1 Introduction and Quick Review of Chem20 Mrs.LeinweberChemistry 20 Review Introduction and Quick Review of Chem20 Mrs.Leinweber

2 Welcome to My ClassroomThe goal in this room is …. SUCCESS !! As a team (student, parent, teacher) we will all work towards the achievement of this goal ! The expectations set out are there to help ensure that students can realistically reach this goal If you have any questions or concerns, make an appointment with me and we can discuss 

3 Do you have the skills ? Let’s go over a few things that you need to make sure you have before you start this course

4 Scientific Knowledge Scientific Knowledge is built on theresults of countless experiments conducted by many investigators. In order to find answers to questions in science, the scientific process must be planned, refined and repeated. A useful, relevant scientific investigation generally has all of the following parts Introduction, Problem, Design, Data collection, Analysis and Conclusion. ** We will be discuss designing and reporting scientific investigations more later on **

5 Scientific Knowledge Science deals with two types of knowledge:Empirical Knowledge – this type of knowledge is OBSERVABLE. (facts) Theoretical Knowledge – this type of knowledge is NON-OBSERVABLE. (interpretations and theories) Empirical (directly observable ) knowledge is considered more certain than interpretations or theoretical explanations.

6 Scientific Knowledge Observations Versus Interpretations in ScienceAn observation is any report from your 5 senses. It does not involve an explanation. An observation can also involve measurements. Such an observation is a quantitative one, as opposed to a qualitative one (no measurements). An interpretation is an attempt to figure out what has been observed. Here are some examples designed to help you distinguish between them.

7 Scientific Knowledge

8 Scientific Knowledge When doing labs, it is important not to confuse observations with interpretations. If you are asked to observe, you should not identify gases. “Hydrogen gas was produced when zinc was added to acid” is not an observation. What you see is the liquid getting cloudy with bubbles rising from the area surrounding the metal. The latter is what you should be reporting as an observation. Later if you collect the gas and it has hydrogen’s characteristic property, then it will be time to conclude that the bubbles contained hydrogen gas.

9 Precision, Error and AccuracyA major component of the scientific inquiry process is the comparison of experimental results with predicted or accepted theoretical values. All measurements have some degree of uncertainty. Uncertainty, often referred to as error , is not the result of a mistake necessarily, but can be caused by limitations of the equipment or experimenter.

10 Precision, Error and AccuracyPrecision – describes the exactness and repeatability of a value or set of values. Accuracy – describes the degree to which the result of an experiment or calculation approximates the true value.

11 Percent Difference & Percent YieldPercent difference – determines how precise a set of measurements is. Compares experimental values to constants or theoretical values. (acceptable values for high school labs are ±10%.) Percent yield – determines how close an actual(experimental) yield is to the predicted or theoretical value. (acceptable values for high school labs are >90%.)

12 Significant Digits

13 Significant Digits - ExamplesHow many significant digits does each number have ?? 4568 3.49 50.17

14 Significant Digits

15 Significant Digits

16 Significant Digits – Calculations ExamplesHow many significant digits would each answer have ?? Don’t forget to round properly  – = 4.88 / 2.3 x 1.16 =

17 Scientific Notation Used to simplify large numbers (and to clarify the number of significant digits), when reporting them and doing calculations.

18 Scientific Notation MAKE SURE YOU CAN PUT THIS INTO YOUR CALCULATOR ACCURATELY!!!

19 Rules for Scientific Notation

20 Rules for Scientific NotationIf you want to make an exponent larger move the decimal to the left If you want to make an exponent smaller move the decimal to the right

21 Scientific Notation Examples Complete the following 6.7 x x = 4.6 x x = (8.9 x 10 2) x (5.2 x 10 4 ) = (4.4 x 10 8) / (1.2 x 10 3) =

22 SI Units The International System of Units, SI units, from the French Name, systéme international d’unites. The measurement and communication system used internationally by scientists Physical quantities are ultimately expressed in terms of 7 fundamental SI units, called base units. All other quantities are derived from these units

23 SI Units – Base units Universality

24 SI Units – Prefixes ConvenienceWith the exception of centi-, we use prefixes that change by multiples of 1000

25 Tables and Graphs Both tables and graphs are used to summarize information and to illustrate patterns or relationships. It is important to prepare them accurately, following accepted conventions, in order to best communicate the information.

26 Tables All tables must have descriptive titles that mention both the manipulated and responding variables. The row or column containing the manipulated variable precedes the responding Headings of rows or columns are labelled with the units in parentheses where applicable and not included in the body of the table.

27 Graphs Descriptive title (manipulated & responding variables are mentioned) Labeled axes with units in parentheses X-axis has the manipulated variable Y-axis has the responding Use appropriate scales, with equal increments, on the axes so the graph takes up at least ½ of the paper Use a pen for data points Draw, in pencil, a line of best fit if the data appears in to be in a straight line, if not use a smooth curve to connect the points

28 Graphs Interpolation ExtrapolationFinding a value between a set of data points Extrapolation Finding a value past a set of data points (dotted line)

29 Manipulating FormulasOpposite Operation What is done to one side must be done to the other.

30 Now down to the Chemistry  Let’s get ready for Unit 1Remember … Chemistry is the physical science that deals with the composition, properties and changes in matter.

31 Classifying Matter Matter: anything that has mass and takes up space. Chemists can classify matter as solid, liquid, or gas. But there are other ways to classify matter, as well — such as pure substances and mixtures. Classification is one of the basic processes in science. All matter can be classified as either a pure substance or a mixture

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33 PURE SUBSTANCES: A pure substance has a definite and constant composition — like water or sulfur. A pure substance can be either an element or a compound, but the composition of a pure substance doesn’t vary. Pure substances have definite properties that are always the same (e.g. mp, bp, density, etc). You cannot distinguish a solution from a pure substance by looking at it, as they appear homogenous. There are two types of pure substances:

34 Elements: are pure substances that cannot be broken down into simpler substances. They are the smallest particle that is the element. Empirical They are composed of one type of atom (Theoretical), and are the building blocks of all compounds. Each element has its own symbol, the first letter always being capitalized, and the second letter being lower case. Currently there are 118 different elements *** Every element has very unique properties (i.e. (s), (l), (g), reactivity, density etc.).

35 Compounds: are pure substances that contain two or more elements in definite proportions (fixed ratios) Theoretical Compounds can be broken down into the elements they are made of. Empirical Even though there are only 118 elements, there are millions of compound possibilities that can be formed from these elements.

36 Mixtures: Mixtures are physical combinations of pure substances that have no definite or constant composition — the composition of a mixture varies according to who prepares the mixture. Although chemists have a difficult time separating compounds into their specific elements, the different parts of a mixture can be easily separated by physical means, such as filtration or evaporation

37 Heterogeneous Mixtures:aka Mechanical Mixtures have visible different parts (e.g. concrete, salt & pepper, soil, granola, chocolate chip cookies, etc.).

38 Homogeneous Mixtures aka solutions are type of mixture where you only see one part. It appears to be one substance (e.g. soft drinks, clear tea, Kool-Aide, salt water, metal alloys, etc.) and cannot be easily distinguished from a pure substance without further testing.

39 Periodic Table Dmitri Mendeleev created the periodic table in 1869Periodic Law Chemical & physical properties of elements repeat themselves in regular intervals when elements are arranged in order of increase atomic #

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41 How its put together

42 Family (or Group): are the vertical columns on the periodic table numbered Elements within a group have similar chemical properties to each other.

43 Periods: are the horizontal rows on the periodic table numbered 1-7Periods: are the horizontal rows on the periodic table numbered 1-7. As you move left to right, you see an increase in the number of protons. On the far left we start with metals, and as we move farther to the right we end up in non-metals. Reactivity in metals usually decreases as we move to the right.

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45 Atoms: Atoms are the basic units of matter and the definingstructure of elements. Atoms are made up of three particles: protons, neutrons and electrons. Electrons are extremely lightweight and exist in a cloud orbiting the nucleus. The electron cloud has a radius 10,000 times greater than the nucleus

46 Atomic number: represents the number of protons an atom contains. The atomic number is found in the top left corner of each element.

47 Atomic Mass: is the mass of each element. This mass is made up of the number of protons plus the number of neutrons found in an atom (remember electrons have very little mass).

48 Isotopes: Atoms of the same element that have a different number of neutrons.

49 Isotopes:

50 IONS Are atoms that have either lost or gained electrons in order to have a full outer shell of electrons. Metals will lose electrons to become positively charged ions(CATIONS). Non-metals will gain electrons to become negatively charged ions (ANIONS). *** It is important to remember that electrons are negatively charged, so if an atoms gain electrons the charge becomes negative.

51 IONS

52 Bohr Model/Energy Level DiagramWrite # of protons and neutrons in the nucleus (or under a line) Draw circles (or lines) to represent electron shells Draw on Electrons Remember: 2 – 8 – 8- 16 If there are more than 4 electrons there will be pairs

53 Bohr Model/Energy Level Diagram Examples

54 Always in pairs. KNOW THESE!!!!Diatomic Elements Always in pairs. KNOW THESE!!!! H2, N2, O2, F2, Cl2, Br2, I2

55 Polyatomic Elements S8 and P4 Know these too! *This is SUPER important for writing and balancing chemical reaction equations*

56 Compounds Ionic & Molecular Compound properties Naming Ionic compoundsNaming Molecular compounds

57 Ionic Compounds Ionic compounds are puresubstances formed from a metal and a non-metal. Ex. NaCl All ionic compounds are solids at room temperature (High mp. And bp.) When an ionic compound is dissolved in water, it will conduct electricity (because of the presence of charged particles) Ionic Bond formed by the transfer of electrons creating ions that are attracted to each other Example: Na + and Cl NaCl

58 Molecular Compounds Molecular compound- when two non-metals combine and electrons are shared They can be solids, liquids or gases at room temperature They tend to be insulators (poor conductors of electricity) even in solutions, because there are no IONS in MC  Have low melting and boiling points

59 Names and Formulas The chemical formula of an elementor compound indicates which elements and how much of each one is present in a compound. Chemical nomenclature is the systematic method of naming substances Ex. table salt has the chemical name sodium chloride and its formula is NaCl.

60 International Union of Pure and Applied ChemistryIUPAC for short  The governing body for scientific communication which specifies rules for chemical names and symbols (**Remember the universality of science ) These are the rules we follow for naming all elements and compounds.

61 Naming Ionic CompoundsWhen you have a chemical formula LiF(s) MgCl2 (s) FeS(s) Fe2S3 (s) Identify the metal first, and name Check the periodic table to see if the metal is multivalent (more than one ion charge) If not move on, if it is use subscripts to determine which charge and use roman numerals between the names Name the non-metal second The name of the non-metal changes its ending to –ide **In high school chemistry, we name BINARY COMPOUNDS, so there is only one cation and one anion**

62 Naming Ionic CompoundsWhen you have a chemical formula LiF(s) Lithium fluoride MgCl2 (s) Magnesium chloride FeS(s) Iron (II) sulfide Fe2S3 (s) Iron (III) sulfide

63 Polyatomic Ions Polyatomic ions are groups of atoms acting as one. Ex. Carbon and oxygen can act as one - CO32- Polyatomic ions are composed of two or more elements covalently bonded with an overall negative or positive charge. These are found on your periodic table.

64 Polyatomic Ions This table is found on the back of your periodic table, and is used as a reference for you to help identify them in the compounds.

65 Polyatomic Ions - ExamplesNaCH3COO (s) – sodium acetate K2Cr2O7 (s) – potassium dichromate MnHPO4 (s) – manganese (II) hydrogen phosphate

66 Steps to writing formulasWrite the metal ION symbol with its ion charge. Next to it write the non-metal elements symbol with its ion charge. Ex. Ca2+ Cl- Balance the ION charges to ZERO. There must be two chlorine atoms each with an ion charge of 1- to balance the 2+ ion charge of one calcium atom. Ex. Cl- Cl Ca2+

67 Steps to writing formulasWrite down the formula by indicating how many atoms of each element are in it. Then write the state in parentheses. Ex. CaCl2 (s) After a chemical formula a subscript in parentheses indicates the compounds physical state. Ex. (g) = gas, (l) = liquid, (s) = solid (aq) = aqueous (dissolved in water)

68 Now you Try Beryllium nitrate Copper(I)sulfate Be(NO3)2 (s) Cu2SO4 (s) Sodium hydroxide Chromium(II)phosphate NaOH (s) Cr3(PO4)2 (s)

69 Ionic Hydrates Some ionic compounds exist as hydratesThis means that the compounds have water molecules attached to them.

