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Chapter 2:Atoms, Molecules, & Ions
Law of Conservation of Mass
Antoine Lavoisier (late 1700’s)
Mass cannot be created or destroyed, merely converted from one form into another.
mass before rxn = mass after rxn
example: 10.62 g mercury is allowed to react with oxygen gas resulting in the formation of 11.47 g of mercury oxide.
Determine the mass of oxygen gas consumed in this reaction.
Law of Definite Proportions
Joseph Proust (late 1700’s - early 1800’s)
A given compound always consists of the same proportion of its constituent elements by mass.
example: sodium chloride (salt)
1.00g salt ! 0.394 g Na & 0.606 g Cl
100 g salt ! 39.4 g Na & 60.6 g Cl
Law of Multiple Proportions
John Dalton (late 1700’s - early 1800’s)
When 2 elements form a series of compounds, the ratio of the mass of the 2nd element that combines with 1 g of the 1st element is a small whole number ratio.
example: Consider three different compounds composed of sulfur and fluorine:
mass of sulfur mass of fluorinecompound A 1.000 g 1.188 gcompound B 1.000 g 2.375 gcompound C 1.000 g 3.563 g
Dalton’s Atomic Theory - 4 basic statements:1. All matter is composed of atoms.
2. Atoms of the same element are identical; atoms of different elements are distinct.
- with respect to physical and chemical properties
3. A compound is a specific combination of atoms of 2 or more elements bonded together in a specific ratio.
4. In a chemical reaction atoms are neither created or destroyed, merely rearranged into new substances.
- Law of Conservation of Matter
In 1808, John Dalton published A New System of Chemical Philosophy.
Early Experiments to Establish Atomic Structure:
J.J.Thomson and Cathode RaysRobert Millikan and the Oil Drop Experiment
Ernest Rutherford and the Nuclear Structure of the Atom
J.J. Thomson and Cathode Rays (1897)
Thomson’s Conclusions:
negatively charged particles (electrons, e!) are common to the internal structure of atoms of all elements
charge to mass ratio of electron: e/m = 1.76 x 108 C/g
Millikan’s Oil Drop Experiment (1909)
Millikan determined the charge of a single electron:1 e! = 1.6 x 10 −19 C
coupled with Thomson’s results: mass of e− = 9.1 x 10−28 g
Ernest Rutherford and α-Rays (1911)
● beam of alpha-particles shot at thin gold film
● Rutherford expected to see very little or no deflection of particle beam
● approximately 1 in 20,000 particles were deflected sharply backward
Rutherford: the Nuclear Structure of the Atomthe nucleus: center of mass and positive charge
very small relative to the whole atom
Atomic Structure: Protons, Neutrons, & Electrons
note: an atom is charge neutral when number protons = number electrons
Individual atoms identified by:name and symbol
atomic number (Z)mass number (A)
atomic number, Z:
Z = number of protons in the nucleus
Z does not change for a given element
mass number, A:
A = number protons + number neutronsor equals total number of nucleons
number of protons is constant for atoms of the same element, but the number of neutrons is variable
Isotopes - atoms with the same number of protons but differing numbers of neutrons
same atomic number ∴ same elementdifferent mass numbers
consider 2 atoms (atom A and atom B are isotopes of one another):
atom A atom Bnumber protons 50 50number neutrons 69 73number electrons 50 50atomic numbermass number
symbol
Elements and the Periodic Tableelements:
◆ pure substances◆ only one type of matter◆ cannot be broken down into simpler substances◆ 118 known; 112 named
each element has a name and a symbol:
◆ an element’s name should not be capitalized unless it is the 1st word in a sentence
◆ for elemental symbol - 1st letter is always capitalized, 2nd letter is never capitalized
★★★ In Chem 1711 you need to know the names (correctly
spelled) and symbols for elements 1 - 56 plus W, Pt, Au, Hg, Pb, & Bi
Periodic Table of the ElementsOrganization of the Periodic Table:
Periods and Groups
horizontal rows: periods
vertical columns: groups or families
numbering of groups - IA to VIIIA; 1A to 8; 1 to 18groups with special names -
IA (1) alkali metalsIIA (2) alkaline earth metalsVIA (16) chalcogensVIIA (17) halogensVIIIA (18) noble (inert) gases
Organization of the Periodic Table: Metals, Nonmetals, and Metalloids
Metalsleft of stair-step line in PT
mostly solids in elemental form (a few liquids)
metallic luster, malleable, and ductile
good conductors of heat and electricity
Nonmetalsright of stair-step line in PT
can be solids, liquids, or gases in elemental form
not ductile or malleable
not conductors
Metalloidselements along the stair-step line in PT
B, Si, Ge, As, Sb, Te, Atintermediate properties
You should be able to . . .
