Lecture 2 Atoms, Molecules, and Ions
Text: Sections B, C, D
(B)
2.1 The Atomic Theory of Matter
Democritus – tiny indivisible particles called “atomos”
Dalton
1. Each element is composed of extremely small particles called atoms.
2. All atoms of a given element are identical; the atoms of different elements are different and have different properties (including masses).
3. Atoms of an element are not changed into different types of atoms by chemical reactions; atoms are neither created nor destroyed in chemical reactions.
4. Compounds are formed when atoms of more than one element combine; a given compound always has the same relative number and kind of atoms.
Atoms are the basic building blocks of matter. They are the smallest particles of an element that retain the chemical identity of the element.
Law of constant composition: in a given compound the relative numbers and kinds of atoms are constant.
The law of conservation of mass: the total mass of material present after a chemical reaction is the same as the total mass before the reaction.
Law of multiple proportions: if two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in a ratio of small whole numbers.
Example: H2O “water” H2O2 “hydrogen peroxide”
2 g hydrogen : 16 g oxygen > 2 g hydrogen : 32 g oxygen
ratio = 16 g oxygen / 32 g oxygen = 1 / 2
2.2 The Discovery of Atomic Structure
Atoms are composed of subatomic particles.
There are two types of electrical charge, positive (+) and negative (-). (Uncharged is “neutral”).
Law of Electrostatic Attraction: like charges repel one another, unlike charges attract.
Cathode rays and electrons
Thomson 1.76x108 coulombs per gram
Millikan oil-drop experiment: charge of an electron 1.6x10-19 C
Mass Spectrometer: If a molecule loses or gains an electron, it becomes charged. The charged particle is attracted to a magnet. The greater the mass of the particle, the less it is attracted. Molecular mass can be determined.
Radioactivity: spontaneous emission of radiation
Three types: a (alpha particles) 2+ charge
b (beta particles) 1- charge
g (gamma rays) no charge
Rutherford found that the alpha particles combined with electrons to form helium atoms.
The Nuclear Atom
Thomson’s “plum pudding” model, a positively charged sphere with electrons embedded in it.
Rutherford’s experiments:
Rutherford bombarded thin gold foil with alpha particles. He found that most of them passed through, some of them were slightly deflected, and some were deflected at high angles.
From these results he deduced that most of the mass of the atom is in a small, positively charged nucleus and that the atom was mostly empty space where electrons moved.
2.3 The Modern View of Atomic Structure
Subatomic particles: proton, neutron, and electron.
The proton and neutron are found in the nucleus.
The charge of an electron is -1.602 x 10-19 C.
The charge on a proton is +1.602 x 10-19 C.
The electronic charge is 1.602 x 10-19 C.
Atoms have equal number of electrons and protons thus they are neutral.
Electrons are attracted to protons in the nucleus by the force that exists between particles of opposite electrical charge.
The mass of atoms is small;
the heaviest atom is about 4x10-22 g.
When discussing the mass of atoms we will use the atomic mass unit (amu). The amu is 1.66054 x10-24 g.
Table 2.1
Particle / Charge / Mass (amu)Proton / Positive (1+) / 1.0073
Neutron / None (neutral) / 1.0087
Electron / Negative (1-) / 5.486 x 10-4
Most of the mass is in the nucleus (protons and neutrons).
The size of atoms is small! Atomic diameters are on the order of 1x10-10 m to 5x10-10 m (100 – 500 pm). A convenient unit (not SI) to express atomic diameters is the angstrom (Å). 1 Å = 10-10 m. Atoms are 1 – 5 Å in diameter.
Isotopes, Atomic Numbers and Mass Numbers
All atoms of an element have the same number of protons in the nucleus.
The number of protons determines the type of atom.
In an atom the number of electrons equals the number of protons.
Atoms of the same element that differ in the number of neutrons, and mass, are called isotopes.
