Dr. Kevin Moore CHM 111

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1 Dr. Kevin Moore CHM 111

Total mass of materials before and after a chemical reaction must remain constant Matter is conserved (neither created nor destroyed) Tin and Sulfur react such that 30.

2 Total mass of materials before and after a chemical reaction must remain constant Matter is conserved (neither created nor destroyed) Tin and Sulfur react such that 30.0 g of Tin consumes 16.2 g of Sulfur. A student completely reacts precisely 44.2 g of Tin. How much sulfur is present in the resulting compound? gsn gs gsn 239. g Sulfur

3 John Dalton Matter is composed of small indivisible particles called atoms Introduced concept of atom Atoms of the same element have the same mass and properties Compounds are composed of atoms of different elements combined in small whole number ratios Chemical reactions are the rearrangement of atoms

4 If two elements form more than one compound, the mass of one element between the two compounds will be in small, whole number ratios. Carbon forms two compounds with Oxygen CO grams of Carbon will require 5.22 grams of Oxygen Ratio is 2.00 : 5.22 or 1.00 : 2.66 CO 5.00 grams of Carbon will require 6.65 grams of Oxygen Ratio is 5.00 : 6.65 or 1.00 : 1.33 The ratio of Oxygen between the two compounds 2.66 : 1.33 or 2:1

5 Electron Cathode Ray Tube Animation

6 Video Rutherford s Gold Foil Experiment

7 Millikan s Oil Drop Experiment Oil Drops passing through air which has been exposed to X-Rays became charged Caused drop to fall between two electrical plates Determined voltage needed to suspend drops x Coulombs Millikan Oil Drop

8 Most of atom is Empty Space Nucleus is very dense and positive Proton large positive charged particle in nucleus (1.672 x g) Neutron large neutral charged particle in nucleus (1.674 x g) Electron small negative charged particle outside of nucleus (9.109 x g)

Atoms is composed of protons and neutrons in the nucleus Electrons are outside of the nucleus Protons and electrons have different mass, but identical, yet

9 Atoms is composed of protons and neutrons in the nucleus Electrons are outside of the nucleus Protons and electrons have different mass, but identical, yet exactly opposite charges Atomic Number (Z) # of protons in the nucleus Unique to the atom # on the Periodic Table Mass Number (A) # of protons and neutrons in the nucleus

10 Isotope - atoms of the same element which have different mass numbers Only the # of neutrons can change in the nucleus Isotopic information on the left Charge and # of atoms on the right A Z # #

11 Almost every element has more than one stable isotope Video 63 Cu 65 Cu 35 Cl 37 Cl 29 p 29 p 17 p 17 p 29 e 29 e 17 e 17 e 34 n 36 n 18 n 20 n

12 Atom is neutral unless a charge is indicated Atoms take on a charge by gaining and losing electrons Changing e- has almost no effect on the atom s mass Cu Cl 29p 27e 36n 17p 18e 18n

13 Atomic mass is expressed in Atomic Mass Units Each isotope has a relative abundance Determined experimentally, known for all elements 35 Cl = 75.4% Each isotope has an isotopic mass Should be very close to mass number 35 Cl = amu Atomic mass is the weighted average mass of all isotopes amu ( isotopic mass) % abundance all isotopes

14 Calculate the average atomic mass for Boron. Boron exists as two isotopes: 10 B(19.9%, amu) and 11 B (80.1%, amu). Video B amu B amu B B amu

15 Calculate the average atomic mass for Chlorine. Chlorine exists as two isotopes: 35 Cl(75.4%, amu) and 37 Cl(24.6%, amu). Cl amu Cl amu Cl * * Cl amu

16 Atom - smallest particle which contains the properties of the element Covalent Bond - simple force created by the sharing of electrons Ionic Bond - strong force created by the exchange of electrons The Chemistry of any element is defined by how its electrons can be manipulated.

