Chapter 2: Atomic Structure
Chemical Nomenclature
if a metal forms two different ions (oxidation numbers are different) ∙ see examples ∙ using Stock method- the changes on those ions are indicated by Roman numerals, placed in parentheses. ∙ using classical system- the ion with the smaller charge (ox. number) is given the ending -ous and ion with greater charge has ending -ic.
Rules Assigned for Oxidation Numbers: #8
in binary compounds (with 2 elements) of metals and non-metals, the non-metal usually takes the least oxidation number (on the periodic table). ∙ see examples ∙ zinc takes ox. number of 2+, Zn²⁺ ∙ silver takes ox. number of 1+, Ag⁺
Rules Assigned for Oxidation Numbers: #7
in combinations involving non-metals, the oxidation number of the less electronegative element is positive and the more electronegative is negative. ∙ ex: OF₂→ O: 3.5, F: 4.0→ O+ and F-
Rules Assigned for Oxidation Numbers: #3
in compounds, the oxidation number of hydrogen is +1
Rules Assigned for Oxidation Numbers: #4
in compounds, the oxidation number of oxygen is -2. ∙ except in peroxide where ox. number of oxygen is -1.
Writing Chemical Formulas for Binary Compounds
metals + non-metals, cation + anion ∙ if ox. numbers are the same, charges will cancel each other. - see examples ∙ if ox. numbers are different, ignore charges and criss cross ox. numbers - see examples - use a parentheses if a polyatomic is present
Prefixes
mono= 1, di= 2, tri= 3, tetra= 4, penta= 5, hexa= 6, hepta= 7, octa= 8, nona= 9, deka= 10 ∙ mono is not written as it is understood, except for CO (carbon monoxide) ∙ see examples
Rules Assigned for Oxidation Numbers: #5
the algebraic sum of the oxidation numbers of all of the atoms in the formula of a compound is zero. ∙ see examples
Rules Assigned for Oxidation Numbers: #6
the algebraic sum of the oxidation numbers of the atoms in the formula of a polyatomic ion is equal to its charge. ∙ (ions have charges and compounds don't) ∙ see examples
Oxidation Number
the number of electron(s) an atom loses, gains, or shares to form a compound. ∙ referred to as the positive or negative charges on the ion ∙ or combining ability of an atom to from a compound ∙ ox. numbers for metals are same as group numbers (groups 1,2,3). Ox. numbers for some non-metals is the group number minus 8 (groups 4-8).
Rules Assigned for Oxidation Numbers
the ox. number of most elements must be calculated according to a few assigned rules.
Rules Assigned for Oxidation Numbers: #2
the oxidation number of a monatomic ion (one atomed) is equal to its charge. ∙ ex: Na⁺→+1, Cl⁻→-1, S²⁻→-2 ∙ ox. number for all diatomic molecules (H₂,N₂,O₂,F₂,Br₂,I₂,Cl₂) is zero.
Rules Assigned for Oxidation Numbers: #1
the oxidation number of an atom of a free element uncombined- pure element) is zero.
