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Chapter 2 Page 1 CHAPTER TWO: ATOMS, MOLECULES, AND IONS Part one: Atomic Theory and Atomic Structure A. Atomic Theory of Matter. (Section 2.1) 1. 400 B.C. - Democritus: matter is not continuous but is composed of unimaginably tiny, discrete, indivisible particles he called “atoms.” 2. 1808 A.D. - John Dalton published first atomic theory: a. All matter is composed of extremely small indivisible particles called atoms, that retain their identity during chemical reactions. b. All atoms of a given element have identical properties, which differ from those of other elements. c. Atoms cannot be created, destroyed, or transformed into atoms of another element. (except by nuclear reactions) d. Compounds are formed when atoms of different elements combine with each other in small whole-number ratios. e. The relative numbers and kinds of atoms are constant in a given compound. 3. This theory explained Laws of Conservation of Matter and Laws of Definite Proportions. B. The Law of Definite Proportions. (applies to compounds) 1. A compound is a pure substance consisting of two or more different elements in a fixed ratio. 2. The Law of Definite Proportions states that: Different samples of any pure compound contain the same elements in the same proportion by mass. 3. Example: Water is always found to have the definite proportion 88.9% Oxygen and 11.1% Hydrogen by mass. Why is this? a. Because water is composed of particles in which 1 atom of O is attached to 2 atoms of H. b. Yet atoms of Oxygen weigh 16 times as much as an atom of Hydrogen.

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Page 1: Atoms, Molecules and Ions

Chapter 2 Page 1

CHAPTER TWO: ATOMS, MOLECULES, AND IONS

Part one: Atomic Theory and Atomic Structure

A. Atomic Theory of Matter. (Section 2.1) 1. 400 B.C. - Democritus: matter is not continuous but is composed of unimaginably

tiny, discrete, indivisible particles he called “atoms.” 2. 1808 A.D. - John Dalton published first atomic theory:

a. All matter is composed of extremely small indivisible particles called atoms, that retain their identity during chemical reactions.

b. All atoms of a given element have identical properties, which differ from those

of other elements.

c. Atoms cannot be created, destroyed, or transformed into atoms of another element. (except by nuclear reactions)

d. Compounds are formed when atoms of different elements combine with each

other in small whole-number ratios.

e. The relative numbers and kinds of atoms are constant in a given compound. 3. This theory explained Laws of Conservation of Matter and Laws of Definite

Proportions.

B. The Law of Definite Proportions. (applies to compounds) 1. A compound is a pure substance consisting of two or more different elements in a

fixed ratio. 2. The Law of Definite Proportions states that:

Different samples of any pure compound contain the same elements in the same proportion by mass.

3. Example: Water is always found to have the definite proportion 88.9% Oxygen and

11.1% Hydrogen by mass. Why is this?

a. Because water is composed of particles in which 1 atom of O is attached to 2 atoms of H.

b. Yet atoms of Oxygen weigh 16 times as much as an atom of Hydrogen.

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c. Thus the 2:1 ratio of H:O by atoms corresponds to a 2:16 ratio by mass. There are 2 parts H and 16 parts O by mass.

% O = 16/18 x 100 = 88.9% O by mass % H = 2/18 x 100 = 11.1% H by mass

C. Law of Multiple Proportions.

1. When two elements form more than one compound with each other, the masses of one

element in these compounds for a fixed mass of the other element are in ratios of small whole numbers.

2. Example. H and O combine to form two compounds: H2O and H2O2 Amount of Oxygen per gram of H in 1st compound = whole # ratio Amount of Oxygen per gram of H in 2nd compound 8.0 grams = 1/2 16.0 grams

That this was true confirmed that compounds form by atoms combined in fixed whole number ratios.

D. Structure of the Atom.

1. Atoms are not the smallest particles. They are composed of three fundamental

particles:

a. electrons (e-) discovered by J. J. Thomson 1897 (cathode ray tube experiment)

b. protons (p or p+) discovered by Goldstein in cathode ray tube

c. neutrons (n or n0) discovered by Chadwick in 1932

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E. The Nuclear Atom Model. (Rutherford) (Section 2.2) 1. Early model of atomic structure. (Thomson)

2. Disproved by Rutherford’s Gold Foil Experiment (1910).

a. α particles very dense, very fast. b. Expected all to pass through with minor deflections from hitting Thomson-like

atoms. c. Actual result:

1.) Nearly all particles passed through the gold as if through empty space. 2.) Amazingly, a few rebounded as if hitting very dense ⊕ charges in the foil

(See Figure 2.8)

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3. Conclusion: atoms are mostly empty space with a pointlike center of ⊕ charge

having virtually all the mass.