17 Compound formed by covalent bonds between the atoms Produced between 2 non-metals Simplest Molecules are Diatomic Elements Hydrogen H 2 Chlorine Cl 2 Nitrogen N 2 Bromine Br 2 Oxygen O 2 Iodine I 2 Fluorine F 2

18 Molecular Formula - elements and their whole # ratios H 2 O, N 2 O, CH 4, C 2 H 5 OH, C 8 H 18 Structural Formula - elements and their actual connection to other elements H-O-H N-N-O

19 Period - Horizontal Row 7 periods Group (Family) - Vertical Column Elements in a group have similar chemical properties 18 groups (32 with Rare Earth Metals) Arranged in Order of Atomic Number Closely follows Atomic Mass

20 Latin Names Element Name Antimony Copper Gold Iron Lead Mercury Potassium Silver Sodium Tin Tungsten Symbol Sb Cu Au Fe Pb Hg K Ag Na Sn W Atomic #

21 1 st column Highly Reactive Reactivity Increases down the table Never found naturally in the pure state Na found as Na + K as K + Soft, Low Densities React with H 2 O to make Base 2Na() s 2H O() l 2NaOH( aq) H ( g) 2 2

22 2 nd Column Reactive, but less than Alkali Metals Reactivity increases down the Periodic Table Soft, but harder than Alkali Metals React with H 2 O to make Base Mg() s H O() l Mg( OH)( aq) H ( g)

23 Next to Last column Highly reactive non-metals Reactivity increases UP the Periodic Table Powerful oxidizing agents Gases, Liquids and Solid Colorful compounds in pure state Combine with Hydrogen to make Acids HF, HCl, HBr, HI H ( g) Cl ( g) HCl( g) 2 2 2

24 Last Column Almost completely non-reactive gases Very Few compounds exist Melting points and Boiling points differ by less than 10 ºC Very narrow range to exist as a liquid Argon is 1.3% of air (by mass) Discovered and obtained from fractional distillation of air

25 Properties Good conductors of heat and electricity Shiny surface (high luster) Malleable Ductile Found on the left side of the table Solids at Room Temperature Most elements on the Periodic Table

26 Properties Poor conductors (heat and electricity) Colorful substances in the pure state Many gases and a few soft solids Found on the extreme right side of the table

27 Also called metalloids Properties are between metals and non-metals Stair-step of elements that separate the two main types Left of the stair step Metal Right of the stair step Non-metal Aluminum is generally not considered to be a semi-metal

28 Main Group or Representative Elements Elements in 1 st two and last 6 columns Make up the most abundant elements Transition Metals Center Block of metals (called d-block ) Inner Transition Metals Block at the bottom of the table ( f-block )

29 Metals Non-metals

30 Standard State physical state of an element at 25 C Solids Liquids Mercury, Bromine Gallium has a very low melting point (29.8 ºC) Gases Noble Gases Hydrogen, Nitrogen, Oxygen, Fluorine, Chlorine

31 Molecules are held together by shared electrons Covalent Bonds Other particles are held together by electrostatics Coulombs Law - Force of attraction of two charged particles increases as the magnitude of the charges increases and decreases as the distance between the charges increases F zz d

32 Ion Atoms or Molecules which carry a charge Cation Positively charged Ion Anion Negatively charged Ion Ions involve changing # s of electrons not protons Ionic Compounds - compounds composed of ions Oxidation state - charge on an individual atom

33 Hard Solids High Melting Points Form an array of connected ions Always represented by the simplest ratio of atoms (Empirical Formula) Interact with polar solvents (WATER!) NaCl, CaCl 2, AgCl, KNO 3, CaCO 3 Ionic Compound = Metal + non metal

34 Molecule which carries a charge Normally anion List on page 62 (need to know) Broken into +1, -1, -2 and -3 groups Most contain oxygen -ate and -ite NH 4 + NO 2 HCO 3 SO 3 2 ClO 4 CN CO 3 2 PO 4 3 NO 3 OH SO 4 2 MnO 4

