# A QUICK OVERVIEW OF CHAPTERS 1-3 AP C H E M I S T R Y M S. G R O B S K Y

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1 A QUICK OVERVIEW OF CHAPTERS 1-3 AP C H E M I S T R Y M S. G R O B S K Y

2 CHAPTER 1 S C I E N T I F I C M E T H O D, SI U N I T S, M E T R I C P R E F I X E S S I G F I G S, D I M E N S I O N A L A N A L Y S I S

3 THE SCIENTIFIC PROCESS VIA SCIENTIFIC METHOD Making observations Qualitative Quantitative (measurement) Formulating hypotheses Possible explanation for observations Performing experiments to test predictions Gathering new information to decide whether hypothesis is valid Experiments always produce new observations and thus, the scientific process begins again

4 THE SCIENTIFIC PROCESS IS ONGOING To understand a given phenomenon, the steps of the scientific method are repeated many times Gradually, enough knowledge is gathered to provide a possible explanation of the phenomenon Theory (a model) A human invention that is continuously refined as new knowledge is acquired Defined as a set of tested hypotheses that gives an overall explanation of some natural phenomenon Different from natural law

5 THEORY VERSUS NATURAL LAW Theory Attempts to explain why the observation happens Natural law A summary of what happens Sometimes, the same observation applies to many different systems

6 QUANTITATIVE OBSERVATIONS Recall that quantitative observations are also known as measurements Consist of two parts: A number A unit The SI system of units is the adopted international system for measurements Based on the metric system The Fundamental SI Units Physical Quantity Name of Unit Abv. Mass Kilogram kg Length Meter m Time Second s Temperature Kelvin K Electric current Amount of substance Ampere Mole A mol

7 PREFIXES OF SI SYSTEM Most Commonly Encountered SI Prefixes Prefix Symbol Meaning Exp. Notation Mega M 1,000, Kilo k 1, Hecto h Deka da The fundamental units are not always convenient Thus, prefixes are used to change the size of the unit Deci d Centi c Milli m Micro μ Nano n Pico p

8 UNCERTAINTY IN MEASUREMENTS In a measurement, you record the certain digits and the first uncertain digit (estimation) Called significant figures (sig figs, SF) Uncertainty of a measurement depends on the precision of the measuring device Due to the fact that different measuring devices have different degrees of precision, rules have been developed for counting the sig figs in each number and for determining the correct number of sig figs when doing calculations

9 EASY WAY TO COUNT SIG FIGS ATLANTIC PACIFIC METHOD If the decimal is PRESENT, start counting from the Pacific side of the number Start counting when you get to the first real digit and count all remaining digits If the decimal is ABSENT, start counting from the Atlantic side of the number Start counting when you get to the first real digit and count all remaining digits

10 COUNTING SIG FIGS

11 CALCULATIONS USING SIG FIGS YOU ARE ONLY AS PRECISE AS YOUR LEAST PRECISE MEASUREMENT! Addition/subtraction When adding or subtracting decimals, the answer must have the SAME NUMBER OF DIGITS TO THE RIGHT OF THE DECIMAL POINT as there are in the measurement having the FEWEST DIGITS TO THE RIGHT OF THE DECIMAL POINT

12 CALCULATIONS USING SIG FIGS Multiplication/division The answer can have no more significant figures than are in the measurement WITH THE FEWEST NUMBER OF SIG FIGS!

13 DIMENSIONAL ANALYSIS FACTOR LABEL METHOD It is often necessary to convert a given result from one system of units to another Write the GIVEN number and unit Set up a conversion factor (fraction used to convert one unit to another) Place the given unit as denominator of conversion factor Place desired unit as numerator Place a 1 in front of the LARGER unit Determine the number of smaller units needed to make 1 of the larger unit Cancel units and solve the problem!

14 DENSITY CALCULATIONS Density is the mass of substance per unit volume of the substance D = m V Units are g/ml (liquid) or g/cm 3 (solid) Know how to do density by volume displacement and how to use density as a conversion factor!

15 CHAPTER 2 H I S T O R Y O F A T O M I C M O D E L, F U N D A M E N T A L C H E M I C A L L A WS, A T O M I C S T R U C T U R E, N O M E N C L A T U R E

16 FUNDAMENTAL CHEMICAL LAWS Law of Conservation of Mass Mass is neither created nor destroyed Verified by Antoine Lavoisier Law of Definite Proportions A given compound always contains exactly the same proportion of elements by mass Discovered by Joseph Proust Law of Multiple Proportions When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers Discovered by John Dalton

17 LAW OF MULTIPLE PROPORTIONS

18 ATOMIC THEORY DEVELOPMENT Democritus (400 B.C) First used the term atom - atomos John Dalton (1803) Credited with being the Father of Modern Chemistry Developed 4 postulates Elements are made up of tiny particles called atoms Atoms of a given element are identical; atoms of different elements are different Chemical compounds are formed when atoms of different elements combine with each other Chemical reactions involve reorganization of atoms; the atoms themselves are not changed in a chemical reaction