F. Atomic Masses (also called “Atomic Weights”) (AM or AW). (Section 2.4)

1. 19th century chemists systematized large body of data establishing mass combining

ratios of elements. Here is an example of how it worked: suppose they had the following initial data: Mg and O combine ~3 to 2 by mass (to make oxide of Mg) H and O 1 to 8 (to make water) H and C 1 to 3 (to make marsh gas). O and C 8 to 3 (to make main oxide of carbon) 2. They then deduced a scale of relative atomic masses (traditionally called atomic

weights). 3. Originally H was established as lightest element and given a value of 1 on the

relative scale; later eventually chose a relative mass scale based on Carbon (easier to work with), and assigned C an Atomic Mass =12 because it was nearly 12 on the Hydrogen scale.

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4. Modern Atomic Masses are based on atomic mass units, assigning the Carbon-12 isotope of C to have mass of exactly 12 amu.

1 amu ≡ 1/12 (mass of 12C atom) OR:

mass of one atom 12C ≡ 12 amu 5. On this scale, Atomic Mass of H is no longer 1 but 1.00794 amu. 6. Atomic Masses are given on Periodic Chart under element symbol. 7. Atomic Mass is an element’s average mass of an atom in amu (averaged over the

stable isotopes of that element). 8. Periodic Table Symbol for Element:

12 Mg 24.305

Here 12 is the Atomic Number, Mg is the elemental symbol, and 24.305 is the Atomic Mass.

9. Atomic Number (Z) = number of protons in the nucleus. (Z determines which element)

10. Atomic Mass = mass of atoms of the element on a relative scale (amu’s).

Future definition = number of grams of the element in one mole.

G. Atomic Number. (Section 2.3) 1. Moseley (1913) - xray experiments showed:

Each element differs from preceding element on the chart by having one more

positive charge in its nucleus.

2. Every nucleus has integer # of protons equal to # of e- (in a neutral atom):

H has one proton / He has 2 protons / Li has 3 protons / etc.

3. Atomic number = Z = number of protons in the nucleus of an atom (determines that atom’s identity).

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H. Neutrons. 1. Chadwick (1932) discovered neutrons by bombarding Be with α-particles. 2. Neutrons = neutral particles in the nucleus having about the same mass as protons. Therefore: nuclei contain neutrons as well as protons.

I. Mass Number and Isotopes. (Section 2.4)

1. Isotopes of a given element contain the same number of protons (Z) but differ in

number of neutrons in the nucleus. 2. Mass number = sum of protons and neutrons = atomic number (Z) + neutron number Hydrogen: mass number = 1 Deuterium: mass number = 2 Tritium: mass number = 3 3. Nuclide symbol:

Example - two stable isotopes of Cl: 17

35Cl 1737Cl

Z 17 17

mass number 35 37

# of neutrons 18 20

% natural abund. 75.77 24.23

mass (amu) 34.969 36.966

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4. Atomic Mass (AM) of an element is actually an average over the different stable isotopes.

Chlorine 35Cl 37Cl mass = 34.969 amu 36.966 amu % 75.77 24.23 Average mass of Cl = 75.77 x 34.969 + 24.23 x 36.966 100 = 35.453 amu = 35.453 g/mol

Average mass 35.45 amu = Atomic Mass of Cl

J. Mass Spectrometry (M.S.) and Isotopic Abundance. (Section 2.4) 1. M.S. measures (charge/mass) ratio of charged particles. 2. Gaseous sample bombarded with high-energy e-, and some e- are knocked off the gas

molecules creating positive ions. 3. These are focused into a beam and passed through magnetic field. 4. Field deflects ions by an angle based on:

a. voltage of field focusing the positive beam. b. magnetic field strength. c. masses of positive particles. d. charges of positive particles.

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Figure 2.13. Mass spectrum of neon (1+ ions only). Neon consists of three isotopes, of which neon-20 is by far the most abundant (90.48%). The mass of that isotope, to five decimal places, is 19.99244 amu on the carbon-12 scale. The number by each peak corresponds to the fraction of all Ne+ ions represented by the isotope with that mass.

K. The Periodic Table. (Section 2.5) 1. Based on the observation called the Periodic Law:

The properties of the elements are periodic functions of their atomic numbers. 2. Examples:

a. Elements Z = 2, 10, 18 have similar properties (He, Ne, Ar are chemically inert

gases.) b. Elements Z = 3, 11, 19 have similar properties (Li, Na, K are chemically active

metals combining with oxygen to form X2O compounds.) 3. First noted by Mendeleev and Meyer (1869). Arranged the 60 known elements in

increasing order of atomic weight. (Atomic number was unknown concept then.) 4. Periodic Law works because Z also equals number of electrons in the neutral atom,

and number of e- determines properties. 5. Vertical columns are groups, horizontal rows are called periods. 6. Groups most strongly correlated with one another are:

a. Alkali metals: Group IA - Li, Na, K, Rb, Cs. (alkaline means basic) b. Alkaline earth metals: Group IIA - Be, Mg, Ca, Sr, Ba.

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c. Halogens: Group VIIA - F, Cl, Br, I. (halogen = “salt formers”) d. Noble gases: Group O - He, Ne, Ar, Kr, Xe, Rn. (inert gases)

7. Broader categorization of elements into metals and nonmetals:

Figure: Trends in metallic character of A group elements. 8. Physical properties of metals and nonmetals: 9. Chemical properties of metals and nonmetals.

Slide 9 p.127

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10. Metalloids form a sort of boundary between metals and nonmetals.

Some act as semiconductors = insulators at low T and conductors at high T.

Part Two - Chemical Substances: Formulas and Names A. Chemical Formulas (Section 2.6)

1. Examples: O2 H2O CH4 NaCl NH3 2. Shorthand for chemical composition of a substance. 3. Chemical formula shows the elements present in a substance and the ratio in which

atoms of those elements are combined. 4. Some substances occur in molecular form, while others occur as ionic compounds.

You'll need to learn to tell the difference.

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5. The chemical formula has a different meaning in those 2 cases. Examples: a. Water occurs as molecules = H2O. It is a molecular substance. Not only does

the compound occur with a 2:1 ratio of H to O atoms, but the substance comes in particles H2O.

Model of H2O. b. The salt calcium fluoride is an ionic substance with formula CaF2, meaning that,

while the atoms of Ca and F are in a 1-to-2 ratio, it does not exist as molecules of CaF2, but as a crystal lattice containing 1 Ca for every 2 F atoms.

Ionic lattice 6. Chemical formula for an ionic compound merely represents atom ratios. 7. Chemical formulas for molecular substances actually represent the make-up of the

molecules themselves.

8. Various representations of molecular substances (See Fig. 2.18):

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9. Molecular models are required to represent 3-D geometric arrangements. Use HyperChem to show molecules

a. ball and stick model. b. space-filling model.

10. The most stable forms of various elements in pure form:

a. Noble gases - He, Ne, Ar, ... stable as individual atoms. (i.e., monatomic gases)

b. Several common elements exist in most stable form as diatomic molecules - H2, N2, O2, F2, Cl2, Br2, I2, ...

c. Some elements exist as polyatomic molecules:

phosphorus exists as P4 (white phosphorus) sulfur as S8 ring. d. Most elements in pure form exist not as molecules at all, but in large repeating

arrays (crystalline solids)

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B. Ionic Compounds. (Section 2.6) 1. Don’t exist as discrete molecules, but in large crystalline arrays of ions. 2. Ions are formed when electrons are:

lost from metal atoms (like Na, K, Mg, Ca,...) gained by nonmetal atoms. (like Cl, Br, I, O, F,...)

Na+ Sodium cation (Na missing 1 e-) Cl- Chloride anion (Cl with 1 e- added) NaCl is not a molecular entity. Exists in a lattice (See Fig. 2.21): 3. Above was an example of monatomic ions. Memorize the others from Table 2.3 and

2.4.