35 Ions ending in ite have 1 fewer oxygen than ions ending in ate 2 Sulfate: SO Sulfite: SO Ions which have 1 more oxygen than ate use the prefix: per- Chlorate: ClO PerChlorate: ClO 3 4 Ions which have 1 fewer oxygen than ite use the prefix: hypo- Chlorite: ClO HypoChlorite: ClO 2

36 Ions containing Chlorine exist with other halogens Bromine Iodine Chlorate ---> Bromate ---> Iodate ClO BrO IO 3 3 3

37 Acid - substance that produces H + in H 2 O Base - substance that produces OH - in H 2 O Hydronium Ion - H 3 O + H + readily attaches to H 2 O Acid + Base = Salt + H 2 O Acid Base H + OH Sour Bitter Corrodes Metals Dissolves Organic

38 Many substances have long established traditional names NaOH ---> Caustic Soda, Lye CaO ---> Lime HCl ---> Muriatic Acid N 2 O ---> Laughing Gas H 2 O > Water NH 3 > Ammonia CH 4 > Methane (Natural Gas)

39 Metals form cations Charge Increases across the period Alkali Metals (IA) - always +1 Alkaline Earth Metals (IIA) - always +2 Aluminum - always +3 Non-metals tend to form anions Charge decreases across the period Halogens (VIIA) - usually -1 Group VIA - usually -2 Group VA - usually -3

40 +1/ Varied Charges

41 Fixed Charged Metals Use the whole name of the cation Al: Aluminum, Ca: Calcium Remove the ending from the anion Add: -ide Oxygen: Oxide Chlorine: Chloride Name the Compound NaCl: Sodium Chloride CaBr 2 : Calcium Bromide MgO: Magnesium Oxide Ba 3 N 2 : Barium Nitride LiF: Lithium Fluoride SrCl 2 : Strontium Chloride

42 Determining the Ionic Formula from the Name Use the Charges All Formulas are Empirical Formulas Compound must be neutral Aluminum Chloride Al +3 Cl -1 AlCl 3 Magnesium Oxide Mg +2 O -2 MgO

43 Create a Compound from Strontium and Nitrogen and give its name Strontium: +2 Nitrogen: -3 Least Common Multiple: 6 (3)(+2) + (2)(-3) = 0 Sr 3 N 2

44 Most transition metals have multiple oxidation states Fe +2, Fe Cr +2, Cr +3, Cr +6 Name must state the charge on the cation Rule for anion is the same (-ide) Use Roman Numerals to indicate charge FeCl 2 Iron (II) Chloride PbO Lead (II) Oxide

45 Empirical Formula smallest ratio of atoms in a compound Determining the Ionic Formula from the Name Charge is given in the Roman Numeral Chromium (III) Chloride > Cr +3 Cl > CrCl 3 Lead (IV) Sulfide Manganese (VII) Oxide Copper (I) Oxide > Pb +4 S 2 > PbS 2 > Mn +7 O 2 > Mn 2 O 7 > Cu +1 O 2 > Cu 2 O

46 Determine the charge on the metal ion using algebra Anion is still a fixed oxidation state 0 # atoms * ch arge all atoms FeCl2 1x 2( 1) 0 x 2 0 x 2 Iron( II) Chloride

47 MnO 2 1x 2( 2) 0 x 4 0 x 4 Manganese (IV) Oxide

48 Molecules: Non-metal and Non-metal First atom listed is always further to the left or below it on Periodic Table Second atom uses the standard convention for ending (-ide) The number of atoms is indicated by a prefix. If the 1st atom is a single atom, mono is not used Mono 1 Tetra 4 Hepta 7 Di 2 Penta 5 Octa 8 Tri 3 Hexa 6

49 Key to Molecules is knowing the prefixes. CO 2 Carbon Dioxide CO N 2 O CCl 4 Carbon Monoxide Dinitrogen Monoxide Carbon TetraChloride SO 3 Sulfur Trioxde ClO 2 Chlorine Dioxide