Writing Chemical Formulas for Binary Molecular Compounds
translate prefix numbers and names of elements into formulas. ∙ ex: dichlorine trioxide = Cl₂O₃
Nuclear Radius
∙ 0.001 pm ∙ density is 2 x 10⁸ metric tons/cm³ (1 metric ton = 1000 kg)
Atomic Radius
∙ 40 to 270 picometers (pm) (1pm = 10⁻¹²m) ∙ Most of the atomic radius is due to the electron cloud
Radiactive Elements
∙ All isotopes of all man-made elements are radioactive ∙ Some naturally occuring isotopes are radioactive - All isotopes of all elements beyond bismuth (atomic number 83) are radioactive
Types of Radioactive Decay: Alpha Emission
∙ Alpha particle (α) is a helium nucleus (⁴₂He), so it has a 2+ charge ∙ Alpha emission is restricted almost entirely to very heavy nuclei
Inferences from the Properties of Elements
∙ Atoms are neutral, so there must be positive charges to balance the negatives ∙ Electrons have little mass so atoms must contain other particles that account for most of the mass
Types of Radioactive Decay: Beta Emission
∙ Beta particle (β) is an electron emitted from the nucleus furing nuclear decay ∙ Beta paticles are emitted when a neutron is converted into a proton and an electron - ¹₀n → ¹₁p + ⁰₋₁β
Properties of Ionic Compounds Continued
∙ Forces are stronger when: the ion charges (Q1 and Q2) are large and the ions are small (d is small) - F (force between ions) = k (constant) Q1(charge on ion 1) Q2 (charge on ion 2) / d² (distance between ions) ∙ Ionic compounds are electrical insulators when SOLID ∙ Many are soluble in water
Types of Radioactive Decay: Gamma Emission
∙ Gamma rays (γ) are high-energy electromagnetic waves emitted from a nucleus as it changes from an excited state to a ground energy state ∙ Gamma rays are produced when nuclear particlces undergo transitions in energy levels ∙ Gamma emission ususally follows other types of decay that leave the nucleus in an excited state
Penetrating Ability of Radiation: Gamma Rays
∙ Greatest penetrating ability ∙ Protection requires sheilding with thick layers of lead, cement, or both
Penetrating Ability of Radiation: Alpha Particles
∙ Least penetrating ability due to large mass and charge ∙ Travel only a few centimenters through air ∙ Cannot penetrate skin ∙ Can cause harm through indigestion or inhalation
Fusion Reactions
∙ More energetic than fission reactions ∙ Source of energy of the hydrogen bomb ∙ Could produce energy for human use if a way can be found to contaian a fusion reaction (magnetic field?)
The Mole
∙ The amount of substance that contains as many particles as there are in exactly *12 grams of carbon-12* ∙ The amount of substance that contains the Avogadro number of particles
Avogadro's Number
∙ The number of particles in exactly one mole of a pure substance ∙ Avogadro's number = *6.022 x 10²³*
Penetrating Ability of Radiation: Beta Particles
∙ Travel at speeds close to the speed of light ∙ Penetrating ability about 100 times greater than that of alpha particles ∙ They have a range of a few meters in air
Isotopes
*Atoms of the same element have different masses* ∙ All elements of the same element have the same number of protons, but may vary in the number of neutrons ∙ Although isotopes have different masses, they do not differ significatnly in their chemical behavior
Gram/Mole Conversions
*Convert from grams to moles:* ∙ grams x (1g/molar mass) = moles *Convert from moles to grams:* ∙ moles x molar mass = grams
Particles in the Atom
*Electron:* ∙ Symbol: ⁰₋₁e; Mass number: 0 *Proton:* ∙ Symbol: ¹₁H; Mass number: 1 *Neutron:* ∙ Symber: ¹₀n; Mass number: 1
Designating Isotopes
*Hyphen notation:* ∙ Mass number is written after the name of the element - *ex:* hydrogen-2, helium-4 *Nuclear Symbol:* ∙ Composition of the nucleus using the element's symbol - *ex:* ²₁H → mass number = 2, atomic number = 1
Ions and Ionic Compounds
*Ions:* charged units ∙ Formed by the transfer of electrons between elements *Cation:* positive ion ∙ Metals form cations ∙ Na → Na⁺ + e⁻ *Anion:* negativd ion ∙ Nonmetals form anions ∙ S + 2e⁻ = S²⁻
The Discovery of the Electron
*Joseph John Thomson (1897)* ∙ Cathode ray tube produces a ray with a constant charge to mass ratio ∙ All cathode rays are composed of identical negatively charged particles (electrons)
