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)