50 PolyAtomic Ions - no changes are made to their names Write the name exactly without modification. Follow other rules as given. If more than one of the polyatomic ion is present in the formula, put the entire ion in parentheses. NaNO 3 Sodium Nitrate Ca 3 (PO 4 ) 2 Calcium Phosphate Fe(NO 2 ) 2 Iron (II) Nitrite HgSO 4 Mercury (II) Sulfate

51 Use Algebra to determine the compound from the name Keep the Polyatomic Ion together as a single fixed charge Cr SO SO x 3( 2) 0 2x 6 0 x 3 Chromium( III) Sulfate

52 Examples Potassium Permanganate K MnO KMnO 4 4 Yttrium( III) Carbonate 3 Y CO Y ( CO )

53 Binary Acids Hydrogen + 1 other non-metal HCl, HBr, H 2 S Change -ide to -ic Put hydro- at the beginning of the name and state that it is an acid HCl (aq) Hydrogen monochloride (Hydrogen Chloride) Hydrochloric Acid H 2 S (aq) Hydrosulfuric Acid

54 Oxy Acids (acids containing Oxygen) Made from a polyatomic ion -ite becomes -ous -ate becomes -ic NO HYDRO! HNO 3 (aq) Contains Nitrate Nitric Acid H 2 SO 4 (aq) Contains Sulfate Sulfuric Acid HClO 2 (aq) Contains Chlorite Chlorous Acid

55 Hydrogen in an acid is always +1 Charge on the anion determines the number of Hydrogens present Polyprotic H 3 PO 4-3 H + because PO -3 4 Triprotic H 2 Se - 2 H + because Se -2 Diprotic HClO 4-1 H + because ClO 4 - Monoprotic

56 Organic Chemistry Study of Carbon containing molecules HydroCarbons molecules containing only Carbon and Hydrogen Carbon tends to bond with 4 other atoms Alkanes (Chapter 23: page 978) Methane single carbon (CH 4 ) Ethane two carbons (C 2 H 6 ) Propane three carbons (C 3 H 8 ) Butane four carbons (C 4 H 10 ) All Alkanes C n H 2n+2

57 A group which replaces one or more Hydrogens on an alkane and gives it special properties Alcohol OH replaces H (R-OH) Adds ol to the end of the name CH 4 {Methane} ---> CH 3 OH {Methanol} C 2 H 6 {Ethane} ---> C 2 H 5 OH {Ethanol}

58 Amine (R-NH 2 ) from ammonia (NH 3 ) makes compound basic replace H with NH 2 drop ane from base name and change to yl methane ----> methyl ethane ----> ethyl add amine to the name CH 4 {Methane} ---> CH 3 NH 2 {Methylamine} C 2 H 6 {Ethane} ---> C 2 H 5 NH 2 {Ethylamine}

59 Alcohols readily form double bonds by dehydrating the alcohol (using H 2 SO 4 ) to form water Unsaturated hydrocarbons Contain double or triple bonds

60 Hydrocarbons which contain a double bond are alkenes -ene (C n H 2n ) Ethene (C 2 H 4 ) Hydrocarbons which contain a triple bond are alkynes -yne (C n H 2n-2 ) Ethyne (C 2 H 2 )

61 When a functional group is present, the carbon position of the group is identified Butylamine = 1-Butylamine Start at the end of the chain which makes the number the lowest 3-Butylamine does not exist 2-propene does not exist Chain branching Always name based on the longest chain 2-Ethylpentane or 1-Methylhexane?

62 Ether (R-O-R) Carboxylic Acid (R-OOH)

63 John Dalton established atom and compounds Atom has protons, neutrons and electrons Atomic mass is average of all isotopes Molecules have shared electrons Ions are charged atoms or molecules Nomenclature Name Ionic Compounds Name Molecules Name Acids

H 2 O. Chapter 9 Chemical Names and Formulas

H 2 O. Chapter 9 Chemical Names and Formulas H 2 O Chapter 9 Chemical Names and Formulas Section 9.1 Naming Ions OBJECTIVES: Identify the charges on monatomic ions by using the periodic table, and name the ions. Section 9.1 Naming Ions OBJECTIVES: