Atoms, Molecules, and Ions

Chapter 2 BLB 12th

Expectations Recognize important steps in the discovery

of the atom and its structure. Work with isotopes. Learn about the periodic table. Differentiate between molecular and ionic

compounds. Name compounds (molecular and ionic).

2.1 The Atomic Theory of MatterOf what is matter comprised? Democritus (400 BC) – tiny, indivisible particles, atomos Plato, Aristotle – NOT! Newton (17th century) – favored atoms as invisible particles Boyle (1660) – gas experiments with pressure & volume Priestly (1774) – isolated oxygen Lavoisier (1789) – Law of Conservation of Mass: Mass is

neither created or destroyed. (p.78) Proust (1800) – Law of Definite Proportions (or constant

composition): A compound always contains the same proportion of elements. (p. 10)

Dalton’s Atomic Theory (1808)1. Elements are composed of small particles

called atoms.2. All atoms of a given element are identical.3. Atoms of an element are not changed in a

chemical reaction.4. Compounds are formed when different

atoms combine.>> Atoms are the building blocks of matter.<<

2.1 The Atomic Theory of Matter

Dalton – Law of Multiple Proportions: element mass proportions in a compound are in a ratio of small whole numbers. (p. 40)

Avogadro (1811) – equal volumes of gases contain the same number of particles (p. 401)

2.2 The Discovery of Atomic Structure

J.J. Thomson (1897) – cathode ray tube experiments; electrons; charge-to-mass ratio of the electron

(1) plum-pudding model of atoms (Fig. 2.9, p. 43) Robert Millikan (1909) – oil-drop experiment; charge and

mass of electron (9.10939 x 10-28 g) Henri Becquerel, Marie Curie (1896, 1899) – radioactivity Ernest Rutherford (1911) – gold foil experiment; nucleus &

protons (1919); (2) nuclear model of atom

3 types of radioactivity: α (heaviest, 2+ charge), β (high-speed electrons, 1− charge), g (lightest, high E, 0 charge)

James Chadwick (1932) - neutrons

Subatomic particles

Separation of Radioactive Particles

Rutherford’s Gold Foil Experiment

pp. 42-43

2.3 The Modern View of Atomic Structure

Subatomic particlesParticle Function Position

(to nucleus)Charge Mass (kg)

electron chemistry outside −1 9.11 x 10-31

proton attract electrons

inside +1 1.67 x 10-27

neutron nuclear glue inside 0(neutral)

1.67 x 10-27

Atoms

Atomic masses ~10-23 g Atomic diameters (e- cloud) ~10-10 m = 1 Å Atomic nuclei ~10-4 Å (very small and dense)

Atoms are neutral: # protons = # electrons

Practice Exercise 2.1How many carbon atoms can be placed side by side across the width of a pencil line that is 0.20 mm wide? C atom diameter = 1.54 Å

Isotopes, atomic and mass numbers Isotopes – atoms with same number of protons

but different numbers of neutrons

Nuclide – a single atom of a particular isotope

symbolelementCmassatomic

12#6#

2.4 Atomic Weights

Based on 12C (assigned a mass of exactly 12 amu) 1 amu = 1.66054 x 10-24 g

(1/12 mass of a 12C atom) Weighted average atomic mass =

Σ(% abundance)(mass of isotope) Atomic mass determined using a mass

spectrometer (p. 49)

Mass Spectrometer

Mass spectrum of Cl

Calculate the (weighted) average mass of magnesium (in amu).

isotope % abund. Mass (amu)

24Mg 78.99 23.98504

25Mg 10.00 24.98584

26Mg 11.01 25.98259

2.5 The Periodic Table 1st table developed by Mendeleev and

Meyer in 1869 Group, period, regions, group names Physical properties of metals and nonmetals

Seaborg (p. 52)

Physical Properties

Metals High electrical conductivity High thermal conductivity Metallic luster Most are solids Malleable, ductile Metallic bonding

Nonmetals Poor electrical conductivity Good heat insulator No metallic luster Solids, liquids, and gases Brittle in solid state Covalently bonded

molecules; noble gases monoatomic

2.6 Molecules and Molecular Compounds

Chemical bonds – forces that hold atoms together in molecules and compounds

covalent bonds – sharing of electronsMolecules – discrete units of covalently bonded atoms;

typically nonmetals, e.g. H2O, CO2, NH3, C2H6

Diatomic elements: H2, N2, O2, F2, Cl2, Br2, I2 (p. 53)

Polyatomic elements: O3, S8, P4

(allotropes – different forms of the same element in the same state, p. 273)

Molecules, cont.

Representation of molecules, CH4

Empirical & Molecular Formulas Molecular formula – actual number of atoms in a

compound Empirical formula – smallest whole number ratio

of atomsMolecular Empirical

C2H4 CH2

P4O10 P2O5

H2O2 HO

2.7 Ions and Ionic Compounds

Ionic bond – attraction between oppositely charged ions; results from a transfer of electrons cation – positively charged ion (metals) anion – negatively charged ion (nonmetals)

Common ions (Fig. 2.20, p. 56)

© 2009, Prentice-Hall, Inc.

Ionic BondsIonic compounds (such as NaCl) are generally formed between metals and nonmetals.

Predicting ionic chargesAtoms will lose or gain electrons to attain a noble gas configuration.

P3–

Ionic Compounds

Ionic compounds – consist of ions; form crystal lattices

+ and − charges balance Formula unit – ratio of

cation to anion

2.8 Naming Inorganic Compounds

1957 IUPAC (Int’l Union of Pure and Applied Chemistry) – devised systematic rules for naming compounds

Binary compounds – consist of two different elements

Don’t capitalize compound or element names.

Ionic compounds - cations

Cations (Table 2.4, p. 60)1. Single metal, single charge

Na+, sodium ion Al3+, aluminum ion

2. Single metal, multiple charges Cr2+, chromium(II) ion Cr4+, chromium(IV) ion

3. Polyatomic ions

Ionic compounds - anions

Anions (Table 2.5, p. 63)1. Monoatomic, -ide ending

Cl-, chloride ion O2-, oxide ion

2. Oxyanions NO3

-, nitrate ion NO2

-, nitrite ion

3. H+ + oxyanion

P3– Phosphide ion PO33- Phosphite ion

NO2– Nitrite ion

Polyatomic Ions to Memorize(formula & charge)

Name Formula Name Formula

acetate ion C2H3O2- hypochlorite ion ClO-

ammonium ion NH4+ nitrate ion NO3

-

carbonate ion CO32- nitrite ion NO2

-

chlorate ion ClO3- perchlorate ion ClO4

-

chlorite ion ClO2- permanganate ion MnO4

-

chromate ion CrO42- phosphate ion PO4

3-

cyanide ion CN- phosphite ion PO33-

hydroxide ion OH- sulfate ion SO42-

sulfite ion SO32-

Ionic compounds

Cation first, anion second Charges (+ and -) must balance

Acids (p. 64)

Acid – substance which produces a H+ when dissolved in water

If anion ends in ____, acid ends with ____. -ide -ic -ate -ic -ite -ous

Molecular Compounds Name as written in

formula. Prefixes denote

number of each atom. Exceptions:

H2O water

NH3 ammonia

CH4 methane

2.9 Some Simple Organic Compounds Hydrocarbon – contain only C and H Alkanes – saturated hydrocarbons with only

C−C single bonds Alkane derivatives:

−OH alcohol−COOH carboxylic acid−COOC− ester−COC− ketone

Organic compounds, cont.

Unsaturated hydrocarbons: Alkenes – contain at least one C=C double bond Alkynes – contain at least one C≡C triple bond