Chemistry: Atoms and Ions - Study Notes
Chapter 2: Atoms and Ions
Early Ideas in Atomic Theory
Dalton's Atomic Theory (Five Postulates)
Matter is composed of exceedingly small particles called atoms.
An atom is the smallest unit of an element that can participate in a chemical change.
An element consists of only one type of atom.
Each atom has a mass characteristic of the element.
The mass is the same for all atoms of that element.
Atoms of one element differ in properties from atoms of all other elements.
A compound consists of atoms of two or more elements combined in a small, whole-number ratio.
In a given compound, the number of atoms of each of its elements are always present in the same ratio.
Atoms are neither created nor destroyed during a chemical change.
Instead, they rearrange to yield different types of matter.
Law of Conservation of Matter
Dalton's atomic theory provides a microscopic explanation for macroscopic properties of matter.
Principle: If atoms are neither created nor destroyed during a chemical change, then the total mass of matter present when matter changes from one type to another will remain constant.
Law of Definite Proportions (Law of Constant Composition)
Principle: All samples of a pure compound contain the same elements in the same proportion by mass.
Example (Isooctane):
Sample A: Carbon = 14.82 \text{ g}, Hydrogen = 2.78 \text{ g} (Mass Ratio consistent)
Sample B: Carbon = 22.33 \text{ g}, Hydrogen = 4.19 \text{ g} (Mass Ratio consistent)
Sample C: Carbon = 19.40 \text{ g}, Hydrogen = 3.64 \text{ g} (Mass Ratio consistent)
Evolution of Atomic Theory
Important Discoveries of the 20th Century
Isotopes: Atoms of the same element that differ in mass.
Neutrons: Uncharged, subatomic particles found in the nucleus with a mass approximately the same as that of protons.
Atomic Structure and Symbolism
Atomic Composition
Nucleus: Contains the majority of an atom's mass.
Composed of protons and neutrons.
Diameter of a nucleus is approximately 10^{-15} \text{ m}.
Electrons: Occupy almost all of an atom's volume.
Much lighter than protons and neutrons.
Diameter of an atom is approximately 10^{-10} \text{ m} (100,000 times larger than the nucleus).
Analogy: If an atom were a football stadium, the nucleus would be the size of a single blueberry.
Units for Subatomic Particles
Atoms and subatomic particles are exceedingly small.
Carbon atom weighs less than 2 \times 10^{-23} \text{ g}.
Electrons have a charge of less than 2 \times 10^{-19} \text{ C}.
Atomic Mass Unit (amu):
1 \text{ amu} = 1.6605 \times 10^{-24} \text{ g}.
Mass of a carbon-12 atom = 12 \text{ amu}.
Fundamental Unit of Charge (e):
e = 1.602 \times 10^{-19} \text{ C} (equivalent to the magnitude of charge on one electron or proton).
Properties of Subatomic Particles
Particle | Mass (amu) | Charge |
|---|---|---|
Proton | 1.0073 | +1 |
Neutron | 1.0087 | 0 |
Electron | 0.00055 | -1 |
Atomic Number (Z)
The number of protons in the nucleus of an atom.
Defining trait of an element: Its value determines the identity of the atom.
For example, any atom with six protons is carbon (Z = 6), regardless of neutron or electron count.
Neutral Atoms
Must contain the same number of positive and negative charges.
Number of protons equals the number of electrons.
Therefore, in a neutral atom, Z also indicates the number of electrons.
Mass Number (A)
The total number of protons and neutrons in an atom.
Calculation of Neutrons: Number of neutrons = A - Z.
Atomic number (Z) = number of protons.
Mass number (A) = number of protons + number of neutrons.
Ions
Atoms (or molecules) that are electrically charged because the number of protons and electrons are NOT equal.
Charge of an atom = number of protons - number of electrons.
Formed by gaining or losing electrons.
Types of Ions
Anion: An atom that gains one or more electrons, exhibiting a negative charge.
Example: A neutral oxygen atom (Z=8) has 8 electrons. If it gains two electrons, it becomes an anion with a 2- charge (8 - 10 = -2).