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4. There are several important molecular ions or polyatomic ions (See Table 2.5):

a. NH4

+ ammonium ion b. SO4

2- sulfate ion c. NO3

- nitrate ion d. NO2

- nitrite ion e. CO3

2- carbonate ion f. OH- hydroxide ion g. PO4

3- phosphate ion

5. These can combine with other ions to form an ionic compound: a. NH4Cl ammonium chloride b. Na2SO4 sodium sulfate

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6. Knowing the charges of various monatomic and polyatomic ions, you should be able to figure out some chemical formulas for thousands of ionic compounds (see Example 2.3 in text): a. potassium bromide K1+ Br1- not K1Br1 but KBr b. silver sulfide Ag1+ S2- Ag2S c. iron(III) sulfate Fe3+ SO4

2- Fe2(SO4)3 d. ammonium phosphate NH4

+ PO43- (NH4)3PO4

e. calcium carbonate marble, chalk, seashells Ca2+ CO3

2- not Ca2(CO3)2 but CaCO3 f. aluminum phosphate Al3+ PO4

3- AlPO4

C. Organic Compounds (Section 2.7)

1. These are molecular compounds that contain carbon combined with other elements

such as hydrogen, oxygen and nitrogen.

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Part Three: Naming Simple Inorganic Compounds (Section 2.8) A. Naming Binary Compounds.

1. Consist of two elements; either ionic or molecular. 2. Name more metallic element 1st and less metallic element 2nd. 3. Less metallic element named by adding “-ide” suffix to element’s stem name. 4. Example: Some binary ionic compounds containing metals that exhibit only one

charged state.

Formula Name KBr potassium bromide CaCl2 calcium chloride NaH sodium hydride

5. Example: Binary ionic compounds with metals that exhibit more than one stable

charge; the charge of the metal is indicated by Roman numeral in parentheses.

Formula Cation Charge

Name

Cu2O +1 copper(I) oxide CuF2 +2 copper(II) fluoride FeS +2 iron(II) sulfide Fe2O3 +3 iron(III) oxide

6. Older method used “-ous” and “-ic” suffixes to indicate lower and higher ox#’s,

respectively.

Formula Cation Charge

Name

CuCl +1 cuprous chloride CuCl2 +2 cupric chloride FeO +2 ferrous oxide FeBr3 +3 ferric bromide

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7. Pseudobinary ionic compounds: one or more of the ions consist of more than one element but behave as simple ions.

Example: hydroxide ion, OH- ; the cyanide ion, CN- ; thiocyanate ion, SCN- . Name of the anion ends in “-ide.” NH4

+, is the common cation that behaves like a simple metal cation.

Formula Name

NH4I ammonium iodide Ca(CN)2 calcium cyanide NaOH sodium hydroxide

8. Binary molecular compounds: involve two nonmetals bonded together. Elemental

proportions are indicated by using a prefix system for both elements (See Table 2.6).

Formula Name Formula Name SO2 sulfur dioxide Cl2O7 dichlorine heptoxide SO3 sulfur trioxide CS2 carbon disulfide N2O4 dinitrogen tetroxide As4O6 tetraarsenic hexoxide

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9. Binary acids dissolved in water. When pure, named as typical binary compounds. Their aqueous solutions are named with the prefix “hydro-” and the suffix “-ic” followed by the word “acid.”

Formula Name of Compound Name of Aqueous Solution

HCl hydrogen chloride hydrochloric acid, HCl(aq) HF hydrogen fluoride hydrofluoric acid, HF(aq) H2S hydrogen sulfide hydrosulfuric acid, H2S(aq) HCN hydrogen cyanide hydrocyanic acid, HCN(aq)

B. Naming Ternary Acids and Their Salts.

1. Ternary compound consists of three elements. 2. Ternary acids (oxoacids) are compounds of hydrogen, oxygen, and (usually) a

nonmetal. (e.g. H2SO4)

3. Nonmetals that exhibit more than one stable charge form more than one ternary acid,

differing in number of oxygen atoms. 4. Most common ternary acid is given the “-ic” name. See following.

Examples:

Nitric acid

Sulfuric acid

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Phosphoric acid

Chloric, Bromic, and Iodic acids

Carbonic acid 5. Acids containing one fewer oxygen atom per central atom are named “-ous.”

