MATINAL PARA ANIMADORES DE CANTO EN LA LITURGIA
QUE PARTICIPARON
A. Basic Trends
1. Metals tend to lose electrons and form cations 2. Nonmetals tend to gain electrons and form anions
3. Metalloids (semi-metals) have properties of both metals and nonmetals a. B, Si, Ge, As, Sb, Te, Po, At
4. Metallic character tends to increase as atomic number increases within a group
B. Atomic Size and Group Anomalies (Anomaly = oddity) 1. Hydrogen vs. Other Group I Elements
a. Very small, relatively high electronegativity (2.1)
b. Forms covalent bonds with nonmetals - other Group I elements form ionic bonds with nonmetals
2. Beryllium vs. Other Group II Elements
a. Small, electronegativity of 1.5 produces covalent bonds with many nonmetals (other group II's form ionic bonds)
b. BeO is amphoteric, other group II oxides are basic (form hydroxides) in solution
3. Boron vs. Other Group III Elements
a. Boron is a nonmetal/semimetal, all others are active metals 4. Carbon vs. Silicon (Group IV Elements)
a. Carbon readily achieves the octet by forming π bonds with oxygen in CO2
b. Silicon does not form π bonds with oxygen in discrete SiO2
molecules
(1) Si 3p orbitals do not easily overlap with oxygen 2p orbitals
(2) Si forms interlocking SiO4 tetrahedra which make up the crystalline structure of quartz
5. Nitrogen and Phosphorus (Group V)
a. Nitrogen forms a diatomic molecule with π bonds (N2)
b. Phosphorus forms aggregates based on the tetrahedral P4
molecule
(1) Single bonds
(2) Large atoms = weak π bonds 6. Oxygen and Sulfur (Group VI)
a. Oxygen forms a diatomic molecule with π bonds (O2)
b. Sulfur forms aggregates such as the cyclic S8 molecule, with all
single bonds 7. Halogens (Group VII)
a. Chlorine has an unexpectedly higher electron affinity than fluorine
(1) Small size of fluorine atoms bring unshared (lone) pairs close together, where they repel each other
C. Abundance and Preparation
1. Earth's Crust, Ocean, Atmosphere
Oxygen 49.2% Titanium 0.58% Silicon 25.7% Chlorine 0.19% Aluminum 7.50% Phosphorus 0.11% Iron 4.71% Manganese 0.09% Calcium 3.39% Carbon 0.08% Sodium 2.63% Sulfur 0.06% Potassium 2.40% Barium 0.04% Magnesium 1.93% Nitrogen 0.03% Hydrogen 0.87% Fluorine 0.03%
2. Major Elements in the Human Body
Oxygen 65.0% Potassium 0.34% Carbon 18.0% Sulfur 0.26% Hydrogen 10.0% Sodium 0.14% Nitrogen 3.0% Chlorine 0.14% Calcium 1.4% Iron 0.004% Phosphorus 1.0% Zinc 0.003% Magnesium 0.50%
3. Metallurgy - Obtaining a Metal from its Ore
a. Reduction of metal ions to atoms, usually using carbon as the reducing agent
2SnO(s) + C(s) + heat à 2Sn(s) + CO2(g)
2PbO(s) + C(s) + heat à 2Pb(s) + CO2(g)
Hydrogen as reducing agent
SnO(s) + H2(g) + heat à Sn(s) + H2O(g)
b. Electrolysis
(1) purification of highly active metals 4. Purification of Nonmetals
a. Liquefaction
(1) sequential expansion (cooling) followed by compression of a gas
b. Electrolysis
(1) Hydrogen from water c. Decomposition
18.2 The Group 1A Elements - The Alkali Metals A. Reactivities
1. With water
a. 2M(s) + 2H2O(l) à 2M+(aq) + 2OH-(aq) + H2(g)
2. Sodium forms oxides or peroxides
a. 4Na(s) + O2(g) à 2Na2O(s) (limited oxygen)
b. 2Na(s) + O2(g) à Na2O2(s) (excess oxygen)
3. K, Rb, Ce react with oxygen to form superoxides, containing the O2-
a. K(s) + O2(g) à KO2(s)
b. Superoxides react with water or carbon dioxide to release oxygen
4. Lithium reacts with nitrogen to form a nitride salt a. 6Li(s) + N2(g) à 2Li3N(s)
B. Biological Importance of Alkali Metals
1. Na+ and K+ are important in nerve conduction and muscle contraction 2. Li+ affects levels of neurotransmitters and is used to treat bipolar
disorder
18.3 Hydrogen A. Properties
1. Colorless 2. Odorless
3. Low boiling (-253°C) and melting (-260°C)points 4. Highly flammable
B. Purification of Hydrogen
1. Decomposition of methane in water, using heat, pressure and a catalyst
CH4(g) + H2O(g) à CO(g) + 3H2(g)
2. Cracking of hydrocarbons in gasoline production