Cation: An atom that loses one or more electrons, exhibiting a positive charge.
Example: A neutral sodium atom (Z=11) has 11 electrons. If it loses one electron, it becomes a cation with a 1+ charge (11 - 10 = +1).
Chemical Symbols
An abbreviation used to indicate an element or an atom of an element.
Example: Hg for mercury.
Most symbols have one or two letters; three-letter symbols for elements past Z=112.
Only the first letter is capitalized.
(See Table 2.3 for common elements and their symbols, e.g., Fe from ferrum, Pb from plumbum, Au from aurum).
Isotope Notation
The symbol for a specific isotope places the mass number (A) as a superscript to the left of the element symbol.
The atomic number (Z) is sometimes written as a subscript to the left of the element symbol, but is often omitted because the element symbol already implies Z.
General form: ^{\text{A}}_{\text{Z}}\text{X} or simply ^{\text{A}}\text{X}.
Example: Magnesium isotopes are ^{\text{24}}\text{Mg}, ^{\text{25}}\text{Mg}, and ^{\text{26}}\text{Mg}.
All have 12 protons (Z=12), but differing numbers of neutrons.
Hydrogen Isotopes:
Protium (^{\text{1}}\text{H}): 1 proton, 0 neutrons.
Deuterium (^{\text{2}}\text{H} or ^{\text{2}}\text{D}): 1 proton, 1 neutron.
Tritium (^{\text{3}}\text{H} or ^{\text{3}}\text{T}): 1 proton, 2 neutrons.
Atomic Mass
The atomic mass of a single atom in amu is approximately equal to its mass number (since protons and neutrons each have mass \approx 1 \text{ amu}).
Most elements occur naturally as a mixture of two or more isotopes.
The periodic table lists the weighted, average mass of all isotopes present in a naturally occurring sample of that element.
Calculation Example (Boron):
19.9\% ^{\text{10}}\text{B} with mass 10.0129 \text{ amu}.
80.1\% ^{\text{11}}\text{B} with mass 11.0093 \text{ amu}.
Average Atomic Mass = (0.199 \times 10.0129 \text{ amu}) + (0.801 \times 11.0093 \text{ amu}).
Mass Spectrometry (MS)
An instrument used to experimentally determine the occurrence and natural abundances of isotopes.
Process: Sample is vaporized and exposed to a high-energy electron beam, causing atoms (or molecules) to lose electrons and become electrically charged cations.
These cations are then separated by their mass-to-charge ratio.
Produces a mass spectrum with peaks corresponding to different isotopes.
Chemical Formulas
Types of Chemical Formulas
Molecular Formula: A representation of a molecule or compound.
Consists of chemical symbols for types of atoms.
Subscripts after the symbol indicate the number of each type of atom (only if more than one).
Example: Methane (CH_4).
Structural Formula: Shows the same information as a molecular formula but also explicitly shows how the atoms are connected.
Can be represented as ball-and-stick models or space-filling models.
Elements That Exist as Molecules
Many elements exist as discrete, individual atoms (e.g., noble gases).
Some elements exist as molecules:
Diatomic Molecules: H2, N2, O2, F2, Cl2, Br2, I_2.
Polyatomic Molecules: Elemental sulfur typically exists as S_8.
Distinguishing Symbols:
H: One hydrogen atom.
2H: Two separate hydrogen atoms.
H_2: One hydrogen molecule (two bonded hydrogen atoms).
2H_2: Two hydrogen molecules.
Empirical Formula
Indicates the simplest whole-number ratio of the number of atoms (or ions) in a compound.
Contrast with Molecular Formula: A molecular formula indicates the actual numbers of atoms of each element in a molecule.
Example: Benzene
Molecular formula = C6H6.
Empirical formula = CH.
Example: Acetic Acid
Molecular formula = C2H4O_2.
Empirical formula = CH_2O.
Note: For some compounds, the molecular and empirical formulas are the same (e.g., H2O, TiO2).
Isomers
Compounds with the same chemical formula but different molecular structures and properties.