Formula Name Formula Name H2SO3 sulfurous acid HClO2 chlorous acid HNO2 nitrous acid

6. Acids that have fewer O atom than the “-ous” acids are named using the prefix

“hypo-” and the suffix “-ous.”

Formula Name HClO hypochlorous acid H3PO2 hypophosphorous acid

7. Acids containing one more oxygen atom per central nonmetal atom than the normal

“-ic acid” are named “per” “ic” acids.

Formula Name HClO4 perchloric acid HBrO4 perbromic acid HIO4 periodic acid

8. The oxoacids of chlorine follow as example:

Formula Ox. No. of Cl Name HClO +1 hypochlorous acid HClO2 +3 chlorous acid HClO3 +5 chloric acid HClO4 +7 perchloric acid

9. Ternary salts: for example, KClO3 - potassium chlorate.

a. Anion derived from “-ic acid” is “-ate.” b. Anion derived from “-ous acid” is “-ite.”

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c. The “-per-” and “hypo-” prefixes are retained. Examples: HClO3 KClO3 ClO3

- chloric acid potassium chlorate chlorate ion

“ic acid” → “ate” HClO2 NaClO2 ClO2

- chlorous acid sodium chlorite chlorite ion

“ous acid” → “ite” HClO NH4ClO ClO- hypochlorous acid ammonium hypochlorite hypochlorite ion HClO4 KClO4 ClO4

- perchloric acid potassium perchlorate perchlorate ion 10. Summary Chart: naming ternary acids and their anions. The stem (XXX) represents

the stem of the name, e.g., “nitr,” “sulfur,” or “chlor.”

11. Examples:

Formula Name (NH4)2SO4 ammonium sulfate KNO3 potassium nitrate Ca(NO2)2 calcium nitrite LiClO4 lithium perchlorate FePO4 iron(III) phosphate NaClO sodium hypochlorite

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12. Ternary acids salts in which one or more acidic hydrogen atoms remain: named with the word “hydrogen” or “dihydrogen” inserted after the cation.

Formula Name

NaHSO4 sodium hydrogen sulfate NaHSO3 sodium hydrogen sulfite KH2PO4 potassium dihydrogen phosphate K2HPO4 potassium hydrogen phosphate NaHCO3 sodium hydrogen carbonate

(sodium bicarbonate)

Part Four: Chemical Reactions: Equations A. Writing Chemical Equations. (Section 2.9)

1. Shorthand for a chemical reaction, showing:

a. substances reacting. (reactants) b. substances formed. (products) c. relative amounts involved. (balancing coefficients)

2. Example: combustion of natural gas, methane. (CH4)

CH4 + 2 O2 → CO2 + 2 H2O a. reactants - CH4, O2 b. products - CO2, H2O c. Reads: 1 molecule CH4 reacts with 2 molecules O2 producing 1 molecule CO2

and 2 molecules H2O. d. Note - both sides of equation contain: 1 atom C 4 atoms H 4 atoms O Therefore, equation is balanced.

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e. Note - smallest whole-number coefficients are typically used to balance:

1 : 2 : 1 : 2 f. Note - mole interpretation is also valid. 1 mole CH4 reacts with 2 moles O2 producing 1 mole CO2 and 2 moles H2O. g. Species formulas in the equation must describe them as they exist. Example: this is improper.

CH4 + 4 O → CO2 + 2 H2O This balances, but oxygen not shown as diatomic molecule.

B. Balancing Chemical Equations. (Section 2.10)

1. Problem: Write down and balance the chemical equation for the combustion of propane, C3H8, in the presence of abundant Oxygen.

__ C3H8 + __ O2 → __ CO2 + __ H2O

a. First balance the elements that appear in only one species on both sides of the

equation (C and H) b. Then complete the balance of O

C3H8 + 5 O2 → 3 CO2 + 4 H2O 2. All combustion reactions can be done this way.

a. Try combustion of ethanol, C2H6O.

__ C2H6O + __ O2 → __ CO2 + __ H2O b. Incomplete combustion of methane: (incomplete combustion refers to

production of CO rather than CO2)

__ CH4 + __ O2 → __ CO + __ H2O