C. Industrial Uses
1. Production of Ammonia by the Haber Process 2. Hydrogenating unsaturated vegetable oils
D. Hydrogen Halides 1. Ionic hydrides
a. Hydrogen and a Group I or II metal b. Hydride ion is H-
c. Hydrides are powerful reducing agents, explosive in water d. Examples include LiH and CaH2
2. Nonmetals + hydrogen (covalent hydrides) a. Examples include H2O, NH3, CH4, HCl
3. Metallic (Interstitial) Hydrides
a. Hydrogen and a transition metal
b. Hydrogen is absorbed by transition metals
(1) Amount of hydrogen depends on length of exposure (2) Potential method of storing Hydrogen fuel
18.4 The Group 2A Elements - The Alkaline Earth Metals A. Basicity of Oxides
1. MO(s) + H2O(l) à M2+(aq) + 2OH-(aq)
2. BeO has amphoteric properties B. Reaction with Water
1. M(s) + H2O(l) à M2+(aq) + 2OH-(aq) + H2(g)
2. Ca, Sr, Ba react at room temperature, Mg in boiling water C. Uses
1. Calcium phosphate in bone structure 2. Mg in metabolism and muscle function 3. Mg metal in flash bulbs and metal alloys D. Removal from "hard" water
1. Cation exchange resin replaces each Mg+2 and Ca+2 in water with 2 sodium ions
Note: Detergents are less soluble in hard water. There is noticeable difficulty, for instance, washing detergent out of one's hair when the concentration of Group II ions is high
18.5 The Group 3A Elements A. Boranes
1. B2H6 (diborane) and others (B5H9) are electron deficient and highly
reactive
B. Aluminum
1. Most abundant metal on earth 2. Oxide (Al2O3) is amphoteric
C. Gallium
1. Largest liquid range of any metal a. melts at 29.8°C
b. boils at 2400°C D. Indium and Thallium
1. The Inert Pair Effect
a. Lose one electron to form +1 ion (full s orbital) b. Lose three electrons to form +3 ion (octet)
18.6 The Group 4A Elements
A. Variation within the Group 1. C is a nonmetal
2. Si and Ge are semimetals 3. Sn and Pb are metals
4. All tend to form 4 covalent bonds to nonmetals (tetravalence) B. Carbon
1. Three allotropic forms (allotropic = two or more distinct forms)
Graphite Diamond Buckminster Fullerene
2. Carbon oxides
carbon monoxide carbon dioxide carbon suboxide
C. Silicon
1. Found in earth's crust in silica and silicates 2. Semimetal used in semiconductors
D. Germanium
1. Rare semimetal used as a semiconductor in electric devices E. Tin
1. Widely used in alloys
Bronze 20% Sn, 80% Cu
Solder 33% Sn, 67% Pb
F. Lead
1. Obtained from the galena ore (PbS)
2. Widely used in the anti-knock agent tetraethyl lead, (C2H5)4Pb
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AP Chemistry Chapter 21 - The Nucleus: A Chemist’s View 21.1 Nuclear Stability and Radioactive Decay
A. Radioactive Decay
1. Decomposition forming a different nucleus and producing one or more particles
a. Total mass number and atomic number must be conserved in any nuclear change n C He Be 1 0 12 6 4 2 9 4 + → + B. Zone of Stability
1. Of 2000 known nuclides, only 279 are stable with respect to radioactive decay
2. All nuclides with more than 83 protons (bismuth) are unstable 3. Light nuclides are most stable when
the neutron/proton ratio is 1 4. Heavier nuclides are most stable
when the neutron/proton ratio is greater than 1
5. Magic numbers
a. Special stability exists when the number of protons or neutrons is: 2, 8, 20, 28, 50, 82, 126
C. Types of Radioactive Decay 1. Alpha Emission
a. Alpha particle (α) is a helium nucleus ( He4
2 ), so it has a 2+ charge
b. Alpha emission is restricted almost entirely to very heavy nuclei
He Pb Po 24 206 82 210 84 + → + 2. Beta Emission
a. Beta particle (β) is an electron emitted from the nucleus during nuclear decay β 0 1 1 1 1 0n → p + −
b. Beta particles are emitted when a neutron is converted into a proton and an electron β 0 1 14 7 14 6C → N + − 3. Positron Emission
a. Positrons are particles that have the same mass as an electron, but a
positive charge
b. Positron emission arises from the conversion of a proton into a neutron and a positron
