Chem 115 Powerpoint Notes Week 1
Chem 115 Powerpoint Notes Week 1 Chem 115
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This 17 page Class Notes was uploaded by Danielle Gibson on Friday August 19, 2016. The Class Notes belongs to Chem 115 at West Virginia University taught by Erin Battin in Fall 2016. Since its upload, it has received 9 views. For similar materials see Fundamentals of Chemistry 1 in Chemistry at West Virginia University.
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Date Created: 08/19/16
Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Chem 115 Notes Chapter 1 Types of Numbers Used in Chemistry o Exact Numbers A number with a value that is exactly known No error or uncertainty in the value Numbers obtained by counting individual objects AND/OR defined numbers within a given measurement system. o Measured Numbers Number with a value that is NOT exactly known due to the measuring process Some error or uncertainty in the value Amount of error depends on the measuring device (increment and distance between markings) and you! Scientific Figures, and Uncertainty o Measured numbers will ALWAYS contain some uncertainty or error o This error is expressed by the number of significant figures (or digits) reported for the measured number o Something to Pay attention to: As the number of significant figures increase, uncertainty/error decreases, and the precision of the measuring device increases Rules for Determining Sig Figs o Rule for determining the Number of Significant Figures (1.) All nonzero digits ARE SIGNIFICANT: (2.) Different Types of Zeros: Trailing Zeros: Any zero to the RIGHT of the last nonzero digit WITH a decimal place. o A decimal place is needed & they ARE SIGNIFICANT Captive Zeros: Any zero between two nonzero digits o A decimal place is not needed & they ARE SIGNIFICANT Leading Zeros: Any zero to the LEFT of the first nonzero digit ARE NOT SIGNIFICANT Zeros with decimals: These essentially become TRAILING zeros and ARE SIGNIFICANT (550.00) Ambiguous Zeros: Measured numbers greater than one that end in zero(s) without a decimal o You can’t say how many sig figs are present (20, 550, etc.) o When you have an ambiguous number is MUST be written in standard/scientific exponential notation: ( 2.0 X 10 5.500 X 10 ) Rounding Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 o If the number that is removed is 5 or greater increase the last retained digit by 1 (round up) o If the number that is removed is less than 5 leave the last retained digit unchanged (round down) o ROUND AT THE VERY END OF THE CALCULATION Significant Numbers and Calculations o Addition/Subtraction Calculations: Rule: The final answer has the same number of digits past the decimal as there are in the measurement with the fewest digits after the decimal place (line them up) o Multiplication and Division Calculations: Rule: The final answer contains the same number of Sig figs as there are in the measurement with the fewest number of sig figs. Find the measurement (from the original numbers) with the fewest sig figs. Multistep Calculations o When performing multistep calculations do addition/subtraction FIRST followed by multiplication/division o DO NOT ROUND UNTIL THE VERY END!! Standard Scientific Exponential Notation o Standard Exponential Form (Scientific Notation): ): # x 10 n # is the coefficient “n” is the exponent o Converting numbers greater than 1 into scientific notation: Move decimal place to the left until you have a number between 110 (This is your coefficient ) Add the times you have moved the decimal place as n (This is your exponent) #’s > 1 have positive exponents o Converting numbers less than 1 into scientific notation: Move the decimal place to the right until you have a number between 110 (This is your coefficient) Subtract the number of times you moved the decimal place as n (This is your exponent) #’s < 1 have negative exponents Calculations with Scientific Notation o Multiplication of Numbers in Scientific Notation: Multiply the coefficients Add the exponents o Division of Numbers in Scientific Notation: Divide the coefficients Subtract the exponents o Addition/Subtraction of Numbers in Scientific Notation: Convert the numbers so they have the same exponent Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Convert the smaller exponent to the larger exponent Add/subtract the coefficients (Remember the SF rules!) Keep the exponent the same o Multistep Calculations of Numbers in Scientific Notation: Do addition/subtraction FIRST followed by multiplication/division… round at the end if necessary! Some More Numbers o The Seven Fundamental SI Units of Measure o The International System of Units: Based on the metric system, which is convenient for most people We use these units so that all chemists can “speak” to each other. o Recognize these units of measure! Measurements o You can use prefixes to account for very large or small #’s o You don’t always obtain a measurement in SI units and so you have to convert the measurement to the desired unit! These prefixes can act as CONVERSION FACTORS in dimensional analysis! Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Measurements: Unit Conversion or Dimensional Analysis o If a conversion factor is correctly applied, a unit will always cancel Measurements: Unit Conversion that Require Multiple Steps o Several conversion steps (and this will usually be the case) may be necessary to solve a problem Measurements: Temperature o Watch the video on eCampus o Temperature (T): A measure of how hot or cold one object is relative to another A thermometer is the most common means for measuring temperature Temperature is NOT the same thing as heat (q), which is the energy that flows from the object with the higher temperature to the colder object with the lower temperature o How do they work? When immersed in a substance hotter than the thermometer, heat flows from the substance through the glass into the fluid in the thermometer causing the fluid to expand and rise Opposite occurs when immersed in a substance cc thermometer Measurements: Three Temperature Scales Measurements: Converting Temperature o Three temperature scales: Celsius, Kelvin, and Fahrenheit Based on the physical state of water Celsius (◦C) Boiling point: 100 ◦C Freezing Point: 0 ◦C Kelvin (K) Absolute zero: 0 K Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Boiling point: 373.15 K Freezing Point: 273.15 K All measurements are positive values and have no ◦ sign Uses the same size degree unit as the Celsius scale o Conversion from K to ◦C or ◦C to K: K= ◦C+273.15 ◦C= K273.15 o Fahrenheit (◦F) – Differs from Celsius and Kelvin scale in its zero point and in the size of the unit of measurement. Boiling point: 212 ◦F Freezing Point: 32 ◦F o When converting from Fahrenheit to Celsius or Kelvin, you must adjust for the start point 180/100 = 9/5 100/180 = 5/9 o Conversion from ◦F to ◦C or ◦C to ◦F: ◦F= (◦C*9/5)+32 ◦C= (◦F32)*5/9 Chemistry: Properties of Matter o Watch the video on eCampus o Chemistry is the study of matter and its changes o Physical Properties: Characteristics that do not involve a change in a sample’s chemical makeup Melting Point Electrical Conductivity Solubility Density Temperature Hardness Odor o Physical Change: Occurs when a substance alters its physical form (rearranges its molecules), NOT its composition o Chemical Properties: Characteristics that do involve a change in a sample’s chemical makeup Rusting Combustion Tarnishing o Chemical Change: Any change that results in the formation of new chemical substances Chemical changes involve making or breaking bonds between atoms Example: Balloons and Traffic Light Demonstration o Intensive Properties: Independent of sample size Temperature Color Melting/boiling point Hardness o Extensive Properties: Dependent on sample size Length Volume Height Chemistry: Matter and Its Three Physical States o Solid – Has a fixed shape that does not conform to the container shape (definite volume and shape) Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 The molecules are very close together (touching) and cannot move around; they are tightly packed (can’t flow) Not compressible o Liquid – Conforms to the container shape but fills the container only to the extent of the liquids own volume (definitive volume but no definite shape) The molecules are close together (could be touching) and they move around (can flow) Not extremely compressible o Gas – Conforms to the container shape, and it fills the entire container (no definite volume or shape) Molecules are widely separated and can move around freely Compressible Chemistry: States of Matter Measurements: Density o Density (d) – The density of an object is its mass divided by its volume o Density is commonly measured in g/cm or g/mL o You can use the knowledge of density to your advantage: A substance whose density is greater than that of the liquid will sink in the liquid A substance whose density is less than that of the liquid will float on the liquid A substance whose density is equal to that of the liquid will remain wherever placed