D V

34 Dimensional Analysis Step by step conversion between units
Convert 10.0µm to inches Conversion factors  1m=1,000,000µm   1m = 39.37inches Start with the given unit, then use you conversion factors to cancel units until to arrive at the unit you want to convert to. 10um x  1m   x  39.37inches = in       1,000,000um   1m   

35 Practice Problems 250.0 cm to inches ? gal in 39L ? cm in 16in
? seconds in 5 days ? ft in 86cm ? cm3 in 2.3gal ? m in 3.5mi

36 Percent Error % Error = measured – accepted x 100 accepted Ex:
Accepted value for density of water = 1 g/ml Measured value for density of water in lab = 0.9 g/ml % Error = (0.9 – 1)/1 * 100 = 10% Error

37 Journal #4 – Dimensional Analysis
Are there REALLY seconds in a day? Show your dimensional analysis to defend your answer.

38 Energy & Matter i Chapter 2

39 Energy Potential to do work or produce heat 3 Main Types – Radiant
Ex: sunlight Kinetic Energy of motion Ex: Mechanical – energy of moving parts Ex: Thermal – energy from internal particle motion in matter Potential Ex: Gravitational – falling water Ex: Electrical – opposite charges Ex: Chemical - battery

40 Energy Units calories Joules Measuring calories
Amount of heat needed to raise 1g of water 1oC 1 Calorie in food = 1000 calories Joules SI unit of energy 1 cal = J Measuring calories Calorimeter

41 Law of Conservation Energy is neither created nor destroyed in any process Energy can be transformed from one form to another Ex: kinetic energy of bat transferred to baseball (kinetic, sound) Ex: Chemical energy of striking match transformed into heat and light

42 Temperature Celsius Kelvin 0oC = freezing pt 100oC = boiling pt
21oC = room temp 37oC = body temp Kelvin SI Unit of temperature oC + 273

43 Matter can be neither created nor destroyed
Matter & Conservation Has mass & volume States Solid Liquid Gas Changes Physical Chemical Just like Energy, Matter can be neither created nor destroyed

44 Elements Meet the Elements
Substances that cannot be separated into simpler substances by chemical change Organized in Periodic Table Combine chemically to form COMPOUNDS Meet the Elements

45 Mixtures Blend of 2 or more pure substances (elements or compounds)
Heterogeneous Visible differences in combined substances Ex: chocolate chip cookies Homogeneous No visible differences in substances Ex: salt water

46 Separation of Homogeneous Mixtures
Distillation Tap water = homogeneous mixture Crystallization Evaporation of liquids from solids Chromatography Separates mixtures by Solubility size charge

47 The Atom Chapter 3

48 Early Models of the Atom

49 Atoms The Greek Philosopher Democritus
Proposed all matter made up of small, indivisible particles Called these “atomos” = atoms Today’s Definition- smallest particles of an element that retain properties of element

50 Democritus Ideas REJECTED
Because he didn’t know what held these particles together They remained rejected until the 17th century when better technology = closer observations

51 1700’s Lavoisier Joseph Louis Proust Law of Conservation of matter
Law of constant composition compounds always contains same elements in same proportions by mass

52 John Dalton Atomic Theory of Matter:
Each element is composed of extremely small particles called atoms All atoms of an element are identical, but differ from those of other elements Atoms are neither created or destroyed in a chemical reaction A compound always has the same relative numbers of atoms.

53 Discovering Atomic Structure

54 Michael Faraday structure of atoms is related to electricity
atoms contain particles that have electrical charge.

55 Static Electricity Benjamin Franklin famous electricity experiment
Conclusions from his kite & key experiment lightning is a static discharge from clouds electricity has two kinds of charges. Positive (+) Negative(-)

56 Cathode Rays & Electrons
Running electricity through a partially evacuated glass tube occurs in a cathode tube. Negative end = cathode Positive end = anode With a fluorescent lining & addition of electricity, particles are visible

57 Where's the rest of the mass??
J.J. Thompson Deflected particles w/magnet Discovered particles were negatively charged Named them electrons Along w/Milikin discovered their mass to be only 1/2000 of full atomic mass … Where's the rest of the mass??

58 Ernest Rutherford Discovers alpha particles Discovers beta particles
Deflects towards negative plate Charge = Discovers beta particles Deflect towards positive plate Discovers gamma particles Not affected by electric plates

59 The Nuclear Atom If electrons are negative, why are atoms neutral?
Must contain positive parts equal to the negative parts. Where are they?

60 The Gold Foil Experiment (figure 3-14)

61 What does this mean? The experiment determined that most of the atoms positive charge, as well as the mass, is in the middle, called the nucleus. Most of the particles pass through the empty space but occasionally one gets close enough to the positive nucleus to deflect

62 Modern Atomic Theory

63 Atoms Smallest units of matter
Composed of protons (+) and neutrons in nucleus and electrons (-) in orbitals

64 1 proton has the mass of about 2000 electrons

65 Periodic Table Information
Atomic Number Represents # of protons Also # of electrons in a stable atom of an element Discovered by Moseley Atomic Mass Sum of protons & neutrons Electron mass is small and almost negligible

66 Electrons Electrons move in space around the nucleus
Rutherford visualized it as a mini solar system.

67 Atomic Mass Measured by Atomic Mass Units (AMU)
atomic mass approximately = protons + neutrons Atomic Mass (AM) = average mass of element’s atoms, including isotopes

68 Unstable Atoms Ions Isotopes Different # of electrons
Atoms with a charge Isotopes Different # of neutrons Often radioactive Used as diagnostic tracers

69 Ions When an atom gains or loses electrons it acquires a charge
Fewer electrons means positive charge More electrons means negative charge Charge of ion = # protons - # electrons

70 Sample Write the chemical symbol for the ion with 9 protons and 10 electrons Answer F- What is the symbol of the ion with 13 protons and 10 electrons? Answer Al3+ 7 Protons and 10 electrons? N 3-

71 Isotopes Dalton said all atoms of an element are the same.
Not quite true, ISOTOPES have a different number of neutrons In nature, elements are almost always found as a mixture of isotopes

72 Identifying Isotopes To identify isotopes more specifically
Use the Mass Number Mass Number = (# protons) + (# neutrons)

73 Fundamental Subatomic Particles
Location Charge (C) Mass (g) Mass (AMU) Proton Inside nucleus x 10-19 1.673 x 10-24 1 Neutron 1.675 x 10-24 Electron Outside nucleus x 10-19 9.109 x 10 –28

74 Changes in the Nucleus

75 Radioactive Elements Discovered by Becquerel in late 1800s
Uranium Pierre & Marie Curie Radium & Polonium

76 Nuclear Reactions Changes in nucleus Unstable nucleus  radioactivity
Changes composition of nucleus Alpha & beta radiation comes from nucleus Unstable nucleus  radioactivity Not many elements radioactive Why not?? Seems like all those + protons would cause a lot of repelling…..

77 Composition of Stable Nuclei
As the number of protons increases, it takes more and more neutrons to remain stable. All atoms above 83 are unstable

78 Radioactive Decay RULE:
sums of mass numbers & atomic numbers are same before & after reaction

79 Types of Radioactive Decay
Alpha Radiation Stream of high energy alpha particles Consists of 2 protons and 2 neutrons Identical to a helium-4 nucleus Symbol 42He2+ or 42He or 42 Do not cause a health risk Do not travel far

80 When an atom emits one of these, it is said to be undergoing radioactive decay
Which brings us to the nuclear equation, or a way to keep track of the components

81 Alpha Decay

82

83 stream of high speed electrons
Beta Radiation stream of high speed electrons neutron changes into 1 proton & electron proton stays in nucleus electron is propelled out at high speed Symbol of 0-1e- or 0-1e or 0-1 damaging to skin

84 Beta Decay

85

86 Extremely energetic form of light energy we cannot see
Gamma radiation Extremely energetic form of light energy we cannot see Symbolized by 00 Does not consist of particles Able to penetrate deeply into substances

87

88 Types of Radiation Name Identity Charge Alpha () Helium-4 nuclei 2+
Penetrating Ability Alpha () Helium-4 nuclei 2+ Low Beta () Electron 1- Medium Gamma () High energy particle None high

89 Practice Problems Write the nuclear equation for the alpha decay of uranium 238. Write a nuclear equation for the beta decay of sodium 24

90 Other Nuclear Reactions
Nuclear Fusion Atoms collide and join together releasing great amounts of energy Like in the Sun Nuclear Fission Splitting the nuclei of large atoms Like in Nuclear reactors

91 Bonding Atoms bond in order to fill their valence shell (outer energy level) Octet Rule The idea that most atoms want 8 electrons in their outer shell and will share, steal, or give away electrons in order to fill the valence shell Exceptions are those who have less than 6 total electrons

92 Covalent Bonding Valence electrons (outer shell) are shared
Form molecules Single, double or triple bonds are possible

93 Ionic Bonding Electrons are lost or gained from the outer shell in order to fulfill the octet rule

94 Hydrogen Bonding Weak bonds formed between molecules that contain polar covalent bonds Bonding animation:


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