Atomic number (subscript): number or protons in nucleus. Often omitted since the atomic symbol indicates this.
Mass number (superscript): total number of protons and neutrons.
14C “carbon 14” 6 protons (carbon) 8 neutrons
An atom of a specific isotope is called a nuclide.
Difference in elements is due to the difference in the number of subatomic particles.
Relative abundance from mass spectrometer.
2.4 The Periodic Table
The elements in a column (up and down) are known as a group.
Several ways to label groups.
North American convention: A and B designations.
Elements belonging to the same group exhibit similar chemical and physical properties.
“Halogens” are in group VII (VIIA or 17), “alkali metals”, “noble gases”…
All elements on the left side are metallic elements.
The metals are separated from the nonmetals by a steplike line that runs from boron (B) to astatine to (At).
Elements that lie along the line that separates the metal from nonmetals have intermediate properties and are often referred to as metalloids.
(C)
2.5 Molecules and Molecular Compounds
Macroscopic View: Compound: substance consisting of two or more elements in definite proportions
Microscopic View: A molecule is an assembly of two or more atoms tightly bound together.
Most common compounds are either
1) molecular compounds or 2) ionic compounds (salts)
Molecules and Chemical Formulas.
Chemical formulas give the composition substances.
The subscripts in the formula tell us the number of that type of atom present in the molecule.
Example: O2 “Oxygen” a molecule with two oxygen atoms
H2O “water” two hydrogen atoms one oxygen atom
Molecules containing two atoms are called diatomic.
Elements that occur naturally as diatomic molecules are: N2, O2, H2, F2, Cl2, Br2, and I2. (know these!)
When we speak of these elements we are referring to the diatomic form listed above.
Molecular compounds are compounds that are composed of molecules.
Composition of compounds is given by their chemical formulas.
Most molecular substances that we will encounter contain only nonmetals. (right side of periodic table)
Most ionic compounds that we will encounter contain both metal and nonmetal (vide infra).
“Metals” are substances containing one or more metallic elements.
Molecular and Empirical Formulas
Molecular formulas indicate the actual number and types of atoms in a molecule.
Empirical formulas give only the relative number of atoms of each type. The subscripts are always the smallest whole number ratio.
Example: Mol. Formula Empirical Formula
H2O2 HO
C2H4 CH2
H2O H2O
Picturing Molecules
The structural formula of a substance shows which atoms are attached to which within the molecule ( a flat, 2D picture).
A “perspective drawing” uses wedges to give 3D character.
2.6 Ions and Ionic Compounds
Addition or removal of electrons from a neutral atom results in the formation of a charged particle called an ion.
An ion with positive charge is called a cation.
A negatively charged ion is called an anion.
The net charge is represented by a superscript.
Superscripts +, 2+, and 3+ mean a net charge resulting from the loss of one, two, or three electrons.
Superscripts -, 2-, and 3- mean a net charge resulting from the gain of one, two, or three electrons.
In general, metal atoms lose electrons; nonmetals tend to gain electrons.
Predicting Ionic Charges
Many atoms gain or lose electrons so as to end up with as many electrons as the closest noble gas.
Group 1A atoms form 1+ ions
Group 2A atoms form 2+ ions
Group 7A atoms form 1- ions
Group 6A atoms form 2– ions
Ionic Compounds
Ionic compounds contain positively charged ions and negatively charged ions. (salts)
Generally, cations are metal ions and anions are nonmetal ions.
Ionic compounds are generally combinations of metals and nonmetals.
(Polyatomic ions: NO3–, SO42-, CO32–, NH4+ non-metals)
Molecular compounds are generally composed of nonmetals only.
Only empirical formulas can be written for most ionic compounds. These will be given such that the total positive charge equals the total negative charge.
