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Chapter 2

Atoms, molecules and ions

Discovery of Electron JJ Thomson (1897) Modified cathode ray tube Made quantitative measurements on cathode rays Discovered negatively charged particles Electrons (e –) Determined charge to mass ratio (e/m) of these particles e/m = –1.76 x 108 coulombs/gram
Discovery of Proton Discovered in 1918 in Ernest Rutherford’s lab Detected using a mass spectrometer Hydrogen had mass 1800 times the electron mass Masses of other gases whole number multiples of mass of hydrogen Proton Smallest positively charged particle
Rutherford’s Nuclear Atom Demonstrated that nucleus: has almost all of mass in atom has all of positive charge is located in very small volume at center of atom Very tiny, extremely dense core of atom Where protons ( ) and neutrons ( ) are located
Discovery of Neutron postulated by Rutherford and coworkers Estimated # + charges on nucleus based on experimental data Nuclear mass based # protons always far short of actual mass About ½ actual mass must be another type of particle mass = p neutral /1932 by James Chad
Properties of Subatomic Particles Three kinds of subatomic particles of principal interest to chemists Electron -1 proton +1 neutron 0
Atomic number (Z) Number of protons that atom has in nucleus Unique to each type of element Element is substance whose atoms all contain identical number of protons Z = number of protons
Isotopes Atoms of same element with different masses Same number of protons ( ) Different number of neutrons ( )
Isotope Mass number (A) A = (number of protons) + (number of neutrons) A = Z + N For charge neutrality, number of electrons and protons must be equal
Atomic Symbols Summarize information about subatomic particles Every isotope defined by two numbers Z and A Symbolized by Ex. What is the atomic symbol for helium? He has 2 e–, 2 n and 2 p Z = 2, A = 4
Isotopes Most elements are mixtures of two or more stable isotopes Each isotope has slightly different mass Isotopes distinguished by mass number (A): e.g., Three isotopes of hydrogen (H) Four isotopes of iron (Fe)
What is the isotopic symbol for Uranium-235? Number of protons ( ) = 92 = number of electrons in neutral atom Number of neutrons ( ) = 143 Atomic number (Z ) = 92 Mass number (A) = 92 + 143 = 235 Chemical symbol = U Summary for uranium-235: 235/92U
An atom of has ___ protons, ___ neutrons, and ___ electrons. 82, 124, 82
What is the correct symbol for an element that has 27 protons, 33 neutrons, and 27 electrons Trick question: the problem has two of the same in but only one is correct because of the way the elements are arranged on the periodic table.
Carbon-12 Atomic Mass Scale Need uniform mass scale for atoms Atomic mass units (symbol u) Based on carbon: 1 atom of carbon-12 = 12 u (exactly) 1 u = 1/12 mass 1 atom of carbon-12 (exactly)
Why was 12C selected? Common Most abundant isotope of carbon All atomic masses of all other elements ~ whole numbers Lightest element, H, has mass ~1 u
Now, for a break, let’s calculate a person’s average in this course 93 homework average @ 15% 72 quiz average @ 15% 64 exam average @ 40% 58 on the final @ 30% What was her average, and what was her course grade? 67.75 C (because this particular grade scale goes from 60--75 as a c. )
Calculating Atomic Mass How do we calculate average atomic mass? Weighted average Use isotopic abundances and isotopic masses (add up all the sums of the mass x % abundance)
How do we calculate atomic mass? Average of masses of all stable isotopes of given element atomic mass = p+n
A naturally occurring element consists of two isotopes. isotope #1 68.5257 u 60.226% isotope #2 70.9429 u 39.774% Calculate the average atomic mass of this element. 69.487 u
Periodic Table Summarizes periodic properties of elements
Early Versions of Periodic Tables Arranged by increasing atomic mass Mendeleev (Russian) and Meyer (German) in 1869 Noted repeating (periodic) properties
Modern Periodic Table Arranged by increasing atomic number (Z ): Rows called periods Columns called groups or families Identified by numbers 1 – 18 standard international 1A – 8A longer columns and 1B – 8B shorter columns
A groups—Longer columns Alkali Metals 1A = first group Very reactive All are metals except for H Tend to form +1 ions React with oxygen Form compounds that dissolve in water Yield strongly caustic or alkaline solution (Na2O
A groups—Longer columns Alkaline Earth Metals 2A = second group Reactive Tend to form +2 ions Oxygen compounds are strongly alkaline (MgO) Many are not water soluble
A groups—Longer columns Halogens 7A = next to last group on right Reactive Form diatomic molecules in elemental state 2 gases – F2, Cl2 1 liquid – Br2 2 solids – I2, At2 Form –1 ions with alkali metals—salts (e.g. NaF, NaCl, NaBr, and NaI)
A groups—Longer columns Noble Gases (not “Nobel”) 8A = last group on right Inert—very unreactive Only heavier elements of group react and then very limited Don’t form charged ions Monatomic gases (e.g., He, Ne, Ar)
B groups—shorter columns All are metals In center of table Begin in fourth row Tend to form ions with several different charges e.g., Fe2+ and Fe3+ Cu+ and Cu2+ Mn2+, Mn3+, Mn4+, Mn5+, Mn6+, and Mn7+ Note: Last 3 columns all have 8B designation
Inner Transition Elements At bottom of periodic table Tend to form +2 and +3 ions Lanthanide elements Elements 58 – 71 Actinide elements Elements 90 – 103 All actinides are radioactive
Lanthanide elements Elements 58 – 71
Actinide elements Elements 90 – 103 All actinides are radioactive
Metals, Nonmetals, or Metalloids metals -most of chart nonmetals-H and then most of right hand corner metalloids-B,Si As, Ge, At, Te, Sb, Po
Classify the following three elements as a metal, non-metal, or metalloid: silicon (Si), vanadium (V), bromine (Br) metalloid, metal, metalloid
Strontium (Sr) is a _______, ruthenium (Ru) is a ________, and iodine (I) is a_________. alkaline earth metal, transition metal, halogen
Metals Properties Metallic luster -- Shine or reflect light Malleable -- Can be hammered into thin sheets Ductile -- Can be drawn into wire Hardness Some hard – Fe & chromium Some soft – Na, lead, Cu
More Properties of Metals Conduct heat and electricity Solids at room temperature Hg only liquid metal (mp = –39 °C) Tungsten (W) (mp = 3400 °C) has highest Chemical reactivity Varies greatly Au, Pt very unreactive Na, K very reactive
Nonmetals 17 elements Upper right hand corner of periodic table mostly compounds vs pure elements Brittle Pulverize when struck Insulators Non-conductors of electricity and heat inert (Noble) Some reactive (F2, O2, H2) (reactive to form ionic comp)
nonmetals gases Many are gases Monatomic (Noble) He, Ne, Ar, Kr, Xe, Rn Diatomic H2, O2, N2, F2, Cl2 Br2
nonmetals solids Some are solids: I2, Se8, S8, P4, C Three forms of carbon (graphite, coal, diamond)
nonmetals liquid One is liquid: Br2
Metalloids 8 Elements on diagonal line btwn metals and non B, Si, Ge, As, Sb, Te, Po, At Between metals and nonmetals Metallic shine Brittle like nonmetal Semiconductors Conduct electricity But not as well as metals Silicon (Si) and germanium (Ge)
Which statement is incorrect? A. Cu is a transition element B. Na is an alkaline earth metal C. Si is a metalloid in group 3A D. F is a halogen B. Na is an alkaline earth metal
All are characteristics of metals except: Malleable Ductile Good conductors of heat Act as semiconductors Act as semiconductors
Molecules and Chemical Formulas Atoms combine into compounds Useful to visualize atoms, compounds, and molecules Atoms = spheres Diff atms have diff colors Atoms combne to frm more complx substncs Discrete particles composed of 2 or more atoms ( O2, CO2, NH3, C12H22O11)
Chemical Formulas Specify composition of substance Chemical symbols Represent atoms of elements present Subscripts Given after chemical symbol Represents relative numbers of each type of atom e.g., Fe2O3 : iron and oxygen in 2:3 ratio
Free Elements Element not combined with another in compounds Just use chemical symbol to represent e.g., Iron Fe Neon Ne Sodium Na Aluminum Al
Diatomic Molecule Molecules composed of two atoms each Many elements found in nature e.g., Oxygen O2 Nitrogen N2 Hydrogen H2 Chlorine Cl2
Depicting Molecules Want to show: Order in which atoms are attached to each other 3-dimensional shape of molecule Three ways of visualizing molecules: 1. Structural formula 2. Ball-and-stick model 3. Space filling model
Structural Formulas Use to show how atoms are attached Atoms represented by chemical symbols Chemical bonds attaching atoms indicated by lines (diagragmic)
"Ball-and-Stick” Model Spheres = atoms Sticks = bonds (look like pins sticking together out of a pincushion but dr. conover thinks they look like macy's parade)
“Space-Filling” Model Shows relative sizes of atoms Shows how atoms take up space in molecule (look like balloons)
3-D Representations of Molecules Use fused spheres to indicate molecules Different colors indicate different elements Relative size of spheres reflects differing sizes of atoms
More Complicated Molecules Sometimes formulas contain parentheses How do we translate into a structure? e.g., Urea, CO(NH2)2 Expands to CON2H4 Atoms in parentheses appear twice
Counting Atoms Raised dot in formula indicates that the substance is a hydrate Example: 1 (CH3)3COH Subscript 3 means 3 CH3 groups So from (CH3)3 we get 3 × 3 H = 9 H C = 3 C + 1 C = 4 C H = 9 H + 1 H = 10 H O = 1 O TL 15 atoms
Hydrates Crystals that contain water molecules e.g., Plaster: CaSO4∙2H2O calcium sulfate dihydrate Water is not tightly held
Dehydration Removal of water by heating Remaining solid is anhydrous (without water)
Counting Atoms Example: 2 CoCl2·6H2O The dot 6H2O means you multiple both H2 and O by 6 So there are: Number of H 6 × 2 = 12 H Number of O 6 × 1 = 6 O Number of Co 1 × 1 = 1 Co Number of Cl 2 × 1 = 2 Cl Total Number of atoms = 21 atoms
Dalton’s Atomic Theory All molecules of compound are alike & contain atoms in same numerical ratio. Example: Water, H2O Ratio of oxygen to hydrogen is 1 : 2 1 O atom : 2 H atoms in each molecule O weighs 16 times as much as H 1 H = 1 mass unit 1 O = 16 mass units
Dalton’s Atomic Theory Explains Law of Conservation of Mass Chemicl reactns = rearranging atoms. Explains Law of Definite Proportions Predicted Law of Multiple Proportions -Not yet discovered
Law of Definite Proportions compound always has atoms of same elements in same ratios
Law of Multiple Proportions Some elements combine to give 2+ compounds e.g., SO2 and SO3 different masses of one element that combine with same mass of other element are always in ratio of small whole numbers. Atoms react as complete (whole) particles.
Chemical formulas Indicate whole numbers of atoms Not fractions
Molecules Small and Large Some are very large, especially those found in nature Same principles apply to all
How Do We Know Formulas? Use results of experiments to determine Formula Chemical reactivity Molecular shape Can speculate once formula is known Determine from more experiments
Visualizing Mixtures Look at mixtures at atomic/molecular level Different color spheres stand for two substances -Homogeneous mixture/solution – uniform mixing -Heterogeneous mixture – two phases (oil and vinegar)
Mole A number equal to the number of atoms in exactly 12 grams of 12C atoms How many in 1 mole of 12C ? -Based on experimental evidence 1 mole of 12C = 6.022 × 1023 atoms 12C
Avogadro’s number = NA Number of atoms, molecules or particles in one mole 1 mole of X = 6.022 × 1023 units of X ***1 mole Xe = 6.022 × 1023 Xe atoms ***1 mole NO2 = 6.022 × 1023 NO2 molecules