◆ find an element in the periodic table
◆ give the name, symbol, and atomic number for the element
◆ identify the position of the element by period and group number
◆ identify (if appropriate) group or family by name
◆ identify if the element is a metal (main group vs. transition), nonmetal, or metalloid
◆ give the proton, neutron, and electron count for the atom if mass number is known
Masses of individual atoms of all elements determined using mass spectrometry.
relative to standard: 1 atom 12C = 12 amu
for all elements, mass range per atom: 1 - 260 amu10!24 - 10!22 g
Average Atomic Mass
average atomic mass - weighted average mass (in amu) calculated using the masses of naturally occurring isotopes, and their abundances
ex. Calculate the average atomic mass of chlorine given the following information:
isotope isotope mass percent abundance
35Cl 34.969 amu 75.78%
37Cl 36.966 amu
average atomic mass = [(mass of 35Cl)(%abund35Cl/100)] + [(mass of 37Cl)(%abund37Cl/100)]
Chromium has four naturally occurring isotopes:
isotope isotope mass abundance50Cr 49.9461 amu 4.350%52Cr 51.9405 amu 83.79%53Cr 52.9407 amu 9.500%54Cr 53.9389 amu 2.360%
Considering only this data given (i.e. do not look at the periodic table), what approximate average atomic mass do you expect for Cr?
Then - calculate the average atomic mass of Cr and compare your answer to the reported average atomic mass in the periodic table.
Nomenclature
naming of chemical compounds
rules for determining the IUPAC names for:
binary molecular compoundsionic compoundsbinary acids and oxoacids
Molecular Compounds
◆ consist of molecules
◆ molecule - neutral group of atoms held together by covalent bonds (shared electrons)
◆ molecular compounds are composed of nonmetals only, or nonmetals and metalloids
◆ binary molecular compounds - 2 elements only
Ionic Compounds◆ consist of ions
◆ ions are charged species*; ions can be:monatomic or polyatomic
cations or anions
◆ individual ions are charged, but overall an ionic compound is charge neutral
◆ cations and anions held together by ionic bonds (electrostatic attraction of opposite charges)
◆ composed of metals and nonmetals
* ions are formed by the gain or loss of electrons ion charge = # protons ! # electrons
Naming Monatomic Cations - Type I Metals
metals tend to form cations2 types of metals
Type I metals - metals that form cations of only one charge
cation charge = metal’s group number
group IA, group IIA, group IIIB, Al3+, Ga3+, Zn2+, Cd2+, Ag+
examples: Na+ sodium ion Cs+ cesium ion Mg2+ magnesium ion Ba2+ barium ion Al3+ aluminum ion Zn2+ zinc ion
Naming Monatomic Cations - Type II Metals
Type II metals - metals that form cations with more than one possible charge
◆ these tend to be the transition metals and heavier main group metals
examples: Fe2+ & Fe3+, Co2+ & Co3+, Cr3+ & Cr6+, Sn2+ & Sn4+
naming type II metal cations: use stock system to indicate charge on cationCr3+ chromium (III) ionCr6+ chromium (VI) ionSn2+ tin (II) ion Sn4+ tin (IV) ion
Naming Monatomic Anionsnonmetals tend to form anions
monatomic anions: primarily nonmentals in groups VA, VIA, and VIIA
anion charge = group number ! 8
naming monatomic anions: _________ ! ide ion
examples: N3! nitride ionAs3! arsenide ionO2! oxide ionSe2! selenide ionF! fluoride ionCl! chloride ion
Naming Polyatomic Anions
◆ mostly oxoanions (i.e. contain oxygen)◆ need to know name, formula, and charge
to name an oxoanion, basic rule: ________-ate ion
examples:CO32! carbonate ionSO42! sulfate ionNO3! nitrate ionPO43! phosphate ionClO3! chlorate ion
Naming Polyatomic Anions
◆ Some elements form more multiple oxoanions, each with a different number of O’s.
name these in the following way:
______-ate ion for the oxoanion with more O’s
______-ite ion for the oxoanion with less O’s
examples:SO42! sulfate ionNO3! nitrate ionClO3! chlorate ion
SO32! sulfite ionNO2! nitrite ionClO2! chlorite ion
Naming Polyatomic Anions
◆ Cl, Br, and I each form series of four different oxoanions with 1 - 4 O’s.