in the liquid and will neither sink nor float o You can use mass and volume to determine density: o You can use mass and density to determine volume: o You can use volume and density to determine mass: Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Chapter 2 Classification of Matter What are elements really made of? o Greek philosophers like Plato, Aristotle, and Democritus speculated about what an element was made of…but technology was in issue in 400 B.C. o Much later it was found that… Elements are composed of atoms o Atomic Theory: provides us with ideas about the structure, properties, and behavior of atoms. We Know What an Atom is made of! (i.e. Structure) o An atom is electrically neutral and consists of three subatomic particles o Different atoms have different amounts of subatomic particles! o In a neutral atom, # protons = # electrons because the overall charge of an atom = 0 The Periodic Table and Atomic Notation o Atomic Number (Z): the number of protons in the nucleus of an atom o Mass Number/Atomic Mass Number (A): the total number of protons and neutrons in the nucleus of an atom Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 o Determining # of Neutrons: # Neutrons = Mass # (A) – Atomic # (Z) o You can determine the number of each subatomic particle for every atom using the periodic table What are Isotopes? o Isotope: atoms of an element having: The same # of protons Different # of neutrons o So…. The same atomic number (Z) Different mass number (A) o Determining # of Neutrons: # Neutrons = Mass # (A) – Atomic # (Z) o Example: Carbon has three naturally occurring isotopes: 12C, 13C, 14C o The number of neutrons in isotopes are different; however, the protons and electrons remain the same! Atomic Mass Unit (amu) o If we consider the actual mass of the proton, neutron, and even electron in grams we would have rather “large” atomic masses for the elements so Atomic masses of all elements were referenced to the atomic mass of the most abundant isotopes of carbon ( C) in amu for convenience. o 6 protons and 6 neutrons 12 o Atomic Mass Unit (amu): Based on Carbon12 or C 1 atom C = 12 amu (exactly) 1 amu = 1/12 the mass of an atom of C = 1.660539 X 10 g 24 Average Atomic Mass Number (A) o Why is the atomic mass of carbon given as 12.011 amu instead of 12 amu? Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 o Mass Number (Atomic Mass # (A)): The weighted average of the isotopic masses of the element’s naturally occurring isotopes Atomic masses shown on periodic table are average atomic masses taking into account the different isotopes of each element and their percent abundances (isotopic abundances) The Atomic Mass Number also tells you which isotope is most abundant! Examples: o Calculating Average Atomic Masses: Example: It is found that carbon consists of two naturally occurring isotopes (12C and 13C) with atomic masses and % abundances given below. Calculate the average atomic mass of carbon. This is a made up problem because you know carbon has 3 naturally occurring isotopes Isotope Atomic Mass % Abundance 12 13 12 amu 98.89% C 13.0034 amu 1.11% o Calculating % of Abundance: Example: A sample of naturally occurring gallium has an average atomic mass of 69.7 and consists of two isotopes, gallium69 and gallium71. Given the information shown below, calculate the % isotopic abundances of the two isotopes Isotope Atomic Mass 6Ga 68.9 amu 71 Ga 70.9 amu Matter: Elements & the Periodic Table o Element is the simplest type of matter with unique physical and chemical properties that consist of only one type of atom Elements cannot be broken down The periodic table organizes the elements and are represented by symbols Organization of the Periodic Table o Groups are vertical columns o Periods are horizontal columns o Metals: Shiny solids at room temperature (Hg is exception) Conduct heat and electricity Approximately ¾ of the periodic table are metals Does not include H o NonMetals: Gases or dull brittle solids at room temperature (Br is exception) Conduct heat and electricity poorly Includes H o Metalloids: Have properties between those of a metal and nonmetal Also called semimetals Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 B, Si, Ge, As, Sb, Te, Po, At Organization of the Periodic Table o Alkali Metals: Metals in Group IA (Li, Na, K, Rb, Cs) Have similar properties o Alkaline Earth Metals: Metals in Group IIA (Be, Mg, Ca, Sr, Ba) Have similar properties o Both Alkali Metals and Alkaline Earth Metals are found in the SBlock of the Periodic Table o DBlock: Transition Metals o FBlock: Lanthanide and actinide Metals o Halogens: Nonmetals in group VIIA (F, Cl, Br, I, At) Have similar properties Exist as diatomic molecules in elemental form at RT (F 2 Cl 2 Br2, 2 ) o Noble Gases: Nonmetals in Group VIIIA (He, Ne, Ar, Kr, Xe, Rn) Have similar properties: not very reactive and do not readily form compounds Exist as monatomic gases at RT (He(g), Ne(g), Ar(g), etc.) o Both Halogens and Noble Gases are found in the PBlock of the Periodic Table The PBlock contains metals, semimetals, and nonmetals Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Matter: Mixtures o Mixture is a group of two or more pure substances (elements and/or compounds) that are PHYSICALLY intermingled No chemical changes of the individual substances NOT chemically combined NOT combined in fixed proportions by mass o Types of Mixtures: Homogenous mixtures Heterogeneous mixtures Matter: Compounds o Compound is a substance that is composed of two or more elements in fixed proportions that are CHEMICALLY combined Compounds can be broken down through chemical changes The chemical and physical properties of compounds are different from the elements that formed them Compounds: Chemical Formulas and Formula Units o Chemical Formula – Shorthand way of writing the chemical symbols of each element in a compound Subscripts are used to indicate the number of atoms of each element present in the compound If no subscript is given, the number 1 is understood o Also called a formula units for ionic compounds o Also called a molecular formula for molecular compounds Compounds: Ionic vs Molecular o Depending on what type of bond is formed will determine the type of compound formed o Ionic Compounds: An electrostatic attraction between a positive ion and a negative ion, where one or more electrons have been transferred from the valence shell of one atom to the valence shell of the other atom An ionic compound typically forms between a metal and non metal Characteristics: good electrical conductors if the ions are mobile (liquids, solutions). Tend to be hard, brittle, often with high melting points o Molecular Compounds: Sharing valence electrons between atoms of different elements form COVALENT bonds Typically formed between NONMETALS Characteristics: poor electrical conductors; small molecules may be gas, liquid or soft solids Ionic Compounds: Cations and Anions o What is an ion and how are they formed? Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 o Ion: A charged species that forms when an atom gains or loses an electron The ion becomes charged either positively or negatively Cation positively charged ion Formed when atoms lose electron(s) Metals tend to form cations Anion negatively charged ion Formed when atoms gain an electron(s) Nonmetals tend to form anions Example: How many of each type of subatomic particle (# e, # p, and # n) are present in the following? 1O 2 32S2 Cs +1 Ionic Compounds: Why Do Atoms Gain or Lose Electrons? o Every element wants to obtain noble gas configuration and an “octet”. Why? Because NGC/octet is the most stable and lowest in energy In order to obtain this, elements will donate or accept electrons (valence electrons) from the valence shell Ionic Compounds: How Many Electrons are Gained or Lost? o Monatomic Ions: Ions consisting of ONE atom Example: Cl, Al , Zn , N 3 o Charges on monatomic ions can be predicted from positions in the periodic table… so in case you forget how to determine the charge of an ion you can use this method Metals Nonmetals Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 IA IIA BGroups IIIA IVA VA VIA VIIA VIIIA +1 +2 Variable +3 4/+4 3 2 1 0 +2 +3 (Fe /Fe ) o Transition and posttransition metals (metals in Groups IIIA –VIA) have variable charges Example Sn /Sn , Sb /Sb+3 +5 Ionic Compounds: Charges of Ions o Some transition and posttransition metals form more than one cation. Some elements would rather SHARE than TRANSFER their electrons! Ionic Compounds: Polyatomic Ions o Polyatomic Ions: Ions consisting of two or more elements The atoms in the polyatomic ions are held together by molecular/covalent bonds BUT then go on to form ionic compounds! So when asked what types of bonds are present in a compound with a polyatomic ion…you say both! The charge shared over all atoms in the ion and atoms stay together as a unit 2 3 2 2 Example NO , 3O , C2 , PO 3 SO ,4SO , e4c. 3 o Brackets are used when 2 or more polyatomic ions appear in the ionic formula unit 2+ Ex. Ca(NO )3 2= 1 Ca cation & 2 NO pol3atomic anions o Sorry, but you are going to have memorize the list of common polyatomic ions You need to know the name, formula, and charge A incomplete list can be found on last page of the syllabus or pg. 62 in the textbook Ionic Compounds: Common Polyatomic Ions You Need To Know!!! Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Compounds: Naming Ionic Compounds o In the name of an ionic compound, the cation is always given first and the anion second The cation is always specified as the name of the metal and is NOT changed The anion is specified by using the first part of the nonmetal name and then adding the suffix “ide” EXCEPTION: Don’t change ending of polyatomic ions!! Examples: KCl Al O MgSe CaCO 2 3 3 o For ionic compounds, you do not indicate the numbers of each type of atom… so NO prefixes! Compounds: Naming Ionic Compounds o Transition metals can often have several stable positive ions (variable charge)… that is why we skip the Dblock when assigning charge! • Example: Fe and Fe , Cu and Cu 2+ ***The modern system for naming indicates the charge on the cation as a roman numeral within the name of the substance: Iron (II) chloride FeCl2 Copper (I) chloride CuCl Iron (III) chloride FeCl3 Copper (II) chloride CuCl2 An older system, still in common use, uses the suffixes –ous and –ic to indicate the charge on the cation: Ferrous chloride FeCl2 Cuprous chloride CuCl Ferric chloride FeCl3 Cupric chloride CuCl2 Ionic Compounds: Transferring Electrons and Neutral Compounds o In the formation of sodium chloride, one electron is transferred from the sodium atom to the chlorine atom BUT…. the overall charge of an ionic compound = 0!! Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 Balancing Chemical Formulas o When balancing the chemical formulas for ionic compounds, you must consider charges…this is different with covalent compounds The overall charge of the compound must be = to zero! o Use the *CrissCross Method* for generating balance chemical ionic formulas o You can also used a balanced chemical formula to determine the charge of a transition element! Molecular Compounds: Covalent Bonds o Molecular Compounds (aka Covalent compounds) have covalent bonds o Covalent Bond: A form of chemical bonding characterized by sharing of valence electrons between atoms Typically formed between NONMETALS o Molecule is the smallest bit of a molecular compound Molecular Compounds: Common Diatomic Molecules o All of the gaseous nonmetallic elements except for the noble gases exist as diatomic (2 atoms) molecules o Be able to recognize these! Compounds: Naming Molecular Compounds o Naming Molecular Compounds: Name the element that comes first in the chemical with prefix Exception: Mono is never used to name the 1 element Name 2 element 2 but add –ide to end of name and use prefix Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 o Note: Covalent compounds with names that contain two vowels sidebyside usually end up dropping the “a” Typical with oxygen o Prefixes Some Special Covalent Compounds: Acids + o Acid – A substance that provides H ions in water (aq) o Nomenclature (Naming) of Acids (aq): Hint: H usually written first in chemical formula of acids Exception: Acetic Acid CH COO3 o Decide whether acid is a BINARY or OXYACID Binary Acid – contains H and one other element. Ex. HCl(aq), HF(aq), H S2aq) Oxyacid – contains H, O, and one other element. Ex. HClO ,4H S2 , 3NO 2 o FYI: Base – A substance that provides OH ions in water (aq) Acids Continued o Naming of Binary Acids Start with hydro prefix Use the root of nonmetal for the middle of the name Add –ic suffix after the root of the nonmetal Add Acid as a separate word o Naming of oxyacids: There are 7 common oxyacids… so 28 names! Usually four different oxyacids can be formed from each nonmetal, specify which is present by adding following prefixes and suffixes PREFIX SUFFIX Hypo ous oxyacid with least oxygens ous ic Per ic oxyacid with most oxygen Compounds: Naming Polyatomic Ions o Most polyatomic ions are anions derived from oxyacids by removal of one or more protons (H )+ Danielle Gibson Chem 115 Erin Battin Class 1 Notes 1/12/16 1/6/16 o “ous” ”ite” o “ic” ”ate” In the Pursuit of Studying the Atom… We Found Out About Compounds Too o Fundamental Chemical Principles: o Law of Mass Conservation – Mass is neither created nor destroyed in chemical reactions Antoine Lavoisier o Law of Definite Proportions – Different samples of a pure chemical substance always contain the same proportion of elements by mass Joseph Proust
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