70 Ionic Hydrates To name them Identify and name the ionic compoundThen use a prefix for the number of water molecules attached and put it in front of the word hydrate. Example BaSO4 • 4H2O (s) Barium sulfate tetrahydrate

71 Naming Molecular CompoundsThe first element in the compound uses the element name The second element has the suffix -ide When there is more than one atom in the formula, a prefix is used which specifies the number of atoms However, when the first element has only one atom the prefix mono is not used. ***Many molecular compounds are known by their common names Ex. water, ammonia

72 Prefix Number of atoms Mono 1 Di 2 Tri 3 Tetra 4 Penta 5 Hexa 6 Hepta 7 Octa 8 Nona 9 Deca 10

73 Example CO2 = carbon dioxide N2O = dinitrogen monoxide CCl4 = carbon tetrachloride

74 Writing Formulas for Molecular CompoundsThey differ from ionic compounds because no ions are present Remember!! A few elements are diatomic meaning that they are always found in the form O2, N2, H2, F2, Cl2, Br2, I2, and P4, S8 when written alone

75 Writing Formulas for Molecular CompoundsP5O10 = pentaphosphorous decaoxide CO = carbon monoxide CH4 = carbon tetrahydride (METHANE  )

76 Naming Acids and Bases Acids – are aqueous hydrogen compounds that make blue litmus paper turn red and are generally written with hydrogen appearing first in the formula Ex. HCl Ex. H2SO4 **** acids containing –COOH, the H is placed at the end of the formula Ex. CH3COOH(s) acetic acid Ex. C6H5COOH(s) benzoic acid

77 1. If the anion name ends in “ide” the acid name is hydro__________ic acidEg. HCl hydrochloric acid Eg. H2S hydrosulfuric acid 2. If the anion name ends in “ate” the acid name is ________ic acid Eg. HNO3(aq) nitric acid Eg. H2SO sulfuric acid 3. If the anion name ends in “ite” the acid name is ________ous acid Eg. H2SO nitrous acid

78 Bases Strong Bases are aqueous ionic hydroxides form electrically conductive solutions and turn red litmus paper blue. Eg. KOH – potassium hydroxide Weak bases – react with water to form solutions containing hydroxide

79 Practice Problems TEXTBOOK Pg 40 Q 20-22, 24-25

80 Types of Chemical ReactionsFormation Decomposition Combustion Single Replacement Double Replacement

81 Formation Reactions (synthesis)2 elements combine to form a compound X + Y = XY Ex Na(s) + Cl2(g)  NaCl(s)

82 Decomposition Reactionsone compound breaks down into two or more simpler compounds or elements (opposite of formation reactions) XY  X + Y Ex. 2 H2O(l)  2H2(g) + O2(g) (electrolysis of H2O)

83 Hydrocarbon CombustionA hydrocarbon (CnHn) (i.e. oil, fuel, natural gas) reacts with oxygen gas to form carbon dioxide gas and water vapor and heat. CnHx + O2(g) = CO2(g) + H2O(g) Ex. CH4(g) + O2(g)  CO2(g) + H2O(g) + thermal energy

84 Single Replacement Reactions Zn(s) + CuCl2 (aq) ZnCl2(aq) + Cu(s)Single replacement reactions replace one element from a compound with another element. A compound and an element react, and the element switches places with part of the original compound. AB + C  AC B Zn(s) + CuCl2 (aq) ZnCl2(aq) + Cu(s)

85 Double Replacement ReactionsDouble replacement reactions swap elements between 2 compounds that react together to form two new compounds. Two compounds react, with elements switching places between the original compounds. See page 262

86 Types: Double replacementExample: MgO + CaS S O  Mg Ca + O S Mg Ca +

87 Summary

88 Solubility A SOLUTION is a homogeneous mixture of a SOLUTE (substance dissolved) and a SOLVENT (substance dissolving, usually a liquid). SOLUBILITY is the maximum amount of a substance that can be dissolved at a specific temperature.