Atomic Theory
*Modern atomic theory:* 1. All matter is made up of very tiny particles called atoms 2. Atoms of the same element are chemically alike 3. Individual atoms of an element may not all have the same mass. However, the atoms of an element have a definite average mass that is characterisitc of the element 4. Atoms of different elements have differnet average masses 5. Atoms are not subdivided, created, or destroyed in chemical reactions
The Nucleus
*Nucleons:* protons and neutrons *Nuclides:* Atoms identified by the number of protons and neutrons in the nucleus ∙ radium-228 or ²²⁸₈₈Ra
Structure of the Nucleus
*Protons* ∙ Positive charge, mass of 1.675 x 10⁻²⁷ kg ∙ The number of protons in the nucles determines the atom's identoty and is called the atomic number *Neutrons* ∙ No charge, mass of 1.675 x 10⁻²⁷ kg *Nucleur Forces* ∙ Short-range attractive forces: neutron-to-neutron, proton-to-proton, proton-to-neutron
Average Atomic Mass
*The weighted average of the atomic masses of the naturally occuring isotopes of an element* ∙ Atomic masses on the periodic table are average masses ∙ In calculations using atomic mass, we will round the masses to two decimal places before doing calculations
Transmutations
A change in the identity of a nucleus as a result of a change in the number of its protons
Nuclear Chain Reaction
A reaction in which the material that starts the reaction is also one of the products and can start another reaction
Nuclear Reactions
A reaction that affects the nucleus of an atom ∙ small amounts of mass are converted to LARGE amounts of energy ∙ *E = mc²*
Oxoanions
A series of related names ∙ -ate = more oxygen, -ite = less oxygen ∙ If they begin with H, add a prefix "hydrogen" ∙ Not a rule, it just happens for some of the ions
Nuclear Fission
A very heavy nucleus splits into more stable nuclei of intermediate mass ∙ The mass of the products is less than the mass of the reactants. Missing mass is converted to energy ∙ Small amounts of missing mass are converted to HUGE amounts of energy (E = mc²)
Unstable Nuclides
All nuclides beyond atomic number 83 are unstable and radioactive
Nuclear Equations
Alpha decay: ⁴₂He Beta decay: ⁰₋₁e
The Rutherford Experiment (1911)
Alpha particles (helium nuclei) were fired at a thin sheet of gold ∙ Assumed that the positively charged particles were bounced back if they approached a positively charged atomic nucleus head-on *(like charges repel one another)* ∙ Very few particles were greatly deflfected back from the gold sheet → *nucleus is very small, dense, and positively charges; most of the atom is empty space*
Molar Mass
The mass of one mole of a pure substance ∙ Units are grams/mole (or g/mol) ∙ Molar mass equals _________ units
Critical Mass
The minimum amount of nuclide that provides th number of neutrons needed to sustain an chain reaction
Atomic Number (Z)
The number of protons in the nucleus of each atom of that element ∙ Atoms are identified by their atomic number ∙ Because atoms are neutral, #protons=#electrons ∙ Periodic Table is in order of increasing atomic number
Radioactive Decay
The spontaneous disintegration of a nucleus into a slightly lighter and more stable nucleus, accompanied by emission of particles, electromagnetic radiation, or both
Half-Life
The time required for half the atoms of a radioactive nuclide to decay ∙ More stable nuclides decay slowly ∙ Less stable nuclides decay rapidly
Mass Number
The total number of protons and neutrons in the nucleus of an isotope
Balancing Nuclear Reactions
Total atomic numbers and mass numbers must be equal on both sides
Naming Binary Compounds
binary compounds of two non-metals ∙ use general formula: prefix, name of 1st element + prefix, stem, ide.
Binary Ionic
cation and anion or metal and non-metal ∙ name of cation followed by name of anion - ex: KBr = potassium bromide - ex: FeCl₂= iron (II) chloride/ferrous chloride ∙ two polyatomic ions are named binary - ex: OH⁻ = hydroxide→ KOH = potassium hydroxide
Binary Compounds
composed of 2 elements only ∙ name of all binary compounds must end with -ide.