β 0 1 1 0 1 1p → n + + β 0 1 38 18 38 19K → Ar + +
2 4. Electron Capture
a. Inner orbital electron is captured by the nucleus of its own atom b. Electron combines with a proton and a neutron is formed
n p e 11 01 0 1 + → − Ag e 10646Pd 0 1 106 47 + − → 5. Gamma Emission
a. Gamma rays () are high-energy electromagnetic waves emitted from a nucleus as it changes from an excited state to a ground energy state
b. Gamma rays are produced when nuclear particles undergo transitions in energy levels
c. Gamma emission usually follows other types of decay that leave the nucleus in an excited state
D. Decay Series
1. In some cases, multiple decays are needed to produce a stable nuclide a. Original nuclide is called the "Parent" nuclide
b. Ensuing decay nuclides are called "daughter" nuclides
21.2 The Kinetics of Radioactive Decay A. Rate of Decay
1. The negative of the change in the number of particles per unit of time
N t N Rate ∝ ∆ ∆ − = kN t N Rate = ∆ ∆ − =
a. This is a first order rate law, so…
kt N N − = 0 ln
N0 = original number of nuclides (at t = 0)
N = the number of nuclides remaining at time t B. Half-Life (t1/2)
1. The time required for the number of nuclides to reach half the original value
k k
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Representative Radioactive Nuclides and Their Half Lives
Nuclide Half-life Nuclide Half-life
H 3 1 12.32 years Po 214 84 163.7 µseconds C 14 6 5715 years Po 218 84 3.0 minutes P 32 15 14.28 days At 218 85 1.6 seconds K 40 19 1.3 x 10 9 years 238U 92 4.46 x 10 9 years Co 60 27 10.47 minutes Pu 239 94 2.41 x 10 4 years 21.3 Nuclear Transformations A. Nuclear Transformation
1. The change of one element into another B. Methods of Transformation
1. Particle accelerators overcome the repulsive forces of the target nucleus a. Cyclotron
(1) Particle is accelerated from the inside and takes the spiral path to the target outside
b. Linear Accelerator
(1) Particle is accelerated down a linear track C. Transuranium Elements
1. Elements beyond Uranium
93 -112, 114, 116, 118 (as of May, 1999)
*** notice the absence of odd atomic numbers in the heavy nuclides
21.4 Detection and Uses of Radioactivity A. Detection
1. Geiger counter 2. Scintillation counter B. Dating by Radioactivity
1. Decay rate of unstable nuclides can be used to determine age of some objects
2. Carbon-14 dating (radiocarbon dating) a. Carbon-12 is stable
b. Carbon-14 decays, with a half-life of 5730 years
β 0 1 14 7 14 6C → N + −
(1) Living things take in carbon-12 and carbon-14, in a fixed ratio (2) When a living thing dies, the amount of carbon-12 does not
4 21.5 Thermodynamic Stability of the Nucleus
A. Mass Defect
1. The difference between the mass of an atom and the sum of the masses of its protons, neutrons, and electrons
For He4
2 :
2 protons: (2 x 1.007 276 amu) = 2.014 552 amu 2 neutrons: (2 x 1.008 665 amu) = 2.017 330 amu 2 electrons: (2 x 0.000 5486 amu) = 0.001 097 amu
total combined mass = 4.032 979 amu
Helium's atomic mass = 4.002 60_amu____ Mass defect = 0.030 38 amu B. Nuclear Binding Energy
1. The energy released when a nucleus is formed from nucleons 2. The energy required to break apart the nucleus
3. Mass defect is related to nuclear binding energy by the equation:
E = mc2 ∆E = ∆mc2
a. ( -0.03038 amu)(1.66 x 10-27 kg/amu) = -5.04 x 10-29 kg b. ∆E = (-5.04 x 10-29 kg)(3.00 x 108 m/s)2 = -4.54 x 10-12 J c. Binding energy per nucleon = 4.54 x 10-12 J = 1.14 J/nucleon
4 nucleons
C. Binding Energy per Nucleon
1. The binding energy of the nucleus divided by the number of nucleons it contains
2. High binding energy per nucleon results in greater stability a. The most stable nucleus is that of iron-56
5 21.6 Nuclear Fission and Nuclear Fusion
A. Nuclear Fission
1. Splitting a heavy nucleus into two nuclei with smaller mass numbers 2. The mass of the products is less than the mass of the reactants. Missing