Structural Isomers: Differ in the bonding arrangement of atoms.
Example: Acetic acid (CH3COOH) and methyl formate (HCOOCH3) both have the molecular formula C2H4O_2.
Spatial Isomers (Stereoisomers): Differ only in the relative orientations of the atoms in space.
Example: Different forms of carvone (spearmint vs. caraway scent).
The Periodic Table
Historical Context
Dimitri Mendeleev (Russia, 1869) and Lothar Meyer (Germany, 1870): Independently recognized a periodic relationship among element properties.
Arranged elements according to increasing atomic mass.
Mendeleev's Contribution: Used his table to predict the existence and properties of unknown elements (e.g., gallium, germanium), which were later discovered, supporting his work.
The Modern Periodic Table
By the 20th century, it was realized that elemental properties are better correlated with atomic numbers.
Periodic Law: The properties of the elements are periodic functions of their atomic numbers.
Organization: Arranges elements in increasing order of their atomic numbers.
Groups (Vertical Columns): 1-18. Elements in the same group have similar chemical properties.
Periods (Horizontal Rows): 1-7. Properties gradually change across a period.
Classifications of Elements
By Physical & Chemical Properties:
Metals: Shiny, malleable, good conductors of heat and electricity (tend to lose electrons).
Nonmetals: Dull, poor conductors of heat and electricity (tend to gain or share electrons).
Metalloids (Semimetals): Possess properties intermediate between metals and nonmetals; conduct heat and electricity moderately well.
By Position on the Table:
Main Group Elements (Representative Elements): Groups 1, 2, 13-18.
Transition Metals: Groups 3-12.
Inner Transition Metals: Two rows at the bottom.
Lanthanides: Top row.
Actinides: Bottom row.
Specific Groups:
Group 1: Alkali Metals (except hydrogen).
Group 2: Alkaline Earth Metals.
Group 15: Pnictogens.
Group 16: Chalcogens.
Group 17: Halogens.
Group 18: Noble Gases (or inert gases).
Molecular and Ionic Compounds
Chemical Reactions and Electron Behavior
In ordinary chemical reactions, the nucleus (and thus element identity) remains unchanged.
Electrons participate by being gained, lost, or shared, leading to the formation of ions or molecules.
Predicting Ion Charge (for Main-Group Elements)
The periodic table guides the prediction of ionic charges for main-group elements.
Metals (Groups 1, 2, 13): Tend to lose electrons to achieve the electron count of the preceding noble gas (forming cations).
Group 1: Lose one electron, form 1+ charge (e.g., Na^+).
Group 2: Lose two electrons, form 2+ charge (e.g., Ca^{2+}).
Group 13 (Al): Lose three electrons, form 3+ charge (e.g., Al^{3+}).
Nonmetals (Groups 15, 16, 17): Tend to gain electrons to achieve the electron count of the next noble gas (forming anions).
Group 17 (Halogens): Gain one electron, form 1- charge (e.g., Br^-).
Group 16 (Chalcogens): Gain two electrons, form 2- charge (e.g., O^{2-}).
Group 15 (Pnictogens): Gain three electrons, form 3- charge (e.g., N^{3-}).
Transition Metals: Less reliable prediction; often form multiple cations with variable charges (e.g., Cu^+ and Cu^{2+}, Fe^{2+} and Fe^{3+}).
Polyatomic Ions
Monatomic Ions: Formed from only one atom (e.g., Na^+, Cl^-, O^{2-}).
Polyatomic Ions (Molecular Ions): Electrically charged molecules; groups of bonded atoms with an overall charge.
Oxyanions: Polyatomic ions that contain one or more oxygen atoms.
Naming Oxyanions (Systematic Approach)
Two Oxyanions from a Nonmetal:
-ate suffix: Used for the ion with the larger number of oxygen atoms (e.g., SO4^{2-} sulfate, NO3^{-} nitrate).
-ite suffix: Used for the ion with the smaller number of oxygen atoms (e.g., SO3^{2-} sulfite, NO2^{-} nitrite).