Moles of Compounds Atoms =Atomic Mass 1 mole of atoms = gram atomic mass = 6.022 × 1023 atoms Molecules =Molecular Mass Sum of atomic masses of all atoms in compound’s formula 1 mole of molecule X = gram molecular mass of X = 6.022 × 1023 molecules
Ionic compounds Formula Mass Sum of atomic masses of all atoms in ionic compound’s formula 1 mole ionic compound X = gram formula mass of X = 6.022 × 1023 formula units when a metal and a nonmetal bond
General Molar mass (MM) Mass of 1 mole of substance (element, molecule, or ionic compound) under consideration 1 mol of X = gram molar mass of X = 6.022 × 1023 formula units
Molar mass mass is our conversion factor between g and moles 1 mole of X = 6.022 × 1023 units of X
What is the molar mass of C2H5COOH? 74.08 g/mol
Learning Check: Using Molar Mass Example: How many moles of iron (Fe) are in 15.34 g Fe? What do we want to determine? 15.34 g Fe = ? mol Fe What do we know? 1 mol Fe = 55.85 g Fe Set up ratio so that what you want is on top and what you start with is on the bottom 15.34g Fe x (1 mol/55.85g Fe) =0.2747 mol fe
Example: If we need 0.168 mole Ca3(PO4)2 for an experiment, how many grams do we need to weigh out? What do we want to determine? 0.168 mole Ca3(PO4)2 = ? g Ca3(PO4)2 Calculate MM of Ca3(PO4)2 3 × mass Ca = 3 × 40.08 g = 120.24 g 2 × mass P = 2 × 30.97 g = 61.94 g 8 × mass O = 8 × 16.00 g = 128.00 g =1 mole Ca_3(PO_4)_2=310.18 g Ca_3(PO_4)_2
Set up ratio so that what you want is on the top and what you start with is on the ____________? bottom
How many moles of CO2 are there in 10.0 g? A. 0.0227 mol B. 4.401 mol C. 44.01 mol D. 0.227 mol D. 0.227 mol
How many grams of platinum (Pt) are in 0.475 mole Pt? A. 195 B. 0.0108 C. 0.00243 D. 92.7 D. 92.7
Using Moles in Calculations Start with either -Grams (Macroscopic) -Elementary units (Microscopic) Use molar mass to convert from grams to mole Use Avogadro’s number to convert from moles to elementary units
Macroscopic to Microscopic How many silver atoms are in a 85.0 g silver bracelet? What do we want to determine? 85.0 g silver = ? atoms silver What do we know? 107.87 g Ag = 1 mol Ag 1 mol Ag = 6.022×1023 Ag atoms g Ag  mol Ag  atoms Ag 85.0g x(1 mol /1017.87g) x ((6.022x10^23atoms)/1 mol) = 4.7 × 1023 Ag atoms
Calculate the number of formula units of Na2CO3 in 1.29 moles of Na2CO3. 1.29 mol Na2CO3 ((6.022x10^23 units)/ 1 mol) = 7.77×1023 particles Na2CO3
How many moles of Na2CO3 are there in 1.15 × 105 formula units of Na2CO3 ? 1.15 x 10^5 formula units (1 mol/(6.022x10^23)) = 1.91×10–19 mol Na2CO3
Calculate the mass in grams of FeCl3 in 1.53 × 1023 formula units. (molar mass = 162.204 g/mol) A. 162.2 g B. 0.254 g C. 1.661×10–22 g D. 41.2 g E. 2.37× 10–22 1.53 x 10^23 (1 mol /(6.022x 10^23)) x (162.2 /1 mol) = 41.2 g FeCl3
Mole-to-Mole Conversion Factors In H2O there are:  2 mol H ⇔ 1 mol H2O  1 mol O ⇔ 1 mol H2O  2 mol H ⇔ 1 mol O  on atomic scale  2 atom H ⇔ 1 molecule H2O  1 atom O ⇔ 1 molecule H2O  2 atom H ⇔ 1 molecule O
Stoichiometric Equivalencies Ex. N2O5 2 mol N ⇔ 1 mol N2O5 5 mol O ⇔ 1 mol N2O5 2 mol N ⇔ 5 mol O Ratios of atoms in chemical formulas must be whole numbers
Calculating the Amount of a Compound A sample is found to contain 0.864 moles of phosphorus. How many moles of Ca3(PO4)2 are in that sample? What do we want to find? 0.864 mol P = ? mol Ca3(PO4)2 What do we know? 2 mol P ⇔ 1 mol Ca3(PO4)2 Solution 0.86 mol P( (1 mol Ca3(PO4)2)/2 mol P) = 0.432 mol Ca3(PO4)2
There are 5.4 moles of sodium carbonate, Na2CO3, in a sample. How many moles of oxygen atoms (O) are present in the sample? A. 7.20 B. 5.40 C. 32.4 D. 16.2
Two Methods of Energy Exchange Between System and Surroundings Heat= q Work=w E = q + w 1. Heat + Heat absorbed, System’s q (goes up) - Heat lost, System’s q (goes down) 2. Work  Is exchanged when pushing force moves something through distance Ex. Compression W>0 expansion W<0
Created by: rhonda83



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