19 ATOMIC THEORY DEVELOPMENT J.J Thomson (1897) Credited with the discovery of the electron Max Planck (1900) Described packets of energy called quanta Albert Einstein (1905) Described photoelectric effect Described wave-particle duality of radiation Robert Millikan (1909) Discovers magnitude of electron charge and mass of the electron through oil-drop experiments Ernest Rutherford (1911) Gold-foil experiment Proposed the nuclear atom Atom is mostly empty space

20 ATOMIC THEORY DEVELOPMENT Neils Bohr (1913) Electrons move in fixed orbits Model only worked for hydrogen atoms Louis de Broglie (1923) Proposes particle wave behavior of electron Particle-wave duality Erwin Schrodinger (1926) Formulates an equation to determine probability of electron location Quantum Theory James Chadwick (1932) Discovers neutron

21 MODERN VIEW OF ATOMIC STRUCTURE Nucleus contains protons and neutrons Protons determine element s identity Electrons move about the nucleus at an average distance of 10-8 cm Electrons determine element s chemical properties

22 ELEMENT SYMBOLS Atomic number Z Number of protons Mass number A Number of protons and neutrons NOT the number that is found on the periodic table Average atomic mass Isotopes Atoms with same number of protons but different number of neutrons

23 CHEMICAL NOMENCLATURE Know how to properly name and write the formula for: Binary ionic compounds (Type I) Contains a metal and non-metal Metal only forms a single type of cation Example NaCl Binary ionic compounds (Type II) Contains a metal and non-metal Contains a metal that forms more than one type of positive ion so charge must be specified using Roman numerals Example HgO Ionic compounds with polyatomic ions Must memorize list of common polyatomic ions See hand-out

24 CHEMICAL NOMENCLATURE Know how to properly name and write the formula for: Binary covalent compounds (Type III) Formed between two non-metals Use prefixes! Example P 4 O 10 Acids When dissolved in water, some molecules produce a solution containing free H+ ions Example HClO 4

25 CHAPTER 3 A V E R A G E A T O M I C M A S S, T H E M O L E, M A S S %, C H E M I C A L E Q U A T I O N S, A N D S T O I C H I O M E T R Y

26 AVERAGE ATOMIC MASS Most elements occur in nature as a mixture of isotopes Thus, atomic masses on periodic table are usually average values Average atomic mass % natural abundance = atomic mass % natural abundance atomic mass

27 THE MOLE By definition, the mole is the amount of substance that contains as many entities as there are in exactly 12 g of carbon-12 One mole of something consists of x units of that substance The mole is just a really big number! The mass of 1 mole of an element is equal to its average atomic mass in grams! Molar mass

28 Welcome to Mole Island 1 mol = molar mass 1 mole = 22.4 STP 1 mol = 6.02 x particles

29 MASS PERCENT There are two common ways of describing the composition of a compound: Numbers of its constituent atoms % by mass of its elements Mass % Obtained by comparing the MASS OF EACH ELEMENT present in 1 mole of the compound to the TOTAL MASS of 1 mole of the compound A pure compound should show the same percent mass of each element consistently So given a formula, you should be able to figure out the percent mass of each element Mass % = mass of element in 1 mole of compound 100 mass of 1 mole of compound

30 CHEMICAL FORMULAS Empirical formula The lowest whole number ratio of elements in a compound CH 2 Molecular formula The actual ratio of elements in a compound C 2 H 4 C 3 H 6 Empirical and molecular formula can be the same! H 2 O

31 CALCULATING EMPIRICAL FORMULAS FROM MASS % Pretend that you have a 100 gram sample of the compound Change the % to grams Convert the grams to moles for each element How do we do this again? Write the number of each element as a subscript in a chemical formula Keep each number as a decimal at this point! Divide each subscript by the smallest number Multiply the result by some integer to get rid of any fractions May not be necessary

32 HOW TO CONVERT BETWEEN EMPIRICAL FORMULAS AND MOLECULAR FORMULAS Since the empirical formula is the lowest ratio, the actual molecule would weigh more Molecular formula can always be obtained by multiplying by some whole number To do so, divide the actual molecular molar mass (usually given in the problem) by the mass of 1 mole of the empirical formula Gives whole number you MUST multiply the empirical formula by to get the molecular formula x molar mass empirical formula mass

33 CHEMICAL REACTIONS A chemical change involves the reorganization of the atoms in one of more substances Chemical equation represents this process with the reactants on the left side of the arrow and the products on the right side of the arrow Chemical reactions follow the Law of Conservation of Mass so the equation must be balanced! Chemical equation for a reaction gives: The nature of the reactants and products Physical states! The relative number of each

34 INFORMATION CONVEYED BY THE BALANCED EQUATION FOR THE COMBUSTION OF METHANE

35 STOICHIOMETRIC CALCULATIONS HOW TO CALCULATE AMOUNTS OF REACTANTS AND PRODUCTS

36 STOICHIOMETRIC CALCULATIONS WITH LIMITING REACTANT

37 THEORETICAL YIELD AND PERCENT YIELD Theoretical yield The amount of a product formed when the limiting reactant is completely consumed Determined using stoichiometry! Percent yield Percentage of theoretical yield that is actually produced in the laboratory Actual yield x 100 = percent yield Theoretical yield

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