NaCl = 1x(Na+) 1x(Cl–) = 1(+1) + 1(-1) = 0
BaCl2 = 1x(Ba2+) 2x(Cl–) = 1(+2) + 2(-1) = 0
Mg3N2 = 3x(Mg2+) 2x(N3–) = 3(+2) + 2(-3) = 0
(D)
2.7 Naming Inorganic Compounds
Names and Formulas of Ionic Compounds
1. Positive Ions (Cations)
a. Cations formed from metal atoms have the same name as the metal.
b. If a metal can form cations of differing charges, the positive charge is given by a Roman numeral in parentheses following the name of the metal: (for transition metals, but NOT groups I, II, or III because there is only one charge for each of them)!!!
Older names used –ous or –ic instead of numerals (see Appendix 3C)
c. Cations formed from nonmetals atoms usually have names that end in -ium:
2. Negative Ions (Anions)
a. Monatomic (one-atom) anions have names formed by dropping the ending of the name of the element and adding the ending -ide.
b. Polyatomic (many-atom) anions containing oxygen have names ending in -ate or -ite. (and sometimes prefixes of hypo- or per-)
c. Anions derived by adding H+ to an anion are named by adding as a prefix the word hydrogen or dihydrogen, as appropriate.
HCO3- “hydrogen carbonate”, CO32- “carbonate”
CO32- plus H+ equals HCO3-
(-2) plus (+1) equals (-1)
(Bicarbonate is the old term for hydrogen carbonate.)
3. Ionic Compound Names
Names of ionic compounds are the cation name followed by the anion name:
BaBr2 barium bromide
Al(NO3)3 aluminum nitrate
Cu(ClO4)2 copper (II) perchlorate (or cupric perchlorate)
The key to naming the metal charge is knowing the charge on the anion. Know Table D.1 !!!
-1) F-, Cl-, Br-, I-, OH-, CN-, NO2-, NO3-,ClOx-, CH3CO2-
-2) O2-, CO32-, SO32-, SO42-,
-3) PO43-
4. Hydrates of Ionic Compounds
Water can enter crystal structure in definite proportions to form distinct compounds: hydrides
CuSO4 “anhydrous copper(II) sulfate” a light blue powder CuSO4·H2O “copper(II) sulfate pentahydrate” dark blue crystal
Names and Formulas of Acids
1. Acids Based on Anions Whose Name End in -ide. Anions whose names end in -ide have associated acids that have the hydro- prefix and an -ic ending as in the following examples.
From Table D.1 KNOW: names, symbols, charges
Important detail:
Anion Acid(disolved in water) (Gas Molecule)
Cl– (chloride) HCl (hydrochloric acid) hydrogen chloride
S2- (sulfide) H2S (hydrosulfuric acid) hydrogen sulfide
2. Acids Based on Anions Whose Names End in -ate or
-ite. Anions whose names end in -ate have associated acids with an -ic ending, whereas anions whose names end in -ite have acids with an -ous ending.
anion / acidHypochlorite, ClO– / Hypochlorous, HClO
Chlorite, ClO2– / Chlorous, HClO2
Chlorate, ClO3– / Chloric, HClO3
Perchlorate, ClO4– / Perchloric, HClO4
“Halide” is the general name for anions of halogens.
Just as there is chlorate, there are bromate and iodate ions, etc.
Names and Formulas of Binary Molecular Compounds
1. The name of the element farthest to the left in the periodic table is usually written first.
2. If both elements are in the same group in the periodic table, the lower one is named first.
3. The name of the second element is given an -ide ending.
4. Greek prefixes are used to indicate the number of atoms of each element. The prefix mono- is never used with the first element. When the prefix ends in a or o and the name of the second element begins with a vowel (such as oxide), the a or o is often dropped. (Table D.2)
Required Skills:
1) Give the name of the compound from its formula.
2) Give the formula of the compound from its name.
3) Know Greek prefixes for numbers 1 – 12. (Table D.2)
AND…
4) Know names of hydrocarbons and their groups
Ch 2 Atoms, Molecules, and Ions 23