to name the ions in these series:
per-______-ate ion (most O’s) ______-ate ion (more O’s) ______-ite ion (less O’s) hypo-______-ite ion (least O’s)examples:
BrO4! perbromate ion IO4! periodate ionBrO3! bromate ion IO3! iodate ionBrO2! bromite ion IO2! iodite ionBrO! hypobromite ion IO! hypoiodite ion
Naming Polyatomic Anions
◆ Protonated oxoanions:have one or more proton (H+)formula, charge, and name all change
examples:
CO32! carbonate ion HCO3! hydrogen carbonate ion
SO32! sulfite ion HSO3! hydrogen sulfite ion
O2! oxide ion OH! hydroxide ion
PO43! phosphate ion HPO42! hydrogen phosphate ion H2PO4! dihydrogen phosphate ion
Names and Formulas of Ionic CompoundsCation - Anion
examples:
Write the IUPAC name for each of the following:
Na3P(NH4)3PO4
V2(SO4)5
Fe(BrO4)3
Names and Formulas of Ionic Compounds
examples:
Write the chemical formula for each of the following given the IUPAC name:
magnesium nitratecopper (I) chloridecobalte (III)oxalatesilver acetate
Naming Molecular Compounds
◆ name the elements in the order that they appear◆ add suffix !ide to stem of 2nd element name◆ use prefixes to indicate how many atoms of each element are presentexamples:
Give the name for each of the following:
SiO2 PCl3 XeF6
Give the chemical formula for each of the following:
dinitrogen pentoxideboron triiodide dichlorine oxide
Naming Acids
◆ Acids - one or more H+ with an anionnumber of H+’s in acid formula = charge on anion
◆ binary acids - anion is a monatomic anion; only 2 elements present
to name a binary acid: hydro-______-ic acid
examples: HI hydrosulfuric acid
Naming Acids
◆ oxoacids - anion is an oxoanion
◆ the name of an oxoacid is determined based on the name of the oxoanion:
if oxoanion name ends in -ate: ______-ic acid
if oxoanion name ends in -ite: ______-ous acid
examples: HClO4 HNO2
acetic acid carbonic acid sulfurous acid
Naming Hydrated Metal Salts
◆ ionic compounds (aka salts) that have one or more water molecule associated per formula unit
◆ “waters of hydration”
name these in the following way:
ionic compound name + ______ - hydrate
examples:
Mg(NO3)2•4H2O magnesium nitrate tetrahydrate
CoSO4•7H2O cobalt (II) sulfate heptahydrate
Writing and Balancing Chemical EquationsReactants ! Products
substances consumed substances formed
chemical equations provide qualitative information:◆ identity of reactants and products; chemical formulas◆ physical states; s, l, g, aq◆ reaction conditions
chemical equations provide quantitative information:◆ stoichiometric coefficients◆ numbers of particles that combine◆ ratio of how reactants combine and products form
Writing and Balancing Chemical Equations
Chemical equations must be balanced before they can provide quantitative information.
◆ atom book-keeping . . . must have equal numbers of atoms of each element of both sides of the arrow
Law of Conservation of Matter - matter cannot be created or destroyed, merely converted from one form into another
Writing and Balancing Chemical Equations
to balance a chemical equation:1. write the skeletal equation
◆ chemical formulas of reactants and productsnote: there are 7 elements that exist as diatomic molecules in their element form:
H2, N2, O2, F2, Cl2, Br2, I2
2. adjust the stoichiometric coefficients until the equation is balanced
◆ do not change subscripts in chemical formulas◆ trial and error
3. DOUBLE CHECK!
4. add details if you know them◆ physical states; reaction conditions
Writing and Balancing Chemical Equations
examples:
Molten aluminum metal and solid barium oxide are combined and heated resulting in the formation of elemental, liquid barium and solid aluminum oxide.
__ Ca3P2 (s) + __ H2O (l) ! __ Ca(OH)2 (aq) + __ PH3 (g)
__ SO2 (g) + __ H2S (g) ! __ S (s) + __ H2O (l)
__ CO2 (g) + __ H2O (l) ! __ C6H12O6 (s) + __ O2 (g)
__ C4H10 (g) + __ O2 (g) ! __ CO2 (g) + __ H2O (g)
Classes of Chemical Reactions
Chemical reactions typically fall into one of the following categories:
1. Formation: formation of a compound from its elements.
ex: 6 Na (s) + N2 (g) ! 2 Na3N (s)
2. Synthesis
3. Decomposition
4. Combustion: organic compound (compound composed of C, H, O) reacts with oxygen to form CO2 and water
ex: CH4 (g) + 2 O2 (g) ! CO2 (g) + 2 H2O (g)
5. Precipitation: formation of an insoluble solid (precipitate)
ex: AgNO3 (aq) + NaCl (aq) ! AgCl (s) + NaNO3 (aq)
6. Acid-Base Neutralization: H+ transfer
ex: NH3 (aq) + HCl (aq) ! NH4Cl (aq)
7. Oxidation - Reduction: electron transfer; changes in oxidation numbers of elements
ex: Zn (s) + 2 HCl (aq) ! ZnCl2 (aq) + H2 (g)
Classes of Chemical Reactions (Cont’d)