89 Solubility

90 Balancing Chemical Reaction EquationsA balanced chemical reaction is one in which the total number of each kind of atom or ion in the reactants is equal to the total number of the same kind of atom or ion in the products. This is the LAW OF CONSERVATION OF MASS

91 Balancing Equations Need equal # of each element on both sides of the reaction

92 Steps for balancing chemical EquationsStep #1 – write the chemical formula for each reactant and product including the state of matter Step #2- balance the atom or ion using coefficients. Make sure to use the lowest common multiple Step #3 Check your balancing

93 Chemical Amount A mole is the unit of chemical amount of a substance with the number of entities corresponding to Avogadro’s number (6.02 x 1023) N2(g) + 3H2(g) → 2NH3(g) One mole of nitrogen gas and three moles of hydrogen gas react to form two moles of ammonia gas

94 Molar mass (M)- the mass of one mole of the substance. Units are g/molHow to calculate molar mass: 1. Write the correct chemical formula 2. Identify the number of atoms of each element present in the formula 3. The molar mass is the sum of the number of atoms multiplied by the atomic molar mass of each element.

95 Calculations: 3 types of problems that you have to be able to do:calculate M (molar mass) – you can do that now! (Remember: Periodic Table) calculate the number of moles of a substance (given the mass in g) calculate the mass of a substance (given the number of moles) The Mole

96 Practice Problems Textbook Pg 66 Q 7-9 Pg 69 Q 18-27

97 pH, pOH and [H3O+] &[OH-] pH = -log [H3O+(aq)] [H3O+(aq)] =10 –pHpOH = -log [OH -(aq)] [OH -(aq)] =10 –pOH The number of digits following the decimal point in a pH or pOH value is equal to the number of significant digits in the corresponding hydronium or hydroxide concentration. For both pH and pOH, an inverse relationship exist between the ion concentration and the pH or pOH. The greater the hydronium ion concentration, the lower the pH is.

98 (Complete ionic equation)Net Ionic Equations Write the net ionic equation for the reaction of aqueous barium chloride and aqueous sodium sulfate. (Refer to the solubility table) 1) BaCl2(aq) + Na2SO4(aq)  BaSO4(s) + 2NaCl(aq) 2) Ba2+(aq) + 2Cl-(aq) +2Na+(aq) + SO42-(aq)  BaSO4(s) + 2Na+(aq) + 2Cl-(aq) (Complete ionic equation) 3) Ba2+(aq) + 2Cl-(aq) +2Na+(aq) + SO42-(aq)  BaSO4(s) + 2Na+(aq) + 2Cl-(aq) 4) Ba2+(aq)) + SO42-(aq)  BaSO4(s) (Net ionic equation) Ions that are present but do not take part in (change during) a reaction are called spectator ions (like spectators at a sports game: they are present but do not take part in the game)

99 Stoichiometry (Measured quantity) solids/liquids m  n(Required quantity) mole ratio

100 Limiting and Excess ReagentsWhen no further changes appear to be occurring, we conclude that the reaction has gone to completion. A limiting reagent is the reactant whose entities are completely consumed in a reaction, meaning the reaction stops. In order to make sure this happens, more of the other reactant must be present than is required An excess reagent is the reactant whose entities are present in surplus amounts, so that some remain after the reaction ends..

101 Practice How much precipitate is produced from a reaction of 10.0g of Silver nitrate and 10.0g of zinc chloride ?

102 Practice In an experiment, a mL sample of sulfuric acid solution reacts completely with 15.9 mL of mol/L potassium hydroxide. Calculate the amount concentration of the sulfuric acid.

103 Practice Problems Textbook Pg 66 Q 7-9 Pg 69 Q 18-27