More than Two Oxyanions from a Nonmetal (using prefixes):
per- (largest number of oxygens): perchlorate (ClO_4^{-}).
-ate (one fewer oxygen): chlorate (ClO_3^{-}).
-ite (two fewer oxygens): chlorite (ClO_2^{-}).
hypo- (smallest number of oxygens): hypochlorite (ClO^{-}).
(See Table 2.5 for a partial list of common polyatomic ions).
Types of Chemical Bonds
Ionic Bond: Results when electrons are transferred between atoms, forming ions.
These are electrostatic forces of attraction between oppositely charged ions.
Covalent Bond: Results when electrons are shared between atoms, forming molecules.
Compounds are classified as ionic or molecular (covalent) based on the types of bonds present.
Ionic Compounds
Typically formed when a metal (readily loses electrons, forming cations) reacts with a nonmetal (readily gains electrons, forming anions).
Contain ions held together by ionic bonds.
Formation Examples:
Na (metal) + Cl (nonmetal) \rightarrow Na^+ + Cl^- \rightarrow NaCl
Ca (metal) + 2 Cl (nonmetal) \rightarrow Ca^{2+} + 2 Cl^- \rightarrow CaCl_2
Properties:
Typically solids at room temperature.
High melting and boiling points.
Nonconductive in solid form.
Conductive in molten (liquid) form or when dissolved in water (due to mobile ions).
Formulas of Ionic Compounds
Ionic compounds are electrically neutral overall.
The formula must show a ratio of ions such that the total positive and negative charges are equal.
These are not molecular formulas; they represent the simplest ratio of ions in the crystalline lattice.
Example 1: Al^{3+} and O^{2-} forms Al2O3.
2 \times (+3) = +6 total positive charge.
3 \times (-2) = -6 total negative charge.
Compounds with Polyatomic Ions: Treat polyatomic ions as discrete units.
Use parentheses to indicate a group of atoms behaving as a unit that appear more than once.
Example 2: Ca^{2+} and PO4^{3-} forms Ca3(PO4)2.
3 \times (+2) = +6 total positive charge.
2 \times (-3) = -6 total negative charge.
Molecular Compounds (Covalent Compounds)
Result when atoms share electrons.
Exist as discrete, neutral molecules.
Usually formed by a combination of two or more nonmetals.
Properties:
Often exist as gases, low-boiling liquids, and low-melting solids (compared to ionic).
Chemical Nomenclature
Overview
Nomenclature: A collection of rules for naming things.
Compounds are identified by both their formula and name.
Focus: Ionic and molecular binary compounds, ionic compounds with polyatomic ions, and acids.
Naming Ionic Compounds
General Rules:
Name the cation first, followed by the name of the anion.
Monoatomic Cation: Simply the name of the element (e.g., sodium).
Monoatomic Anion: Name of the element with its ending replaced by the suffix -ide (e.g., chloride, oxide, sulfide, nitride, phosphide, carbide).
Polyatomic Ion: Given the name of the ion directly (from Table 2.5).
Examples:
NaCl: sodium chloride
Na_2O: sodium oxide
KBr: potassium bromide
Mg3N2: magnesium nitride
KC2H3O_2: potassium acetate
(NH_4)Cl: ammonium chloride
CaSO_4: calcium sulfate
Al2(CO3)_3: aluminum carbonate
Naming Ionic Compounds with Metal Ions of Variable Charge
Primarily transition metals and some main group metals can form multiple cations with different charges.
The charge of the metal ion is specified by a Roman numeral in parentheses after the name of the metal element.
Example: FeCl2 is iron(II) chloride; FeCl3 is iron(III) chloride.
Example: Hg2O is mercury(I) oxide (note: Hg2^{2+} is a specific polyatomic ion); HgO is mercury(II) oxide.
Naming Ionic Hydrates
Hydrate: An ionic compound that contains one or more water molecules bound within its crystals.
Can be dehydrated by heating to yield the anhydrous compound (without water).
Naming Rules:
Name the anhydrous compound according to usual rules.
Add the word