mass is converted to energy B. Chain Reaction
1. A reaction in which the material that starts the reaction is also one of the products and can start another reaction
C. Critical Mass
1. The minimum amount of nuclide that provides the number of neutrons needed to sustain a chain reaction
D. Nuclear Fusion
1. Combining two light nuclei to form a heavier, more stable nucleus
A. Fusion Reactions
1. More energetic than fission rxns
2. Source of energy of the hydrogen bomb
3. Could produce energy for human use if a way can be found to contain a fusion rxn (magnetic field?)
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AP Chemistry Chapter 22 - Organic Chemistry 22.1 Alkanes: Saturated Hydrocarbons
A. Straight-chain Hydrocarbons
1. Straight-chain alkanes have the formula CnH2n+2
2. Carbons are sp3 hybridized
The First 10 Alkanes
# of Carbons Name Formula (CnH2n+2)
1 Methane CH4 2 Ethane C2H6 3 Propane C3H8 4 Butane C4H10 5 Pentane C5H12 6 Hexane C6H14 7 Heptane C7H16 8 Octane C8H18 9 Nonane C9H20 10 Decane C10H22 B. Structural Isomers
1. Same formula, but the atoms are bonded together in a different order 2. Different bonding order results in different properties
C4H10 Butane C4H10 2-methylpropane
C. Rules for Naming Alkanes (Nomenclature)
1. For a branched hydrocarbon, the longest continuous chain of carbon atoms gives the root name for the hydrocarbon
2. When alkane groups appear as substituents, they are named by dropping the -ane and adding -yl.
3. The positions of substituent groups are specified by numbering the longest chain of carbon atoms sequentially, starting at the end closest to the
branching.
4. The location and name of each substituent are followed by the root alkane name. The substituents are listed in alphabetical order (irrespective of any prefix), and the prefixes di-, tri-, etc. are used to indicate multiple identical substituents.
2 D. Reactions of Alkanes 1. Combustion reactions a. 2C2H6(g) + 7O2(g) à 4CO2(g) + 6H2O(g) 2. Substitution reactions a. CH4 +Cl2 →hv CH3Cl+HCl methane chloromethane 3. Dehydrogenation reactions a. 500 2 2 2 3 3 3 2 CH CH H CH CH CrOat°C→ = + ethane ethylene
E. Cyclic Alkanes (Cycloalkanes)
1. Alkanes in which the carbon atoms are arranged in a ring, or cyclic, structures
a. The 90° angle in cyclobutane is not nearly tetrahedral, therefore the molecule is quite unstable
2. Nomenclature
a. Rings are numbered to give the smallest substituent numbers possible
b. Largest substituents are given the lowest possible numbers
22.2 Alkenes and Alkynes A. Alkenes
1. Hydrocarbons that contain double bonds
a. The simplest alkene is ethene, or ethylene (C2H4)
b. Alkenes are nonpolar molecules B. Geometric Isomers
1. Isomers in which the order of atom bonding is the same but the arrangement of atoms in space is different
2. A molecule can have a geometric isomer only if two carbon atoms in a rigid structure each have two different groups attached
3
3. In some isomer pairs, one isomer is biologically active, while the other is not (specificity of enzymes is the cause)
C. Alkynes
1. Hydrocarbons with triple covalent bonds
a. The simplest alkyne is ethyne, or acetylene (C2H2)
b. Alkynes are nonpolar molecules D. Reactions of Alkenes and Alkynes
1. Addition reactions a. Hydrogenation 3 2 3 2 3 2 CHCH H CH CH CH CH = + Catalyst → Propene Propane b. Halogenation 3 2 2 2 2 3 2 2 2 CHCH CH CH Br CH BrCHBrCH CH CH CH = + → 1-Pentene 1-2-dibromopentene c. Polymerization
(1) small molecules are joined together to form a large molecule
22.3 Aromatic Hydrocarbons A. Structure of Aromatics
1. Hydrocarbons with six-membered carbon rings and delocalized electrons a. The simplest aromatic hydrocarbon is benzene (C6H6)
b. Aromatic hydrocarbons are nonpolar molecules B. Geometric Isomerism
1. ortho (o-) = two adjacent substituents
2. meta (m-) = one carbon between substituents 3. para (p-) = two carbons between substituents
4 C. Reactions of Aromatic Hydrocarbons
1. Substitution reactions
22.4 The Petrochemical Industry A Brief Narrative:
Petroleum contains molecules that vary from short chain (1 to 4 carbons) to very long chains (greater than 25 carbons). In the nineteenth century, kerosene and gas oil were the most desirable "fraction" of petroleum. With the advent of the internal combustion engine, the shorter chain molecules that make up gasoline became more important. Until that time they were considered a waste product of the purification of kerosene.
Rather than waste the kerosene-gas oil fraction, refineries take the long chain
molecules and break them into the smaller molecules of gasoline in a process called "cracking".
22.5 Hydrocarbon Derivatives
Classes of Organic Compounds
Class Functional Group General Formula
Alcohol
hydroxyl group (-OH) Alkyl halide Ether Aldehyde carbonyl group Ketone carbonyl group Carboxylic acid carboxyl group
5 Ester Amine amine group Examples:
Class: Ether Name: dimethylether
Class: Ether Name: diethylether
Class: Carboxylic acid Name: ethanoic acid
Class : Alcohol Name: ethanol (ethyl alcholol)
Class: Aldehyde Name: methanaldehyde
Class: Carboxylic acid Name: propanoic acid