Ternary Compounds
compounds composed of more than two elements ∙ naming the ternary compounds: name the metal, then the polyatomic/ name 1st polyatomic then 2nd - see examples, reverse criss cossing ∙ writing formulas: translate names into formulas. follow same method as binary compounds - use parentheses if ox. numbers are different
Isotopes
Atoms of the same element with different masses (A) are called *isotopes*. ∙ They have: equal number of protons, different numbers of neutrons ∙ ex: ²⁰₁₀Ne, ²¹₁₀Ne, ²²₁₀Ne
Recognizing Ionic Compounds
Compounds are ionic when they contain: ∙ A metal cation and a non-metal anion ∙ One or more polyatomic ions
Metals that Form Two Different Ions
Cu→ +1,+2 Fe→ +2,+3 Cr→+2,+3 Hg→ +1,+2 Sn→ +2,+4 Pb→ +2,+4
Hydrogen Isotopes
Hydrogen-1, hydrogen 2, and hydrogen-3
Acid Nomenclature
If the anion in the acid *ends in -ide*, change the ending to *-ic acid* and add the *prefix hydro-* . ∙ ex: *hydro*brom*ic acid*(HBr); *hydro*chlor*ic acid*(HCl) If the anion in the acid *ends in -ite*, change the ending to *-ous acid*. ∙ ex: hypochlor*ous acid* (HClO); nirt*ous acid*(HNO₂) If the anion in the acid *ends in -ate*, change the ending to *-ic acid*. ∙ ex: phosphor*ic acid* (H₃PO₄); nitr*ic acid* (HNO₃)
Ionic Compound Properties
Ionic compounds are generally: ∙ Solids at room temperature with high melting/boiling temperatures ∙ Hard ∙ Brittle and easily cleaved ∙ Poor heat conductors ∙ Poor electrical conductors (unless molten) ∙ Ionic crystals have distinctive shapes and can be cleaved ∙ Ionic compounds held together by strong *Coulombic forces* ∙ A lot of energy required to disrupt the crystal lattice and so melting and boiling points are high
Ionic Compounds
Ionic compounds are: ∙ Held together by electrostatic forces and they are always electrically neutral
Molecular Compounds
Ionic compounds contain metal ions. Compounds formed with only non-metal ions are usually *molecular* ∙ At the nanoscale: two or more elements combine ∙ Molecular formula shows the number and tye of elements combined ∙ Individal, independent units (molecules) form *Binary* ompounds contain two different elements If one element is hydrogen: ∙ It's written 1st in the formula and named first; other element is renamed with an -ide ending Other binary compounds (without H) ∙ Name elements in the formula order; 2nd element's name ends in -ide, *use prefixes*
The Crystal Lattice
Ionic compounds exist as *crystal lattices* ∙ Each is surrounded by many others ∙ *Formula unit:* smallest ratio of anions to cations
Conversions with Avogadro's Number
Review sample problems on pages 84 and 85 - solutions vary depending on the given data and the desired quantity
Nuclear Fusion
Light-mass nuclei combine to form a heavier, more stable nucleus
Monatomic Ions
Main group elements Add/lose enough electrons to "get to" the nearest noble gas ∙ Charge on ion = group A# or (grpA#-8) ∙ Each electron lost produces one positive charge ∙ Each electron gained adds one negative charge Transisiton Metals ∙ lost varying numbers of electrons ∙ old (and new) group number not very helpful
Mass Spectrometer
Measures the mass to charge the ratio of charged atoms and molecules ∙ number of peaks = number of different types of element; peak areas give relative amounts of each
Average Atomic Mass
Most elements a=occur as a mixture of isotopes ∙ *Atomic weight (average atomic mass) = ∑(fractional abundance)(isotope mass)* ∙ ∑ = sum of (sigma) ∙ Fractional abundance = fraction of each isotope ∙ Add together masses of the abundance of each isotope together to get the atomic mass
Polyatomic Ions
Multiple atom "units" with a net electrical charge ∙ Ex: NH₄⁺ = ammonium ion
Nuclear Radiation
Particles of electromagnetic radiation emitted from the nucleus during radioactive decay
Naming Ions and Ionic Compounds
Positive Ions ∙ Most are metal ions (except for NH₄⁺) Negative ions ∙ Monatomic ion ∙ Add -ide to name stem
